Chemistry 120 Handouts/Notes

Chemistry 120
Chapter Two Notes
Chapter Two: Measurements
Types of Measurements
Suppose I ask the question:
“How far is it from here to San Diego”
Your answer must contain two pieces information
A numerical value
A unit
Examples
About 70 miles
About 1.5 hours
Length
Distance from one location to another
Distance between the “center” of one atom and one it is attached to
Mass
The quantity of matter in whatever it is we are measuring
The more matter an object contains, the greater its mass
Mass vs. Weight
Weight measures the gravitational attraction between an object and the Earth (or moon,
Jupiter, etc.)
Weight is not the same as mass!
Example: Consider a 1-ton (2,000 pound) rock on the surface of the Earth.
How will its weight change if we move it to the moon To Jupiter
How will its mass change
Certain and Uncertain Digits
In every measurement we take, digits we are sure of are said to be certain.
A digit which must be estimated is said to be uncertain.
In taking a measurement, first determine which digits you know without question.
Then estimate exactly one uncertain digit.
When & Where To Estimate
Suppose you are taking a measurement with a ruler that has tenths (0.1) of a centimeter as
its smallest markings.
– You will estimate between the tenths mark, giving you an uncertain digit in the
hundredths (0.01) place.
Suppose you are using a thermometer which has whole degrees as its smallest markings.
– You will estimate between each degree to the tenth of a degree.
– Be careful not to drop the estimated digit!
• If the thermometer is exactly at the 15 degree mark, report 15.0 degrees.
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Chemistry 120
Chapter Two Notes
Certain and Uncertain Digits
In summary, in a given measurement, digits which are unambiguous are considered as
certain.
In any measurement there is always exactly one uncertain digit which results from estimation.
It is always the right-most digit.
In general, when taking a measurement on a device which measures to a given decimal
place, estimate one digit to the right of this decimal place.
– This would not apply for most electronic devices (like digital balances), which
already estimate for you.
Significant Figures
Clearly, when we take measurements, there is a limit as to how accurate our data can be.
For example, compare
A bathroom scale
A truck scale
Significant figures are those values in a measurement which we can rely on for accuracy.
The Rules for Determining Which Digits are Significant
Non-zero digits
– All non-zero digits (1-9) are always significant.
– Example: How many significant digits are there in each value below
48.9 7231.228
Leading zeros are zeros to the left of the first non-zero digit.
– Leading digits are not significant.
– So, the zeros in 0.000825 and 0.0134 are not significant.
– Why
Enclosed zeros are zeros between other non-zero digits.
– Enclosed zeros are always significant.
– Examples of enclosed zeros:
5,002 901.0802 101
Trailing zeros are those zeros which come at the end (to the right) of a value.
– Trailing zeros are significant if there is a decimal point in the value. If there is no
decimal, the zeros are ambiguous and generally not considered significant.
– Why
Practice With “Sig Digs”
How many significant digits do each of these values have
462.0 1230 1230.
0.030900 0.0005
Calculations: Multiplication & Division
When performing multiplication and division, count the number of significant digits in each
part of the calculation.
Determine which value has the least number of significant digits.
Your answer will be rounded to that many significant digits.
Examples
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Chemistry 120
Chapter Two Notes
Calculations: Addition and Subtraction
When adding or subtracting, consider the value with the least precision. Your answer must
extend to that level of precision.
– For example, suppose you are adding 60.82 and 7.831
• 60.82 goes to the hundredths place
• 7.831 extends to the thousandths place
• Therefore, your answer will round to the hundredths place.
VERY IMPORTANT: Note that in adding and subtracting, the number of significant figures in
the values does not matter!
Example
More Examples
23.649 – 17.2 =
8.75 + 9.414 + 105.32 =
(6.834 × 8.32) – 3.45 =
Scientific Notation
Many measurements in chemistry involve numbers that are extremely large or small.
Examples:
• Number of atoms in a glass of water.
• Distance between two atoms in a sugar molecule.
• Scientific notation is a convenient mathematical notation which simplifies working with
numbers of this magnitude.
How it Works
• Take the value under consideration, and move the decimal point after the first non-zero
digit.
Example: 582 5.82
• Next, consider how many places the decimal was moved, and in what direction it was
moved.
The decimal moved 2 places to the left.
Our new value is 10 2 , or 100 times, smaller than it was to begin with.
• Undo this by multiplying by 10 to the power of how many places the decimal moved.
This power is positive if the decimal moved to the left, negative if it moved to the right
Our value is 5.82 × 10 2
• Do not omit any significant digits!
Examples
Express these numbers in scientific notation.
23,894. 0.004289
20.000
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Chemistry 120
Chapter Two Notes
Calculations Involving Significant Digits
• In calculations involving the multiplication and division of numbers expressed in scientific
notation, it is helpful to remember the algebraic expressions:
Examples
Units of Measurement
• Units are used in measurements to tell us
– What type of measurement we are making
• Distance
• Time
• Energy
– The general magnitude of the measurement
• Use feet to measure a person’s height
• Use miles to measure distance between distant cities
• Use light years to measure distance between distant galaxies.
The English System
• The English System of units includes many familiar units
Inches, Feet, and Miles
Liquid Ounces
Pounds and Tons
• Despite its popularity, the English System has a serious flaw
SI Units
• SI units (also called metric units) are based on powers of 10
– There are 10 millimeters in a centimeter
– There are 10 centimeters in a decimeter
– There are 10 decimeters in a meter.
• You only need to know what each prefix means to be able to easily convert from one unit
to another.
• These prefixes are used in all types of measurements, including distance, volume, time,
and energy among many others.
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Chemistry 120
Chapter Two Notes
Converting Into Different Units
• Consider the following well known equality:
1 foot = 12 inches
• Suppose we divide both sides by 12 inches:
• Consider the following:
“Convert 18.5 feet to inches.”
• We need to get rid of feet, and end with inches.
– Strategy: Divide by feet (in denominator), multiply by inches (in numerator)
• Consider this example, using SI units:
“How many centimeters are in 9.86 m”
• Another example using the SI system.
“How many nanoseconds (ns) are there in 2.83 milliseconds (ms)”
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Chemistry 120
Chapter Two Notes
Multiple Unit Conversions
• In many measurements, such as in this problem, there are units in both the numerator
and the denominator:
“A car is traveling at 60. miles per hour. How fast is this in centimeters per
second”
A Side Note…
• By now you have noticed that units can be factored (or divided out) just like numbers
• You should also note that
– Units also can be multiplied
Ex. 2 ft × 2 ft = 4 (ft × ft) = 4 ft 2
Ex. 3 ft. × 2 lbs. = 6 ft.·lbs. (read “foot pounds”)
• You can only add and subtract measurements with the same units!
– Before adding and subtracting, perform any unit conversions to make the
measurements have the same units (of course, they must be the same type of
measurement!).
Area & Volume
• Translating length to area involves squaring the units as well as the values which
accompany them.
• For example, what is the area of a square which is 5.0 cm on each side
5.0 cm × 5.0 cm = 25 cm 2
• Let’s convert that to square meters:
• Common units of volume you should know
Liter(L) – the SI unit of volume
milliliter(mL) – another common unit
Cubic centimeter (cm 3 or cc) – a unit equivalent to the milliliter
1 mL = 1 cm 3 = 1 cc
• Dealing with volumes involves a similar procedure to that of areas.
A Volume Problem
• A rectangular block is 14.6 in. by 7.2 in. by 6.8 in. What is its volume in cm 3
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Chemistry 120
Chapter Two Notes
Density
• Density (d) is the ratio between mass (m) and volume(V):
density =
mass
volume
• Units of density for liquids are usually g/mL.
• For solids, we use the equivalent g/cm 3 .
• If something has high density, then a small volume of it will have a large mass.
– Ex. Which has greater mass, a liter of rocks or a liter of feathers
Density Problems
• The mass of an empty graduated cylinder is found to be 22.57 g. 7.25 mL of a liquid is
added to it. The graduated cylinder and liquid have a combined mass of 30.79 g. What
is the density of the liquid
• A 73.43 g cube of gold is dropped into a graduated cylinder whose volume reads 27.8
mL. After the cube sinks to the bottom, what volume reading will the graduated cylinder
have Note that gold has density 19.3 g/cm 3 .
Temperature
• Temperature is a measurement of the average kinetic energy of a body.
Kinetic Energy is the energy of motion, so the molecules in hot water are, on average,
moving more rapidly than those in colder water.
Note, temperature is not the same as heat!
Units of Temperature
• Fahrenheit (°F)
– Commonly used to measure temperatures in the U.S., but not very practical for
scientific use.
• Celsius (°C)
– Unit most commonly used in determining temperature in the laboratory.
• Kelvin (K)
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Chemistry 120
Chapter Two Notes
– The SI unit of temperature.
– 0 K is called “absolute zero.” It is impossible to for a system to have a temperature of
0 K (or lower).
The Temperature of the Freezing and Boiling Points of Water
Temperature
Scale
Freezing Point
Boiling Point
°F 32 212
°C 0 100
K 273.15 373.15
Converting Temperatures
• To convert between Celsius and Kelvin, use the following relationship:
T(in K) = T(in °C) + 273.15
To convert between °C and °F, use the equation:
°F = (1.8 × °C) + 32
A Temperature Problem
• The temperature in Paris is 23 °C. What is the equivalent temperature in Kelvin In
Fahrenheit
Accuracy & Precision in Measurements
• If data is accurate, this means that the results obtained are close to the “true and correct”
value(s).
• If a group of measurements give data which are close in value, these measurements are
said to be precise.
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Chemistry 120
Chapter Two Notes
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Chemistry 120
Chapter Three Notes
Basic Concepts About Matter
• Chemistry is the study of the properties and changes of matter
• What exactly is matter
◦ Matter is anything which has mass and takes up space (volume)
◦ Examples of matter:
• Sand (a solid)
• Water (a liquid)
• Air (a mixture of gases)
How do we learn chemistry
• Chemistry is an empirical science, meaning that it is based on the results of experiments.
• In the lecture we will study theories and laws based on many years of observations and experiments.
• In the laboratory we will verify many of these principles.
A Conceptual Approach
• During this course, we will focus on the fundamental concepts which define chemistry.
• Specifically, we will often look at matter at the smallest level (the submicroscopic level) to try to
reason why matter behaves the way it does on a larger, or macroscopic, scale.
• This is an essential skill for the aspiring chemist; a good imagination is all you need to develop it!
Atoms: A Brief Overview
• Atoms are the most fundamental units of matter we consider in this course
• All matter which we encounter in our daily lives contains an extremely large number of these tiny
particles
• There are many different types of atoms
◦ Some common examples are hydrogen, oxygen, gold, and sodium
• We will look closely at the structures of atoms at a later point in this course
Molecules
• Two or more atoms may join together to form a molecule
• Molecules are held together by bonds
◦ Specifically, these are called covalent bonds; more on these later!
• A diatomic molecule is made up of exactly two atoms (which may be the same or different)
Molecules
Physical States of Matter
• There are three common physical states of matter which we will consider in this class
◦ Gases
◦ Liquids
◦ Solids
• We compare the three states by asking two questions
◦ Does the substance have a definite shape, or does it take the shape of its container
◦ Does the substance have a definite volume
Solids
• Solids have a definite shape
◦ They do not assume the shape of their container
• Solids have a definite volume
• The particles making up a solid
◦ are close together
◦ do not move about, but vibrate in place
◦ tend to form organized patterns
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Chemistry 120
Chapter Three Notes
Liquids
• Liquids have an indefinite shape
◦ They take the shape of their container
• Liquids have a definite volume
• The particles making up a liquid
◦ are fairly close together, but not to the same extent as solids
◦ move freely throughout the liquid
◦ are not organized in any particular pattern
Gases
• Gases have an indefinite shape
◦ Like liquids, they take the shape of their container
• Gases have indefinite volume
• The particles making up a gas
◦ are generally far apart from one another
◦ move freely throughout their container in a random fashion
Classification of Matter
• We classify matter as either a pure substance or a mixture
• There are two types of pure substances
◦ Elements
◦ Compounds
• There are two types of mixtures
◦ Homogeneous mixtures
◦ Heterogeneous mixtures
Elements
• A pure sample of an element contains only one type of atom
◦ A sample of gold—an element—contains only gold atoms
◦ Helium contains only helium atoms
• There are over a hundred known elements
• Each element is assigned a name and a symbol
◦ Each symbol consists of one to three letters
◦ The first letter of the symbol is always capitalized; any other letters are always written in lowercase
• The symbols (and occasionally the names) are catalogued on the Periodic Table of the Elements
Some Atomic Symbols
H: Hydrogen O: Oxygen
C: Carbon N: Nitrogen
Na: Sodium
Cl: Chlorine
Cu: Copper
K: Potassium
• Note that some of these symbols are unusual!
Elements
• Some elements are generally only found as diatomic molecules
• We will call these seven elements the diatomic elements
H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2
Be sure you know these!
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Chemistry 120
Chapter Three Notes
Compounds
• Atoms come together in whole number ratios to form compounds.
• The chemical formula lists the ratio of these elements in the compound.
◦ Each unit of water contains two hydrogen atoms and one oxygen atom H 2 O
◦ Other examples: NaCl, SiO 2 , C 6 H 12 O 6
◦ This formula/ratio is always the same for a chemical compound
◦ Different compounds may have the same formula.
◦ The compound is said to have a definite composition
Mixtures
• The composition of a mixture is not fixed.
◦ Consider salt water, a mixture of H 2 O and NaCl.
• Is salt water always found in the same proportion The same atom-to-atom ratio
• Mixtures can be classified into two types:
◦ Homogeneous mixtures have all parts in the same state (gas, liquid or solid) and all parts must be
mixed together.
• If the parts of the mixture are visually inseparable, we will call the mixture homogeneous.
◦ Heterogeneous mixtures are simply those which are not homogeneous.
Separations
• Mixtures can be separated by physical methods
◦ A heterogeneous mixture of coffee grounds and water can be separated by filter paper
◦ The water can be removed from a salt water solution (a homogeneous mixture) by boiling the
water off. The salt will remain in the container.
• Compounds can only be separated into their individual elements by chemical means (i.e. through the
result of a chemical reaction)
Mixtures
• Ex. Are each of these mixtures homogeneous or heterogeneous Why
◦ Vodka (a mixture of water and ethyl alcohol)
◦ Cheerios in milk
◦ A mixture of oil and water
◦ A salt water solution
• Note that the term solution is often used to refer to homogeneous mixtures, especially for
compounds dissolved in water.
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Chemistry 120
Chapter Three Notes
◦ Air
◦ Dirty, dust-filled air
Classification Problems
• Classify each of these as an element, a compound, a homogeneous mixture, or a heterogeneous
mixture.
◦ Tap water
◦ Steel (an alloy of several metals)
• Note that alloys can be mixed in different proportions.
◦ Helium
◦ Mud
◦ Carbon dioxide
Properties of Matter
• We can describe matter in two ways:
◦ By its chemical properties, which describe how a type of matter interacts (or “reacts”) with another
type of matter.
• For example, hydrogen is able to react with oxygen to form water. Helium reacts with
virtually nothing.
◦ By its physical properties, which include all non-chemical properties.
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Chemistry 120
Chapter Three Notes
• For example, water is a liquid at room temperature, freezes at 0 °C, has a density of 1.0 g/mL,
and is both clear (we can see through it) and colorless.
Changes of Matter
• Common changes of matter are described in essentially the same way as properties are
• A chemical change is a change which involves a chemical reaction
◦ Bonds are formed and/or broken in a chemical change
◦ The chemical substances you end with are fundamentally different than what you began with
• A physical change is a change which does not involve a chemical reaction
◦ All changes of state of a given substance are physical changes
• Examples include ice melting and liquid water boiling
• Classify each change as either a physical change or a chemical change
◦ Gasoline evaporates off of the ground
◦ A glass vase is shattered
◦ Sodium reacts with chlorine, forming sodium chloride
◦ Dry grass burns in a large fire
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Chemistry 120
Elements and Symbols You Must Know
Group
1A
H
Hydrogen
You must memorize the symbols and names (with the correct spelling) for each element
listed on this periodic table. I recommend making flashcards to help you in this.
2A 3A 4A 5A 6A 7A
He
Helium
8A
Li
Lithium
Be
Beryllium
B
Boron
C
Carbon
N
Nitrogen
O
Oxygen
F
Fluorine
Ne
Neon
Na
Sodium
Mg
Magnesium
3B 4B 5B 6B 7B 8B 1B 2B
Al
Aluminum
Si
Silicon
P
Phosphorus
S
Sulfur
Cl
Chlorine
Ar
Argon
K
Potassium
Ca
Calcium
Sc
Scandium
Ti
Titanium
V
Vanadium
Cr
Chromium
Mn
Manganese
Fe
Iron
Co
Cobalt
Ni
Nickel
Cu
Copper
Zn
Zinc
Ga
Gallium
Ge
Germanium
As
Arsenic
Se
Selenium
Br
Bromine
Kr
Krypton
Rb
Rubidium
Sr
Strontium
Pd
Palladium
Ag
Silver
Cd
Cadmium
Sn
Tin
Sb
Antimony
I
Iodine
Xe
Xenon
Cs
Cesium
Ba
Barium
W
Tungsten
Pt
Platinum
Au
Gold
Hg
Mercury
Pb
Lead
Bi
Bismuth
U
Uranium

Chemistry 120
Standard States of the Elements
Group
1A
H 2
Hydrogen
This periodic table shows you the state (gas, liquid, or solid) of many important elements at room
temperature. Notice that all metals and metalloids except mercury are solids at room temperature.
The only two elements which are found as liquids are bromine and mercury. All the elements in the
last group (8A) are gases. The rest are fairly easy to memorize. On this periodic table I have
indicated the diatomic elements with their formulas; you will not see this on a "normal" periodic
table. I have left the boxes of the elements which are not important in this course blank.
2A 3A 4A 5A 6A 7A
He
Helium
8A
Li
Lithium
Be
Beryllium
Key: solid liquid gas
B
Boron
C
Carbon
N 2
Nitrogen
O 2
Oxygen
F 2
Fluorine
Ne
Neon
Na
Sodium
Mg
Magnesium
3B 4B 5B 6B 7B 8B 1B 2B
Al
Aluminum
Si
Silicon
P
Phosphorus
S
Sulfur
Cl 2
Chlorine
Ar
Argon
K
Potassium
Ca
Calcium
Sc
Scandium
Ti
Titanium
V
Vanadium
Cr
Chromium
Mn
Manganese
Fe
Iron
Co
Cobalt
Ni
Nickel
Cu
Copper
Zn
Zinc
Ga
Gallium
Ge
Germanium
As
Arsenic
Se
Selenium
Br 2
Bromine
Kr
Krypton
Rb
Rubidium
Sr
Strontium
Pd
Palladium
Ag
Silver
Cd
Cadmium
Sn
Tin
Sb
Antimony
I 2
Iodine
Xe
Xenon
Cs
Cesium
Ba
Barium
W
Tungsten
Pt
Platinum
Au
Gold
Hg
Mercury
Pb
Lead
Bi
Bismuth
U
Uranium

Chemistry 120
Chapter Four Notes
Chapter Four: Matter and Energy
Matter
Recall that matter
Has mass
Takes up space (i.e. has volume)
Matter also has one of four states
Gas
Liquid
Solid
Plasma (We will not be concerned with the plasma state in this course!)
States of Matter
The states of matter are classified by two parameters
Does it take the shape of its container
Does it completely fill its container
Gases
Take the shape of their container
Completely fill their container
Liquids
Take the shape of their container
Do not fill their container completely
Solids
Do not take the shape of their container
Do not fill their container completely
Which state a substance is in depends on two factors: Temperature and Pressure
Gases are comprised of molecules which generally are far apart from one another, and travel in
random paths.
The particles in liquids and solids are much closer together than those of gases
The particles of solids are generally “stuck in place”, but are able to vibrate
Liquid particles freely move across one another
Changes of State
Special terms are associated when matter changes from one state to another.
You may be familiar with the term “freezing,” which describes a change from liquid to solid.
Solid
Liquid
Gas
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Chemistry 120
Chapter Four Notes
Properties of Matter
We can describe matter in two ways:
By its chemical properties, which describe how a type of matter interacts (or “reacts”) with
another type of matter.
For example, hydrogen is able to react with oxygen to form water. Helium reacts with
virtually nothing.
By its physical properties, which include all non-chemical properties.
For example, water is a liquid at room temperature, freezes at 0 °C, has a density of 1.0
g/mL, and is both clear (we can see through it) and colorless.
Changes of Matter
Similarly, changes can be classified in the same way as properties:
Chemical changes involve a rearrangement of atoms, producing chemical compounds that were
not there before.
When iron (Fe) is exposed to oxygen on a wet day, rust (Fe 2 O 3 ) is formed.
Physical changes are those which do not involve a chemical change.
Boiling water (liquid H 2 O changes to gaseous H 2 O), glass shattering, a chemical
evaporating.
Metals
The metals include many elements found on the left side of the periodic table.
Many metals are known for having the following properties
They are malleable (meaning they are soft and easily shaped)
They are ductile (they can be twisted and drawn into a wire)
They can conduct both electricity and heat.
They tend to be lustrous (shiny).
All metals are solids at room temperatures except for mercury, which is a liquid.
Nonmetals
Nonmetals are generally found on the right side of the periodic table (except hydrogen, which is
placed on the left).
Their properties are generally the opposite of the metals.
Those which are solids tend to be brittle.
Most are poor conductors of electricity at room temperature (insulators) and do not
conduct heat well.
Some are gases at room temperature; others are solids. Bromine is the only other
element which is a liquid.
Metalloids
Metalloids are found between the metals and the nonmetals on the periodic table.
They include B, Si, Ge, As, Sb, and Te.
The properties of the metalloids are often a cross between those of the metals and the
nonmetals.
All the metalloids listed are solids at room temperature.
(I do not include Po and At with the metalloids. They are rather unstable and unimportant
to our studies at this point.)
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Chemistry 120
Chapter Four Notes
The Law of Conservation of Mass
One of the most fundamental statements about matter is described by this law, which states:
“Matter can neither be created nor destroyed in any chemical process”
Another consequence of this law is that one type of atom cannot be changed into another through
a chemical reaction.
If you start a chemical process with 10 million hydrogen atoms and 10 million oxygen
atoms, then you will end the process with 10 million hydrogen atoms and 10 million
oxygen atoms. ALWAYS!
Energy
Energy is an important subject in chemistry, as chemical reactions either give off or take in
energy.
The Law of Conservation of Energy states that “Energy can neither be created nor destroyed in a
process.”
Energy can be transferred between two systems in two ways
As work (w), which we will consider as energy put to some specific use, like making a
motor run.
As heat (q), which will be random energy(not directed to some useful purpose), like that
given off by your car’s engine.
Any change in energy is just the sum of the work and heat changes.
The SI unit of energy is the Joule (J). Other units include the kJ, and the calorie (cal)
1 calorie = 4.184 Joules
For example, suppose that burning a certain amount of gasoline produces 50 kJ of energy. If
only 20 kJ of it goes into doing work in a car engine, the remaining 30 kJ must be lost as heat.
Remember, energy cannot be destroyed!
An exothermic process is one which gives off more heat than it takes in.
For an exothermic reaction, q < 0.
For example, consider burning gasoline.
An endothermic process is one which brings in more heat than it gives off.
For an endothermic reaction, q > 0.
Specific Heat
Some substances require more energy to raise their temperatures than others.
For example, it requires much less energy to raise the temperature of 50. g of aluminum
by 10 °C than it would to raise 50. g of water by the same amount.
This difference is represented by a constant called the specific heat, c.
Its units are J/(g ·°C) or cal/(g ·°C) .
Transferring Heat
The amount of heat (q) absorbed or given off by a substance when it changes temperature can
be found using the following equation:
q = m × ΔT × c
m is the mass of the substance
ΔT is the change in the temperature; it equals (final T – starting T)
c is the specific heat of the substance
Note that heat always flows from an area of high temperature to one of lower temperature!
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Chemistry 120
Chapter Four Notes
Example
How much heat is required to raised the temperature of 50.0 g of aluminum from 32 °C to 47 °C
The specific heat of aluminum is 0.215 cal/(g ·°C).
A More Difficult Example
A 50.0 g aluminum block is heated to 75 °C, then dropped into a sealed container containing 350.
g of water at 15 °C. What will be the temperature of the block when it is finished cooling
More On Conservation
Note that in the last problem, heat was transferred from the aluminum to the water, but never
destroyed.
Einstein is credited with the famous equation:
E = mc 2
Where E is energy, m is mass, and c is a constant (the speed of light, not the specific heat!)
This shows that matter can be converted to energy.
More on Conservation
Note however, this will not apply to chemical reactions. We see this with nuclear reactions.
In this class, mass is always conserved.
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Chemistry 120
Chapter Four Notes
We can now state the Law of Conservation of Mass-Energy:
“Mass and energy can neither be created nor destroyed in a process, though they may
be interchanged.”
Kinetic and Potential Energy
Kinetic Energy (K.E.) is the energy of motion.
The amount of K.E. an object has is described by the formula
K.E. = ½mv 2
where m is the object’s mass, and v is its velocity.
Potential Energy is the energy of position.
Objects may be in a “high energy” position, and can convert this potential energy to kinetic energy
while moving to a “lower energy” position.
Examples:
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Chemistry 120
Chapter Five Notes
Chapter Five: Atomic Theory and Structure
Evolution of Atomic Theory
• The ancient Greek scientist Democritus is often credited with developing the idea of the atom
• Democritus proposed that matter was, on the smallest scale, composed of particles described
as atomos, meaning “indivisible”
• While atoms can in fact be broken down into smaller particles, the atoms of each element are
distinct from each other, making them the fundamental unit of matter.
Dalton’s Atomic Theory
• Compiling experimental information available during his lifetime, John Dalton described an
accurate picture of atoms in 1803 with his atomic theory.
• Dalton’s Atomic Theory maintains that
◦ All elements are made up of tiny, indivisible particles called atoms, which can be neither
created nor destroyed in reactions.
◦ Atoms of the same type of element are the same; those of different elements are
different.
◦ Atoms of different elements form compounds by combining in fixed, whole number ratios.
• This is also called “The Law of Definite Proportions”
◦ If the same elements combine to make more than one compound, there can be a
different, but definite, atom ratio for each compound.
• This is also called “The Law of Multiple Proportions”
• In actuality, atoms combining in the same ratio can make different compounds.
◦ A chemical reaction does not involve a fundamental change of the atoms themselves, but
of the way they are combined together.
Probing Further into Atomic Structure: The Electron
• One of the three particles which make up all atoms is the electron
• Benjamin Franklin observed that, when a cloth was rubbed across a glass rod, a charge was
developed on each.
• The charges, called positive (+) and negative (-), show an attractive force between opposites
and repulsion between identical charges.
The Atom: A Complete Picture
• Each atom contains at its core a nucleus, a region of positive charge.
• Positively charged particles, called protons, are contained in the nucleus.
• Electrons (negatively charged particles) “orbit” around the nucleus throughout the atom.
• Later experiments also confirmed that all atoms except hydrogen must contain one or more
neutral (non-charged) particles called neutrons.
• Note that the protons and neutrons are each almost 2,000 times more massive than an
electron; therefore, most of the mass of an atom is in the nucleus.
• The protons and neutrons are called the nucleons.
Characteristics of an Atom
• The number of protons in an atom is called the elements atomic number, which is symbolized
by Z.
• The number Z indicates which type of element that atom represents.
• These values are found in order on the periodic table.
• Example: Which element contains 5 protons 10 34
5-1

Chemistry 120
Chapter Five Notes
• Recall that the “Law of Conservation of Mass” indicated that atoms cannot be changed in
different elements by chemical means.
• So, protons are never “changed around” in a chemical reaction!
Characteristics of an Atom
• Atoms that have no charge must have the same amount of positive charge as negative
charge.
• Therefore, the number of electrons in an atom is equal to the number of protons in an atom
(Z) for any neutral atom.
• Atoms which contain more or less electrons than protons therefore must have a charge.
These are called ions.
Ions: A Lesson in Thinking Backwards
• Suppose an ion has exactly one more electron than it does protons.
◦ Will the ion be positively or negatively charged
• What if an atom lost two electrons
◦ Will the ion be positively or negatively charged
• For each electron an ion has more than it does protons, we indicate it with a – as a
superscript.
• For each electron an ion has less than it does protons, we indicate it with a + as a
superscript.
Examples of Ions
• A bromine atom normally has_______protons and________electrons.
• Suppose we add exactly one electron; now we have________protons and_______electrons.
• We symbolize this as Br - (or Br 1- ).
• An aluminum atom normally has________protons and_________electrons.
• Suppose the atom loses three electrons.
• We symbolize this ion as Al 3+ .
• Note that losing electrons is indicated with +, and gaining electrons is indicated with -.
Neutrons and Isotopes
• The number of neutrons in an atom cannot be easily predicted or found on the periodic table.
• Atoms of the same element (i.e. have the same number of protons) can have different
numbers of neutrons. They are called isotopes of one another.
• Examples:
◦ A carbon atom must contain 6 protons, but it may contain 6, 7, or 8 neutrons.
◦ Almost all hydrogen atoms contain zero neutrons. The isotopes of hydrogen which
contain one and two neutrons are called deuterium and tritium, respectively; they are
symbolized as D and T, but do not appear on the periodic table.
More on Isotopes
• The sum of the number of protons and neutrons in an atom is called the mass number of that
atom, and is symbolized by A.
• Isotopes are named by stating the name of the element, followed by A.
5-2

Chemistry 120
Chapter Five Notes
◦ Carbon (Z=6) with 6 neutrons is “carbon-12”
◦ Carbon with 7 neutrons is “carbon-13”
◦ Carbon with 8 neutrons is “carbon-14”
More on Isotopes
• Isotopes can be written in shorthand notation in two different ways,
A
Z
X
or
A
X
Examples
• For each of the following, indicate how many protons, electrons, and neutrons each atom or
ion possesses.
◦
35 Cl
◦
81 Br -
◦
27 Al 3+
The Mass of an Atom
• Recall that virtually all of the mass of an atom comes from its nucleus.
• Knowing the mass of protons and neutrons allows us to calculate the mass of one atom of a
particular isotope.
• Since most elements have more than one isotope, and these isotopes are mixed in nature, it
is not possible to provide an exact mass for the element.
• Instead, we consider a weighted average mass, based on the weight of each individual
isotope and its abundance in nature.
• The periodic table provides the atomic weight of each element, which corresponds to this
weighted average mass.
• These values are in atomic mass units (amu), a very small unit of mass.
◦ 1 amu = 1.66 × 10 -24 g
• For example, the mass of a single helium atom is 4.003 amu.
Calculating Atomic Weights
• The atomic weight of an element can be calculated if we know:
◦ The atomic weight of each isotope of that element, and
◦ The percent abundance of that element in nature.
• For example, suppose that 2 isotopes exist for an element X, and that the isotopes have
atomic weights of 64. amu and 66. amu. If there are 50.% of each in nature, what is the
atomic weight of X
5-3

Chemistry 120
Chapter Five Notes
Examples
There are three isotopes of potassium, as the data in the table below indicates. Calculate the
atomic weight of potassium from this data.
Isotope Mass(amu) % Abundance
39 K 38.963707 93.2581
40 K 39.963999 0.001171
41 K 40.961826 6.7302
Chlorine has only two naturally occurring isotopes. The atomic weight of 35 Cl is 34.96885 amu,
and its % abundance is 75.53%. What is the atomic weight of the other isotope, 37 Cl
Percent Composition
• The percent composition of a compound indicates the percentage of each element in the
compound by mass
• For example, suppose that a compound containing only nitrogen and oxygen is 36.85%
nitrogen
• From this, we see that, for every 100. grams of the compound, 36.85 grams of it comes from
the nitrogen in the compound
• What about oxygen
5-4

Chemistry 120
Chapter Five Notes
Suppose a compound is analyzed and is found to contain 6.00 grams of carbon, 2.00 grams of
hydrogen, and 8.00 grams of oxygen. What is the percent composition of the compound
What mass of hydrogen is in a 1.450 kg sample of this compound
5-5

Chemistry 120
Chapter Ten Notes
Chapter Ten:
Modern Atomic Theory
Light and Energy
• Electromagnetic radiation(EM) is an especially important form of energy for scientific study.
◦ Many types of “radiant” energy are included under this description, including visible light,
X-rays, and radio waves.
◦ EM can be described as waves, which can be characterized by their wavelengths (λ), the
distance between the peaks of each wave.
◦ We also may describe light as a particle, called a photon
Wavelength and Frequency
• The frequency (ν) of waves is measured as the number of waves which occur in a period of
time, usually a second.
◦ For example, there are 445 peaks in a second, we would indicate this as 445 s -1 .
◦ Alternatively, the Hertz(Hz) may be used
• 1 Hz = 1 s -1
• The frequency and the wavelength are inversely proportional
◦ That is, as the wavelength increases, the frequency decreases, and vice versa.
The Mathematical Relationship Between Wavelength and Frequency
c
ν =
λ
Where c is the speed of light
c = 3.00 ×10 8 m/s
Energy of EM Radiation
• The total energy of a single photon depends on the frequency (or wavelength) which that
photon is associated with.
• This is described by the equation
E = hν
where E is the energy (in J), νis the frequency (in s -1 ), and h is
Planck’s Constant (6.626 ×10 -34 J·s).
• Greater frequency (and shorter wavelengths) indicate higher energy photons.
• Provide a formula relating energy to wavelength rather than frequency:
The Electromagnetic Spectrum
• EM radiation can be classified into different regions based on its wavelength or its frequency.
◦ The longest wavelenths (smallest frequencies) correspond to radio waves.
◦ The shortest wavelengths (greatest frequencies) correspond to gamma rays.
10-1

Chemistry 120
Chapter Ten Notes
Electrons & Energy Levels
• Evidence suggested to scientists studying atomic structure that electrons “orbit” around the
nucleus in specific energy levels (also called shells).
◦ The farther the energy level is from the nucleus, the higher the energy of its electrons.
• Normally, electrons “fill up” the lower energy levels first, and as new electrons are added they
go into higher and higher energy levels.
◦ An atom is said to be in the ground state when this is true.
• An energy source, such as light or heat, can give electrons the energy necessary to “jump”
from its energy level to a higher one.
◦ In this situation, the atom is said to be in an excited state.
Emission of Light
• An atom in the excited state is unstable and must release the energy it gained in the first
place to return to the ground state.
◦ According to the Law of Conservation of Energy, this extra energy cannot simply
disappear
• One way in which this is accomplished is for the atom to give off this energy as light. This
process is called emission.
◦ The released light is called a quanta (energy packet) or a photon.
• Excited electrons return to lower energy levels.
◦ Depending on the atom itself and which energy levels are involved, EM radiation of
different wavelengths are given off.
Example
Suppose that a photon has 2.45 × 10 -19 J of energy.
(a) What is its wavelength in meters
(b) What is its frequency in s -1
(c) What type of EM radiation is this
So What is Light, Really
• The answer to this is not as simple as you might think
• Light (EM in general) is correctly described as both
◦ Particles (photons)
◦ Waves
• Light is not one or the other; it is both simultaneously!
10-2

Chemistry 120
Chapter Ten Notes
Quantization
• Notice that, for a given atom, electrons are only allowed to jump to discrete energy levels.
◦ In other words, an electron in the third level may fall to the first or second level, but not “in
between” the levels.
• As a result, only specific energy changes are allowed and can be observed for each atom.
◦ We say that these energy changes are “quantized” since only specific quantities of
energy can be emitted,
• As an analogy, compare a ball moving down a set of stairs (step-by-step) to one moving
down a ramp.
◦ Which corresponds to “quantized” energy values
Orbitals & Subshells
• In truth, electrons do not simply rotate about the nucleus in simple patterns.
• Instead, electrons are mostly contained in orbitals, which are shapes which describe the
regions where an electron is likely to be found.
• We will call a complete “set” of orbitals in a given energy level a subshell.
• Four orbital types are commonly encountered in modern chemistry.
◦ These are designated s, p, d, and f, and are listed here from lowest energy orbital type to
highest.
• Each individual orbital can hold, at most, two electrons in its ground state.
The s Orbital
• The s orbital is the least energetic of the orbitals, and, as a sphere, is also the most simplelooking.
• All energy levels have exactly one s orbital.
◦ Since any orbital can only hold two electrons, each s orbital is limited to this.
The p Orbital
• Starting with the 2 nd energy level, each level possesses three “sets” of p orbitals.
◦ There are three 2p orbitals, three 3p orbitals, etc.
• Recall that each orbital can hold up to two electrons, so the three orbitals combined will hold
up to six electrons.
• The p orbitals are more complex in shape than the s orbital, and electrons residing in them
typically have greater energy than those in an s orbital.
10-3

Chemistry 120
Chapter Ten Notes
d and f Orbitals
• Energy levels three and higher have a total of 5 sets of d orbitals.
◦ A full set of d orbitals can therefore hold up to_____electrons.
• Energy levels four and higher have a total of 7 sets of f orbitals.
◦ A full set of f orbitals can therefore hold up to_____electrons.
Electron Configurations
• Each element has its own unique electron configuration which describes which orbitals have
electrons in them and how many.
◦ For example, the electron configuration of boron is 1s 2 2s 2 2p 1 , which means that there are
• 2 electrons in an s orbital on the first energy level (called the 1s subshell)
• 2 electrons in the 2s orbital (the 2s subshell)
• 1 electron in a 2p orbital (the three p orbitals collectively make up the 2p subshell)
• Atoms always fill their lowest energy orbitals first, and successively higher ones as more
electrons are added.
◦ This is known as the Aufbau principle; aufbau is German for “building up”
• The order in which the electrons fill can be found on the periodic table.
◦ Notice that the periodic table is broken up into 4 distinct “blocks,” each of which indicates
where the highest energy electrons are.
Examples
Determine the electron configuration for each of the following atoms/ions:
Be
N
Na
V
10-4

Chemistry 120
Chapter Ten Notes
Shorthand Notation for Electron Configurations
• You can abbreviate the electron configuration with a special notation
◦ Consider V for example.
◦ The last noble gas before V was Ar (element 18)
◦ So, we can write [Ar]4s 2 3d 3 as vanadium’s electron configuration.
◦ What is the electron configuration of antimony in shorthand notation
Unusual Electron Configurations
• Unfortunately, the electron configurations we predict are not always correct
• For example, predict the electron configuration of copper
• The actual electron configuration is:
There are several other elements which do not “follow the rules”; however, at this point in your
study of chemistry, you do not need to be concerned with these exceptions
Spin
• Each electron is able to rotate about its central axis, a process we call spin
• Electrons may spin in one of two possible directions
• We refer to these directions as “spin up” and “spin down”
The Pauli Exclusion Principle & Hund’s Rule
• In a ground state atom, two electrons in the same orbital must have opposite spins; this is a
simplified version of a rule called the Pauli Exclusion Principle
• Another important principle, Hund’s Rule, states that, for ground state atoms, electrons will fill
unoccupied orbitals within a subshell before filling singularly-occupied ones
• For example, in a nitrogen atom, with electron configuration 1s 2 2s 2 2p 3 , the three
electrons in the 2p subshell reside in different orbitals (i.e. they do not pair up)
Energy Level Diagrams
• Energy level diagrams are a graphical form of the electron configuration, representing
electrons as arrows (up arrow for spin-up, down for spin-down)
and orbitals as lines
• The lines are listed from lowest energy at the bottom (i.e. the 1s orbital) to highest
10-5

Chemistry 120
Chapter Ten Notes
Example
• Draw the energy level diagrams for
a. titanium b. S 2- 10-6

Chemistry 120
Orbital Blocks
Group
1A
Here you see the periodic table separated into blocks, which should assist you in your study
of electron configurations. Note that helium belongs in the s block.
8A
2A 3A 4A 5A 6A 7A
He
Helium
s
block
3B 4B 5B 6B 7B 8B 1B 2B
d
block
p
block
f
block

Chemistry 120
Chapter Eleven Notes
Chapter Eleven
Chemical Bonding
Valence Electrons
• The electrons occupying the outermost energy level of an atom are called the valence
electrons; all other electrons are called the core electrons.
• The valence electrons, as we will see, are responsible for chemical bonding.
◦ Knowing the number of valence electrons an atom has is the single most important
information you can have in predicting how an atom will react chemically.
• Note that the valence electrons are not always the electrons with the highest energy, as we
will see with the transition metals
• For elements in group IA through VIIIA, the number of valence electrons an atom has is
simply its group number.
◦ So phosphorus, in group VA, has five valence electrons.
• Helium is the only significant exception; it has 2 valence electrons, even though it is in group
VIIIA
Lewis Dot Diagrams
• Lewis dot diagrams are simple symbols which show the symbol of an element surrounded by
as many dots as that atom has valence electrons.
• The first four electrons are drawn (in any order) on each of the four “sides” of the symbol.
• The next four electrons are “paired up” with the other four atoms (again, in any order).
• These symbols will be used to model chemical bonding.
Examples
• Draw the Lewis dot diagrams for selenium, rubidium, and manganese atoms.
The Octet Rule
• The Octet Rule states that atoms react in such a way as to give them eight electrons in their
outermost energy level (their valence shell).
• The Noble Gases, those elements in the period on the far right of the Periodic Table,
generally do not react with other atoms, as their valence shell is filled with eight electrons (or
two in the case of helium).
• By gaining or losing enough electrons to give an atom eight electrons in its valence shell, the
atom gains the stable features of a noble gas.
1

Chemistry 120
Chapter Eleven Notes
Getting an Octet
• Atoms usually achieve an octet in the following way:
◦ Metals lose their valence electrons, giving them an empty outer shell.
• Note that the ion now has the electron configuration of a noble gas.
◦ Nonmetals gain or share enough electrons to give themselves eight electrons in their
outer shell.
• Nonmetals generally share with other nonmetal atoms, or take electrons away from
metal atoms.
Example
• How many electrons will usually be gained or lost by each of the following atoms
◦ Bromine
◦ Barium
◦ Potassium
◦ Aluminum
◦ Oxygen
◦ Neon
The Duet Rule
• Hydrogen atoms attempt to pick up another electron, giving them the same electron
configuration as the noble gas helium; this is called The Duet Rule.
• Likewise, when lithium loses an electron to become Li + , it has the same electron configuration
as helium.
A Cautionary Note
• Gaining and losing electrons does not occur unless there are other atoms around to
encourage this.
◦ A sodium atom cannot just toss its valence electron away; it must give it to another atom
which wants to take it, like chlorine.
• Considering this, be sure that you do not confuse an atom of an element (which has its usual
number of electrons) with an ion (which has gained or lost electrons).
What is Bonding
• Bonding describes how atoms interact with each other in an attractive sense.
• There are three types of bonding:
◦ Ionic bonding
◦ Covalent bonding
◦ Metallic bonding
2

Chemistry 120
Chapter Eleven Notes
Ionic Bonding
• Ionic bonding occurs between ions of different charges.
• Often, but not always, ionic bonding occurs between the ion of metal with the ion of a
nonmetal.
◦ Sometimes, one of these ions is a polyatomic ion, such as sulfate (SO 4 2- ).
• Ionic bonding is generally the strongest of the bonding types.
◦ It requires a great amount of energy to separate the ions from each other, resulting in
high melting points.
Examples of Ionic Compounds
NaCl CaO CuBr 2 LiF
Formulas of Ionic Compounds
• An ionic compound should have no total charge in its formula.
• Ions of opposite charge are paired in such a way as to cancel out charges.
◦ Positively charged ions are called cations.
◦ Negatively charged ions are called anions.
• Example:
One Na + combines with one Cl - to give the neutral ionic compound NaCl.
Two Na + combine with one O 2- to give the neutral ionic compound Na 2 O.
• Notice that the cation is always listed first in the formula.
Formulas of Ionic Compounds
• Further examples
◦ Note that the formula of an ionic compound shows the smallest whole number ratio
between the ions.
◦ So, when Ca 2+ and O 2- combine, the formula of the product is CaO, not Ca 2 O 2 .
◦ Sometimes, the ratio between ions is more complex, like with Al 3+ and O 2- .
Example
• What is the formula of the ionic compound which contains ions of each of the two elements
below
sodium and sulfur
barium and nitrogen
selenium and calcium
3

Chemistry 120
Chapter Eleven Notes
Covalent Bonding
• Covalent bonds result from the sharing of electrons between two atoms.
• Recall that covalent bonding is usually found between nonmetal atoms.
◦ This is a bit oversimplified, but will work well for our purposes in this class.
• In following the octet rule, atoms share enough electrons with each other to give each eight
valence electrons (or two in the case of hydrogen).
• Covalent bonds are generally not as strong as ionic bonds; still, many are relatively strong.
◦ Example: Compare heating salt (ionic) and sugar (covalent).
Molecules
• Compounds which are held together by covalent bonds are called molecules
• This term is not appropriate for ionic compounds, which are made up of repeating unit cells
• For example an H 2 O molecule contains exactly three elements, but there is no “NaCl
molecule” consisting of only two atoms
Examples of Covalent Compounds
H 2 O
C 6 H 12 O 6
Cl 2 (an element)
C 6 H 6
CO 2
H 2 S
SO 3
Diatomic Elements
• Seven elements are essentially always found in nature as diatomic molecules, which are
molecules which contain exactly two atoms
• The formulas of these seven elements are:
H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , and I 2
• When we refer to oxygen gas for example, we mean O 2 , not individual oxygen atoms.
• These seven formulas will appear frequently throughout the course and must be learned!
Phosphorus, Sulfur, and Ozone
• The most stable form of phosphorus is a molecule containing four P atoms, which is given the
formula P 4 .
• Similarly, sulfur forms molecules which contain eight sulfur atoms each, and has the formula
S 8 .
◦ Unlike for the diatomic elements, chemists do not consistently use these formulas when
writing chemical reactions
• Oxygen also forms the molecule known as ozone, O 3 .
◦ Although both O 2 and O 3 contain only oxygen atoms, they have entirely different chemical
properties
4

Chemistry 120
Chapter Eleven Notes
Electronegativity
• The electronegativity of an atom indicates how strong its attraction for electrons is.
◦ The greater an atom’s electronegativity, the more strongly it pulls electrons toward it.
• Like the atomic radius and the ionization energy, electronegativity has a periodic trend:
◦ Electronegativity increases up a group and from left to right across a period.
• Fluorine is the most electronegative atom; in general, the closer a main group element is
to fluorine on the periodic table, the more electronegative it is.
• The noble gases have virtually no electronegativity.
Lewis Structures
• Lewis Structures are simple drawings that show the covalent bonds between atoms.
• As in Lewis dot diagrams, electrons are represented by dots.
◦ Since two electrons are required to make a bond, a pair of electrons between two atoms
indicates a bond.
Multiple Bonds
• Remember, atoms bond in order to gain an octet (or duet for hydrogen).
• In order to do this, atoms can also share more than one pair of electrons, up to a maximum of
three pairs.
◦ Sharing one pair of electrons forms a single bond.
◦ Sharing two pairs forms a double bond.
◦ Sharing three pairs forms a triple bond.
◦ There are not quadruple bonds!
Examples of Multiple Bonds
More on Lewis Structures
• It is much simpler when drawing Lewis structures to represent electron pairs as lines rather
than dots.
◦ However, always keep in mind that the lines stand for electron pairs!
5

Chemistry 120
Chapter Eleven Notes
Drawing Lewis Structures
• There are several steps involved in taking a formula and converting it to a Lewis structure;
just keep in mind the goals set out by the octet and duet rules.
1. Add up the total number of valence electrons in the formula.
Ex. NH 3 has 5 + (3×1)=8 v.e.
Note that if the molecule is an ion, each negative charge adds one electron to this count, and
each positive charge takes one away.
2. Draw a simple “skeleton” of the molecule, showing which atoms are connected to which. If
there is one “central” atom (the one surrounded by the others) it is usually the least
electronegative.
Note: Hydrogen is never a central atom!
3. Figure out how many electrons you have left. Remember that each bond stands for two
electrons, so subtract two from the total number of valence electrons for every bond you drew
in the previous step.
8 – 6 = 2 electrons left
4. Give enough electron pairs to each atom in the structure to give each atom (except
hydrogen) an octet, using the following guidelines:
-Electrons go to the most electronegative atoms first.
-You must stop adding electron pairs when you run out of electrons (from the original number
you started with in step 1).
5. This step is not always necessary:
If you have run out of electrons, and some atoms still need octets, change a lone pair on the
least electronegative adjacent atom (except hydrogen) to a double bond between the two
atoms. If necessary, you may need to make a triple bond by taking away another lone pair.
Note that fluorine will not share its lone pairs.
6

Chemistry 120
Chapter Eleven Notes
Examples
Draw Lewis Structures for each of the following molecules:
•
CH 4
H 2 O
HCN
NO 3
-
N 2 H 2
C 2 H 2
7

Chemistry 120
Chapter Eleven Notes
Exceptions to the Octet Rule
• Boron and beryllium often will not be able to gain a complete octet.
◦ For example, consider BH 3 and BF 3
• Atoms in period 3 and below only may exceed the octet rule for reasons we will not consider
here. You will often see this in cases where the central atom is P, S, Cl, Br, or I; others exist
as well.
◦ For example, consider PCl 5, SF 6 ,and SO 3 .
8

Additional Chapter 11 Topics
Periodic Trends
• By comparing the position of one element to another on the periodic table, it is often possible
to make comparisons between the properties of those elements; these are a result of periodic
trends.
• In this class we consider three of these trends:
◦ Electronegativity
◦ Atomic radius
◦ Ionization Energy
Atomic Radius
• The radius of an atom tends to increase as you move down a group in the periodic table.
◦ Explanation: Elements in lower groups of the periodic table have electrons in outer
energy levels. These electrons are not held as tightly by the nucleus, so they can travel
further away from it.
• The radius of an atom tends to decrease from left to right across a period.
◦ Explanation: The valence electrons of atoms in the same period fill the same energy
levels. However, the positive charge in the nucleus is increasing as we move from left to
right on the periodic table as the number of protons is likewise increasing. This pulls the
electrons in closer, making the radius smaller.
Atomic Radius: Ions
• Recall that when an ion is formed, an atom gains or loses electrons, while the number of
protons remains the same
• When an atom loses an electron to form a cation, the positive charge of the nucleus pulls the
remaining electrons closer towards the nucleus
• A monatomic cation is smaller than the corresponding neutral atom
◦ Na + has a smaller atomic radius than Na
◦ Mg 2+ has a smaller atomic radius that Mg
• When an atom gains an electron to form an anion, the increased repulsion of the additional
electrons results in electrons moving farther from the nucleus
◦ Cl - has a larger atomic radius than Cl
◦ O 2- has a larger atomic radius than O
• When comparing atoms and ions with the same number of electrons, the largest ions are
those with the most negative charge
Atomic Radius: Ions
Rank each of the following atoms/ions in order from largest to smallest
I. Ne Al 3+ O 2- F - Na + Mg 2+
II. Li K N Ne Cs

Ionization Energy
• The ionization energy is the amount of energy required to remove an electron from a mole of
atoms in their gas state of a particular element.
X (g) X + (g) + e -
◦ The first ionization energy is the amount of energy required to remove the first electron.
◦ The second ionization energy is the amount of energy required to remove the second
electron. Etc.
• Ionization energy tends to decrease down a group.
◦ Explanation: Electrons in the outer shells are further from the nucleus and feel its pull less
than the core electrons do. Less energy is required to remove them.
• Ionization energy tends to increase from left to right across a period.
◦ Explanation: Protons are added to the nucleus from left to right across the table, and
these protons exert a greater pull on the electrons, requiring more energy to remove
them.
◦ Why is the first ionization energy of oxygen less than that of nitrogen
Comparing Ionization Energies
• Generally, the first ionization energy is less than the second, which is less than the third, etc.
• If an atom has lost enough electrons to give it an octet in its outer shell, it requires a
tremendous amount of energy to remove any electrons beyond this.
Covalent Bonds
• Recall that
◦ Covalent bonds result from the sharing of electrons between atoms.
◦ Nonmetals bond with other nonmetal atoms through covalent bonding.
◦ The strength of a covalent bond is generally less than that of an ionic bond.
• A question:
◦ Do atoms involved in covalent bonding share electrons equally
Polarity and Dipole Moment
• If two atoms which are covalently bonded have substantially different electronegativity values,
then the bond is said to be polar covalent.
◦ The electrons are not shared equally in the bond.
◦ The electrons will tend to stay closer to the more electronegative atom.
• The molecule might have a dipole moment, which is shown by drawing a special arrow
pointing towards the more electronegative atom. More on this later…

Shapes of Molecules
• From Lewis structures we can usually figure out the three-dimensional structure of simple
molecules.
• VSEPR theory (stands for Valence Shell Electron Pair Repulsion) helps us to understand the
shapes that molecules take.
• According to VSEPR, electron pairs move as far away from each other (in three dimensions)
as possible.
◦ Remember that electrons all have negative charge, so they repel one another.
The Basic Geometries of Molecules
• Depending on the number of attached atoms and lone pairs, a basic geometry can be
described for an atom in a molecule.
◦ Linear: 2 attached groups
◦ Trigonal Planar: 3 attached groups
◦ Tetrahedral: 4 attached groups
◦ Trigonal Bipyramidal: 5 attached groups
◦ Octahedral: 6 attached groups
Assigning Shapes to Molecules
• A shape is assigned to each central atom in a molecule.
◦ To be a central atom, an atom must have two or more atoms bonded to it.
• The number of atoms bonded to the central atom is added up. Call this m.
◦ Careful! Add the number of atoms, not the number of bonds!
• The number of lone pairs on the central atoms is added up. Call this n.
Assigning Shapes to Molecules
• Take these numbers and put them in the following formula:
AX m E n
• The A in this formula stands for the central atom, the X for the attached atoms, and the E for
the lone pairs.
• From this code we can name the shape of the molecule.
Linear
• AX 2 is called the linear shape.
• The bond angle (atoms between the three bonds) in this shape is 180°.
Trigonal Planar and
Bent/V-Shaped
• The shape associated with AX 3 is trigonal planar because it resembles a flat triangle.
• The shape associated with AX 2 E 1 is called bent/V-shaped.
◦ When naming a shape we pretend to only “see” the atoms, not the lone pairs. That is
why we have two different shape names for the same geometry.
• Both of these shapes have bond angles of approximately 120°, which may vary slightly in
some molecules.

Tetrahedral Geometries
• The shape associated with AX 4 is called tetrahedral, with all atoms at the four points of a
pyramid.
• Replacing an atom with an electron pair gives AX 3 E 1 , which is called trigonal pyramidal.
• Changing another atom to a lone pair gives the the AX 2 E 2 code, which is called the bent/Vshaped
shape.
◦ This is the same shape name as we saw earlier, but it has a different geometry and
different bond angles.
• The bond angles associated with these three are approximately 109.5°.
Trigonal Bipyramidal and Octahedral Shapes
• The AX 5 shape is called trigonal bipyramidal, and resembles two pyramids sharing a common
face.
◦ The bond angles associated with this shape are both 90° and 120°, depending on where
you look.
• The AX 6 shape is called octahedral.
◦ The bond angles for this shape are 90°.
• Other shapes within these geometries are not included in this course, but they do exist.
Dipole Moments
• Molecules which have fairly electronegative atoms may have a dipole moment.
• Since electrons are drawn towards more electronegative atoms, a “partial negative charge”
(δ-)develops in this region of the molecule, and a partial positive charge (δ+)develops on the
other side.
• Consider CH 3F. .

Dipole Moments
• In thinking about dipole moments, we have to consider the three-dimensional structure of the
molecule.
• Dipoles can cancel each other out, if the same type of bonds are oriented in opposite
directions.
Example
• Example: Which of these molecules have dipole moments Which, if any, have polar bonds
but no dipole moment
CCl 4 NH 3 H 2 O CO 2

Chemistry 120 Molecular Geometry
Molecular Geometry
AX m E n
Code
Approximate Bond
Angles (degrees)
Hybridization
of Central Atom*
2-dimensional
Representation
3-dimensional
Representation
Linear AX 2 180 sp X A X X A
X
X
X
Trigonal Planar AX 3 120 sp 2 A
X
X
X
A
X
Bent/V-shaped AX 2 E 1

Chemistry 120 Molecular Geometry
Molecular Geometry
AX m E n
Code
Approximate Bond
Angles (degrees)
Hybridization
of Central Atom
2-dimensional
Representation
3-dimensional
Representation
X
X
Trigonal Bipyramid AX 5 90 & 120 dsp 3 X
A
X X
X
X
A
X
X
X
X
See-saw AX 4 E 1 90 &120 dsp 3 X A
X X
X
A
X
X
X
X
X A
T-shaped AX 3 E 2 90 dsp 3
X
X
A
X
X
X
Linear AX 2 E 3 180 dsp 3 X A X A
X
Note: None of the shapes on pages 2 or 3 of this handout will be tested in CHEM 120, as they all have central atoms which exceed an octet.
These are included to assist you in your later studies in General Chemistry.
Page 2

Chemistry 120 Molecular Geometry
Molecular Geometry
AX m E n
Code
Approximate Bond
Angles (degrees)
Hybridization
of Central Atom
2-dimensional
Representation
Octahedral AX 6 90 d 2 sp 3 A
Square Pyramid AX 5 E 1 90 d 2 sp 3 X
A
X
X
X
X
X
X
X
X
X
X
3-dimensional
Representation
X
X
X
X
X
A
X
X
A
X
X
X
X
Square Planar AX 4 E 2 90 d 2 sp 3 A
X
X
X
X
X
X
A
X
X
Note: None of the shapes on pages 2 or 3 of this handout will be tested in CHEM 120, as they all have central atoms which exceed an octet.
These are included to assist you in your later studies in General Chemistry.
Page 3

Chemistry 120
Chapter Six Notes
Chapter Six:
Nomenclature—The Names of Chemical Compounds
Types of Ions
• Ions are classified according to how many atoms they contain
◦ If an ion is derived from a single atom, it is called a monatomic ion.
• Examples include Na + , O 2- , and Pb 4+ .
◦ Ions which are derived from two or more atoms are called polyatomic ions.
Polyatomic Ions
• Polyatomic ions can be described as molecules which have collectively gained or lost
electrons to become an ion.
• Being molecules, the ions themselves are held together by covalent bonds.
• From this, you would expect the atoms in a polyatomic ion to all be nonmetals.
◦ Examples: NO 3 - , ClO 4 - , NH 4
+
• However, some polyatomic ions contain metals which are able to covalently bond.
◦ Examples: Cr 2 O 7 2- , MnO 4
-
• All but one of the more common polyatomic ions are anions.
• The ammonium ion, NH 4 + , is the only common polyatomic cation.
• The common polyatomic ions must be memorized, and you must learn to recognize them in a
formula on sight.
1- anions
OH - hydroxide
-
NO 3
CN - cyanide
-
ClO 3
OCN - cyanate
-
BrO 3
SCN - thiocyanate
-
IO 3
-
MnO 4 permanganate
-
C 2 H 3 O 2
2- anions
2-
CO 3 carbonate
2-
SO 4
2-
CrO 4 chromate
2-
Cr 2 O 7
2-
C 2 O 4 oxalate
2-
SiO 3
2-
S 2 O 3 thiosulfate
2-
O 2
3- anions
3-
PO 4 phosphate
3-
BO 3
nitrate
chlorate
bromate
iodate
acetate
sulfate
dichromate
silicate
peroxide
borate
1

Chemistry 120
Chapter Six Notes
Classifications of Compounds
• When naming inorganic compounds, two different naming systems are used, depending on
the type of compound you are considering.
• Binary nonmetals (sometimes called molecular compounds) contain atoms from exactly two
different nonmetals.
◦ Examples: NO, CO 2 , CO, SO 2 , H 2 S.
• Ionic compounds are composed of a cation and an anion.
◦ For the purposes of this course, if a compound is not a binary nonmetal, it is an ionic
compound.
◦ The ions may be monatomic and/or polyatomic.
◦ Examples: NaCl, CaCl 2 , NH 4 Cl, CaCO 3 , Na 3 PO 4 .
Examples
• Classify each of these compounds as binary nonmetals or as ionic compounds.
◦ KBr
◦ Na 2 CO 3
◦ NaCN
◦ N 2 O 4
◦ NaH 2 PO 4
◦ P 2 O 5
◦ NH 4 NO 3
Naming Binary Nonmetals
• Nonmetals can bond with one another in many possible combinations.
◦ For example, nitrogen and oxygen make compounds such as NO, NO 2 , and N 2 O.
• Although these compounds contain the same elements (and may even contain them in the
same ratio), these chemicals have considerably different chemical and physical properties.
• Therefore, each must be given its own distinct name to distinguish it.
• The compound is named by putting prefixes before the name of the element to indicate how
many atoms of each type are in the molecule.
• The ending of the name of the second element is changed to “-ide”.
prefix[1 st element] prefix[2 nd element]
2

Chemistry 120
Chapter Six Notes
Naming Binary Nonmetals
The first ten prefixes must be memorized:
Number of Atoms
in Molecule Prefix
1 mono-
2 di-
3 tri-
4 tetra-
5 penta-
6 hexa-
7 hepta-
8 octa-
9 nona-
10 deca-
• For example, a compound with formula S 2 O 3 would be named disulfur trioxide.
• Exceptions:
◦ Do not use the prefix “mono-” with the first element in the formula; just stating the name of
the element implies that there is only one of it in the formula.
• For example, CO 2 is carbon dioxide, not monocarbon dioxide.
• What is the name of SO 3
◦ If the last letter of the prefix is an “a” or an “o”, and the first letter of the element after the
prefix is an “o”, drop the last letter of the prefix; it looks and sounds awkward.
• Example: CO is carbon monoxide, not carbon monooxide.
• P 2 O 5 should be named diphosphorus pentoxide, which is preferable to diphosphorus
pentaoxide (which is technically acceptable).
Common Names
• In three cases, the naming rules are totally ignored, and the common name is used.
• These three are
◦ H 2 O, which is always called water.
◦ NH 3 , which is always called ammonia.
◦ PH 3 , commonly known as phosphine.
• Other common names do exist, but are omitted here (they are less common).
Hydrogen-Containing Formulas
• When hydrogen appears first in the formula of a binary nonmetal, do not use prefixes at all.
Simply name the compound without them.
◦ There will be as many hydrogens bonded to the other nonmetal as are necessary to give
it a complete octet.
◦ So, H 2 S (g) is named hydrogen sulfide, and HCl (g) is named hydrogen chloride.
◦ Note that the (g) means “gas”, an important fact which must be specified in the formula
for these types of compounds.
• If the hydrogen-containing binary nonmetal is dissolved in water, a different name is used.
3

Chemistry 120
Chapter Six Notes
• In the formula, the notation (aq) is used, meaning aqueous, to indicate that it is dissolved in
water.
• These compounds are named as acids.
◦ These are not to be confused with oxyacids, which will have a different naming system.
Naming Acids of Binary Nonmetals
• The prefix “hydro-” is put before the name of the element.
• The ending of the name of the element is changed to “-ic”, followed by the word “acid”.
◦ Examples:
HCl (aq) is hydrochloric acid
HBr (aq) is hydrobromic acid
HI (aq) is hydroiodic acid
H 2 S (aq) is hydrosulfuric acid
CAUTION: Do not confuse this with sulfuric acid, which is H 2 SO 4 !
Examples
Name each of the following:
N 2 O 5
H 2 Se (aq)
SO 2
NH 3
HF (aq)
XeF 4
HF (g)
H 2 S (aq)
Naming Ionic Compounds
• The naming system for ionic compounds is different than that for binary nonmetals.
• Most importantly: The prefixes used for binary nonmetals (mono, di, tri, etc.) are never used
to name ionic compounds.
◦ This is, by far, the most common mistake made by students in naming compounds.
Naming Individual Ions: Cations
• Some cations always have the same charge in virtually all compounds.
• These include
◦ Group I cations, which only form 1+ ions
•
Ex.: Li + , Na + , K + , Rb + , Cs +
◦ Group II cations, which only form 2+ ions
• Ex.: Mg 2+ , Ca 2+ , Sr 2+ , Ba 2+
◦ Group III cations, which only form 3+ ions
• Al 3+ , (Ga 3+ if it is forming an ionic compound)
• Note that this is not true for elements below Ga
4

Chemistry 120
Chapter Six Notes
• Three additional common metals also only form one ion.
• These must be memorized. They are
◦ Zinc: Zn 2+
◦ Cadmium: Cd 2+
◦ Silver: Ag +
• For all other metals, the charge must be specified in naming the cation.
• The name of the individual cation is stated by simply stating the name of the metal, followed
by the word “ion”.
• Examples:
◦ Na + is the sodium ion.
◦ Ba 2+ is the barium ion.
◦ Zn 2+ is the zinc ion.
• The new method for naming all other cations places the charge of ion in Roman numerals in
parenthesis after the name of the metal.
◦ Cu 2+ is the copper (II) ion.
◦ Cu + is the copper (I) ion.
◦ Pb 4+ is the lead (IV) ion.
◦ Pb 2+ is the lead (II) ion.
• Remember, you only use Roman numerals for those ions that form multiple charges.
• The old names for many ions are still in use and must be learned.
• The names of the ions must be memorized, noting the following:
◦ The ion possessing the higher charge ends in “-ic”
◦ That possessing the lower charge ends in
“-ous”
◦ Ex.: Fe 3+ is the ferric ion, Fe 2+ is the ferrous ion.
Ion Common Name Ion Common Name
Cu 2+ cupric ion Cu + cuprous ion
Fe 3+ ferric ion Fe 2+ ferrous ion
Pb 4+ plumbic ion Pb 2+ plumbous ion
Sn 4+ stannic ion Sn 2+ stannous ion
Hg 2+ 2+
mercuric ion Hg 2 mercurous ion
Cr 3+ chromic ion Cr 2+ chromous ion
Co 3+ cobaltic ion Co 2+ cobaltous ion
Mn 3+ manganic ion Mn 2+ manganous ion
Examples
Name the following cations:
Rb +
Hg 2
2+
Fe 3+
Cd 2+
Hg 2 +
Ca 2+
5

Chemistry 120
Chapter Six Notes
Naming Individual Ions: Anions
• The charge on a monatomic anion is simply 8 minus the group number of the element.
◦ For example, nitrogen is in group 5, so it will only make a N 3- ion. (8-5=3)
• Monatomic anions are named by simply changing the ending of the element’s name to “ide”.
• Examples:
◦ S 2- is the sulfide ion
◦ O 2- is oxide ion
◦ P 3- is the phosphide ion
• CAUTION: Do not confuse these monatomic ions (nitride, sulfide, phosphide, etc.) with the
similar-sounding oxoanions (nitrate, sulfate, phosphate, etc.)
Naming Ionic Compounds
• To name the ionic compound, simply state the cation, followed by the anion. Do not use the
word “ion” in the name of a neutral compound.
• Examples:
◦ NaCl is sodium chloride
◦ CaBr 2 is calcium bromide
◦ CuO is copper (II) oxide, or cupric oxide
◦ Cu 2 O is copper (I) oxide, or cuprous oxide
• Again, notice that we never use the prefixes used for binary nonmetals
(mono-, di-, tri-, etc.) when naming ionic compounds!
• This is because ionic compounds always form predictable ratios, as we have already seen.
◦ Na + and Cl - can only come together as NaCl.
◦ Cu 2+ and O 2- come together as CuO.
◦ Cu + and O 2- come together as Cu 2 O.
• Recall that the total charges of the cations and anions must neutralize.
Examples
Name each of the following ionic compounds:
AgCl
NH 4 NO 3
KNO 3
Hg(ClO 3 ) 2
PbCl 2
CdSO 4
ZnBr 2
SnCl 4
Oxoanions
• The oxoanions include many polyatomic anions which contain oxygen.
◦ OH - and O 2 2- are not considered oxoanions.
• So far, we have only encountered those oxoanions whose names end in “-ate.”
• Related oxoanions have different numbers of oxygen atoms but the same charge.
• Examples:
◦ SO 4 2- is the sulfate anion, SO 3 2- is the sulfite anion.
◦ ClO 3 - is the chlorate anion, ClO 4 - is the perchlorate anion.
6

Chemistry 120
Chapter Six Notes
• Let us consider the chlorate ion, which we know has formula ClO 3 - .
• It is very important that we know how many oxygen atoms are in the “-ate” ion.
• If the ion has one more oxygen in its formula than that of the “-ate” ion, add the prefix “per-” to
its name.
• So ClO 4 - is perchlorate, SO 5 2- would be persulfate, etc.
• Again, consider the chlorate ion, ClO 3 - .
• Taking away one oxygen from the “-ate” anion changes its ending from “-ate” to
“-ite.”
◦ Therefore, ClO 2 - is the chlorite ion, and SO 3 2- is the sulfite ion.
• Taking away two oxygens from the “-ate” anion changes its ending from “-ate” to
“-ite,” and you must add the prefix “hypo-”
◦ So, ClO - is the hypochlorite ion, and SO 2 2- is the hyposulfite ion.
Oxoanions Containing Hydrogen
• Some common polyatomic anions have an H + “within” their formula.
◦ In these ions, the H is always listed first, and the normal charge of the ion is changed by
+1.
◦ The name of the ion is changed by adding one of the following:
• The prefix “bi-”, or
• The word “hydrogen” before the name of the anion (sometimes you might see
“monohydrogen” instead.)
◦ For example, HCO 3 - is commonly called the “bicarbonate ion” or the “hydrogen carbonate
ion.”
◦ What would HSO 3 - be called
• Occasionally you may also see two hydrogen atoms added to the ion. This is rare, except in
the case of phosphate.
• This would increase the charge of the original anion by +2.
• The word “dihydrogen” is added to the name of the anion.
• So H 2 PO - 4 is the “dihydrogen phosphate” ion.
• KH 2 PO 4 is called “potassium dihydrogen phosphate.”
• NOTE: In all these examples, H + is not the cation; we should consider it an “attachment” of
the anion for naming purposes.
Examples
Name each of the following compounds:
KIO 4
LiH 2 PO 4
NaHCO 3
NaClO
Al 2 (SO 3 ) 3
Co(NO 2 ) 3
CuHPO 3
7

Chemistry 120
Chapter Six Notes
Oxoacids
• Oxoacids are compounds which contain exactly as many H + ions as are necessary to cancel
out the negative charge of an oxoanion.
◦ For example, SO 4 2- would need two H + ions to balance out the 2- charge of sulfate; the
oxoacid has the formula H 2 SO 4 .
◦ Note that the H + ions are
• the only cations in the formula, and
• always listed first in the formula.
Naming Oxoacids
• The naming rules for oxoacids are similar to those for oxoanions.
◦ The name of the oxyacid depends on how many oxygen atoms are in the corresponding
oxoanion.
• For those oxoanions which end in “-ate”, change the ending to “-ic acid.”
◦ So, CO 3
2-
is the carbonate ion
◦ H 2 CO 3 is carbonic acid
◦ HNO 3 is named _________________________________
◦ Slightly “weird” cases: H 2 SO 4 is sulfuric acid, and H 3 PO 4 is phosphoric acid.
• An oxoanion that begins with “per-” follows the same rules.
• Simply change the ending to “-ic acid.”
◦ HClO 4 is called perchloric acid.
◦ HIO 4 is called periodic acid.
• Any oxoanion that ends in “-ite” has its ending changed to “-ous acid”
◦ H 2 SO 3 is called sulfurous acid.
◦ HClO 2 is called chlorous acid.
• The acids of oxoanions which begin with “hypo-” likewise have their ending changed to “-ous
acid”.
◦ HBrO would be called hypobromous acid.
◦ HNO would be called hyponitrous acid.
Examples
Name the following oxoacids:
H 3 PO 4
HIO 2
HNO 3
H 2 CO 3
HNO 2 HIO
8

Chemistry 120
Chapter Six Notes
The Cyanide Ion
• One last random bit of information…
• The cyanide ion (CN - ) often behaves like a halogen, so naming rules for it are similar to those
of F - , Cl - , etc.
• HCN (g) is called “hydrogen cyanide.”
• HCN (aq) is called hydrocyanic acid.
Hydrates
• A hydrate is an ionic compound which has “trapped” a fixed number of water molecules within
its structure
• For example, the following compound is called copper (II) sulfate pentahydrate:
CuSO 4 · 5 H 2 O
◦ The formula tells us that there are five water molecules in the structure for every unit of
copper (II) sulfate
• The naming of other hydrates is similar:
◦ State the name of the salt
◦ End with a prefix indicating how many water molecules are present, followed by the word
“hydrate”
Name each of the following hydrates:
BaCl 2 · 2 H 2 O
MgSO 4 · 7 H 2 O
Final Examples
Name each of the following compounds.
BrO 4
NH 3
Cu(ClO 4 ) 2
N 2 O
H 2 SO 3
F 2
HF (aq)
HIO 3
LiHCO 3
9

Chemistry 120
Monatomic Ions I
This periodic table contains the formulas of what I consider to be the most important ions
which only form one charge. You'll notice a few important details:
(1) The charge on almost all of the metal ions on this list is the same as the group number.
Group
1A
(2) The charge on a nonmetal is simply its group number minus eight (i.e. oxide is
derived from oxygen, which is in group 6; 6 - 8= -2 )
(3) Zinc, cadmium, and silver ions are generally only encountered with the charge indicated.
(4) As a cation, hydrogen makes a 1+ ion; when hydrogen makes an ionic bond with a metal
H + /H - 2A the result is a "hydride" ion, which has charge 1-.
3A 4A 5A 6A 7A
(5) Boron and beryllium are excluded, as they generally form only covalent bonds.
(6) Many other elements form ions with two or more different charges; they have also been
Li + excluded.
C 4-
N 3- O 2-
carbide
nitride oxide
(rare)
Na + Mg 2+ 3B 4B 5B 6B 7B 8B 1B 2B Al 3+ P 3- S 2-
phosphide sulfide
K + Ca 2+ Zn 2+ Ga 3+ Se 2-
selenide
Rb + Sr 2+ Ag + Cd 2+ I -
F -
fluoride
Cl -
chloride
Br -
bromide
iodide
8A

Chemistry 120
Monatomic Ions II
Group
1A
This sheet lists the possible charges of many common elements which have more than one stable charge sta
You should know both the new name and old name for each (i.e. copper (I) ion and cuprous ion). Note that
the mercury (I) ion is a strange case; it is always found as a pair of atoms with a net charge of 2+ shared
between each atom in the pair.
8A
2A 3A 4A 5A 6A 7A
3B 4B 5B 6B 7B 8B 1B 2B
Cr 3+
Fe 3+ Co 3+
Cr 2+ Fe 2+
Co 2+ Cu 2+
Cu + Sn 4+
Sn 2+
Hg 2+
Hg 2
2+
Pb 4+
Pb 2+

Chemistry 120
Chapter Seven Notes
Chapter Seven
Formula Calculations
Introducing the Mole
• The dozen is a unit of quanity
◦ If I have a dozen atoms, I have 12 atoms by definition.
• The mole(mol) is a very important unit of quantity in chemistry. It is used to count large
numbers of atoms, molecules, and other submicroscopic pieces of matter.
• If you have 1 mole of something, you have 6.02 × 10 23 of it.
Examples
• How many eggs are in 5.5 dozen eggs
• How many atoms are in 5.5 mole of atoms
• How many hydrogen atoms are in one molecule of water How many oxygen atoms
• How many hydrogen atoms are in one mole of water How many oxygen atoms How many
atoms total
Practice
• If you have 18.45 mol of oxygen gas, how many individual molecules do you have How
many individual atoms
• If a sample of methane contains 7.35 × 10 25 atoms, how many moles of methane do you
have
1

Chemistry 120
Chapter Seven Notes
Atomic Weights and The Mole
• The atomic weights provided on the periodic table can be used in much more convenient
units than amu.
• The values on the periodic table can be interpreted as grams per mole of the atom.
◦ For example, 1 mol of calcium atoms has a mass of 40.08 g; 1 mol of neon atoms has a
mass of 20.18 g.
• Note that atomic weights are often also called molar masses.
Molecular Weights
• The mass of molecules can be calculated by adding up the atomic weights of the individual
atoms making up the molecule.
• For example, suppose we wanted to know the atomic weight of CO 2 .
◦ Every 1 mol of CO 2 contains_____mole of C atoms and_____moles of O atoms.
◦ Calculate the total mass of all of these atoms and add them up to get the molecular
weight of CO 2 .
◦ Aside: what is the mass of a single molecule of CO 2 in amu
• Note that the term formula weight is applied for those compounds which do not form true
molecules (ionic compounds like NaCl and CaO).
Examples
Determine the molecular weights of glucose (C 6 H 12 O 6 ) and acetic acid (CH 3 CO 2 H).
2

Chemistry 120
Chapter Seven Notes
Percent Composition
• One common technique used in specialized chemical laboratories is elemental analysis,
which can be used to determine the percent each element in a compound contributes to its
mass.
• Since this type of data is very common, it is important that we know how to interpret it and put
it to practical use.
• The percent composition is the percent of the total mass percent of each element in a
compound.
• To determine its value, we use the following formula for each element in the compound:
total mass of element
% composition =
× 100%
molar mass
Examples
What is the percent composition of each element in acetic acid
A 19.82 g chunk of an ore is analyzed and found to have a percent composition of 5.75% silver,
64.33% iron, and the remainder is silicon. What mass of each element is contained in a 4.87 kg
sample of this ore
3

Chemistry 120
Chapter Seven Notes
Percent compositions can also be determined for mixtures, such as alloys.
Suppose that 5.50 mols each of copper and zinc are blended with 2.43 mols of tin to make an
alloy. What is the percent composition of the alloy
A student determines the mass of a hydrate sample to be 6.873 g. After severely heating the
sample for 20 minutes to remove water, he reweighs the sample, finding the mass to be
5.276 g. What percent of the mass of the original hydrate is water
Molecular and Empirical Formulas
• The molecular formula of a compound tells us exactly how many atoms of each element are
contained in a compound.
• In contrast, the empirical formula only tells us the lowest whole-number ratio between each
element
• For example, glucose has molecular formula C 6 H 12 O 6 (each molecule of glucose contains 6
carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms)
• Glucose then has empirical formula CH 2 O (1 C to 2 H to 1 O is the lowest whole-number
ratio).
Empirical Formulas
• While the molecular formula ultimately is more useful for most applications, it is often more
difficult to determine than the empirical formula.
• When a sample is analyzed we can easily determine the percent composition, and from here
find the ratio of the number of atoms of each element.
• Information on the exact number of atoms in each molecule cannot be found in this way.
4

Chemistry 120
Chapter Seven Notes
Determining the Emprical Formula
• A simple series of steps can be used to determine the empirical formula of a compound:
◦ Find the mass of each element in the compound
◦ Convert the masses into moles of each element
◦ Express the moles of atoms as the smallest possible whole-number ratio
◦ Use the numbers from these ratios in the empirical formula for each element.
Examples
A sample of a compound is analyzed, and found to contain 1.61 g of phosphorus and 2.98 g of
fluorine. What is the empirical formula of this compound
The mass of a piece of iron is 1.62 g. The iron is exposed to oxygen and reacts to form a pure
oxide of iron, now with mass of 2.31 g. What is the empirical formula of this oxide
A compound is found to contain 20.0% carbon, 2.2% hydrogen, with the remainder of the mass
derived from chlorine. What is the empirical formula of this compound
5

Chemistry 120
Chapter Seven Notes
Molecular Formula
• Given the empirical formula and the molecular weight of a compound, it is possible to
determine the molecular formula.
• For example, suppose we have a compound with empirical formula CH.
◦ Its molecular formula could be CH, C 2 H 2 , C 3 H 3 , etc.
• Now, let’s see what the formula weight would be for each:
◦ CH = 13.02 g/mol
◦ C 2 H 2 = 26.04 g/mol
◦ C 3 H 3 = 39.06 g/mol
◦ Notice that each is a multiple of 13.02!
• To determine the molecular formula from this information
◦ Find the formula weight of the empirical formula
◦ Divide the molecular weight by this value (you should get a whole number).
◦ Multiply the subscripts in your empirical formula by this whole number.
• You now have your molecular formula.
Example
An unknown compound is determined to be 91.8% silicon and 8.2% hydrogen. The molar mass
of this compound is found to be 122 g/mol. Determine the molecular formula of the
compound.
6

Chemistry 120
Chapter Eight Notes
Chapter Eight
Chemical Reactions
What is a Chemical Reaction
• A chemical reaction involves the conversion of one or more substances into one or more
different substances.
• The substance(s) which we begin with are called the reactant(s)
• The substance(s) which we end with are called the product(s)
Writing Chemical Reactions
• Chemical reactions are written with the reactants to the left, the products to the right, and an
arrow between them to indicate the change.
◦ Occasionally symbols or values may be written over or under the arrow to indicate the
reaction conditions.
• An Example:
H 2 + O 2 H 2 O
Balancing Chemical Equations
• Consider the last reaction:
H 2 + O 2 H 2 O
• There is a problem with this equation!
• It indicates that we started with two oxygen atoms, but ended with one.
• What does this contradict
• We must balance chemical equations, which is to say that there must be equal numbers of
each type of atom on either side of a chemical reaction.
• To accomplish this, we put coefficients in front of the chemical formulas whose atom numbers
we wish to increase.
• Note that you may never change the subscripts already in place in a chemical formula!
◦ Why
• To balance chemical equations first count the number of each type of atom you have on both
sides of the reaction.
• Identify any lone elements (as opposed to compounds) in the formulas; you will balance
these last.
• From here, each equation requires its own logic; by trial and error, you should be able to
balance the equation.
• The only other real “tip” I can give you on this subject is that practice makes perfect!
1

Chemistry 120
Chapter Eight Notes
Examples
Balance each of the following chemical reactions:
Mg + O 2 MgO
HgO Hg + O 2
Na + MgCl 2 Mg + NaCl
Na 3 PO 4 + BaCl 2 Ba 3 (PO 4 ) 2 + NaCl
C 6 H 12 O 6 + O 2 CO 2 + H 2 O
• Write the balanced equation that corresponds to each statement
◦ When methane reacts with oxygen gas, carbon dioxide and water vapor are produced.
◦ Calcium metal reacts with ferric oxide, yielding calcium oxide and iron metal.
◦ Treatment of carbon monoxide gas with oxygen gas produces carbon dioxide.
Types of Chemical Reactions
• There are five main classification types of chemical reactions.
• Knowing how each of these reaction types “works” gives you the ability to predict what
chemical products will be formed in a given reaction in most cases.
• Types of Chemical Reactions
◦ Combustion
◦ Single Substitution (Displacement)
◦ Double Substitution (Displacement)
◦ Combination (or Synthesis)
◦ Decomposition
• Within these five classification types, even more specific reaction types can be identified;
these will be described as we come to them.
2

Chemistry 120
Chapter Eight Notes
Combustion Reactions
• In a combustion reaction, a chemical reacts with oxygen gas, forming various products.
• In this class, we will only consider the combustion of organic compounds containing C, H, and
sometimes O.
• In these reactions, the compound reacts with oxygen gas, producing carbon dioxide and
water vapor.
[organic compound] + O 2(g) CO 2(g) + H 2 O (g)
Examples of Combustion
• Combustion of benzene
2C 6 H 6(l) + 15O 2(g) 12CO 2(g) + 6H 2 O (g)
Combustion of formaldehyde
•
CH 2 O (l) + O 2(g) CO 2(g) + H 2 O (g)
• What is the balanced equation for the combustion of glycerol (C 3 H 8 O 3(l) )
Single Substitution Reactions
• In a single substitution reaction, an element replaces another element which is present as an
ion in a compound.
• There are two types of substitutions we should be aware of:
◦ substitution of one metal with another
◦ substitution of one halogen with another
• Format of the first type of single substitution reactions:
A + BX AX + B
Where A and B are metals/metal ions, and X is an anion.
• Here are some examples:
Mg + ZnCl 2 MgCl 2 + Zn
Zn + 2AgNO 3 Zn(NO 3 ) 2 + 2Ag
2Na + 2H 2 O 2NaOH + H 2
Activities of Metals
• We can use the activity of a metal to describe how readily it loses electron(s) in a reaction.
◦ The more active the metal, the more readily it loses its electrons.
• A more active metal will displace a less active metal in a single substitution reaction; the
reverse will not occur.
◦ In other words, an active metal will force the cation of a less active metal to take its
electrons away from it.
• We look to the activity series to see the relative activities of the metals.
◦ The activity series should not be memorized, but you should become familiar with trends
within it.
Displacing Hydrogen
• Notice that many metals can displace the H + from acids, changing it into H 2 .
◦ Notice that the charged hydrogen ion is transformed into the neutral hydrogen molecule.
• Some very active metals can even displace H + from water, leaving OH - behind.
◦ In these cases, think of water as HOH.
3

Chemistry 120
Chapter Eight Notes
When Does the Reaction “Go”
• So let’s consider one single substitution that works well…
Zn + 2AgNO 3 Zn(NO 3 ) 2 + 2Ag
• Since zinc is more active than silver (higher on the activity series), zinc will displace silver.
• Consider the reverse reaction…
Ag + Zn(NO 3 ) 2 no reaction
• Since silver is less active than zinc, it cannot displace it; therefore, no reaction can occur.
Oxidation & Reduction:
An Introduction
• Single substitution reactions involve an oxidation and a reduction.
◦ When a metal atom or ion is oxidized, it loses one or more electrons, resulting in an
increase in its charge.
• Example: Zn Zn 2+ + 2e -
◦ When a metal ion is reduced, it gains one or more electrons, causing its charge to
decrease.
• Example: Ag + + e - Ag
• An oxidation reaction is always accompanied by a reduction reaction.
• Redox is the common term for these reactions.
A Look at Some Redox Reactions
• Consider the following single substitution reaction:
3Mg + 2FeCl 3 3MgCl 2 + 2Fe
◦ What is being oxidized
◦ What is being reduced
• Try this reaction:
2Na + 2H 2 O 2NaOH + H 2
◦ What is being oxidized
◦ What is being reduced
Single Substitution Reactions:
Substitution of Halogens
• Anions derived from halogens (Group VII) can be displaced by a more active halogen.
• Activity of the halogens decreases down the group:
F 2 > Cl 2 > Br 2 > I 2
most active least active
• An example:
Cl 2 + 2NaI 2NaCl + I 2
• Will the reverse reaction proceed
4

Chemistry 120
Chapter Eight Notes
Examples
Predict the products of the following reactions. Write “no rxn.” if none is expected to occur.
Ba + CoBr 3
Ni + NaCl
I 2 + KF
Li + H 2 O
Ni + HNO 3
Double Substitution Reactions
• The general format for a double substitution reaction involves two ionic compounds trading
their anions:
AX + BY AY + BX
• AX and BY are both aqueous solutions, meaning both ionic compounds are dissolved in
water.
• These reactions will proceed only if at least one of the following is true:
◦ A solid precipitate is produced
◦ A covalent (molecular) compound is produced, such as H 2 O, H 2 , CO 2 , SO 3 , etc.
Examples of Double Substitution Reactions
AgNO 3(aq) + NaCl (aq) AgCl (s) + NaNO 3(aq)
3Ba(OH) 2(aq) + 2FeCl 3(aq) 3BaCl 2(aq) + 2Fe(OH) 3(s)
HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq)
Solubility of Compounds
• Solubility is best described as the degree to which a compound will dissolve in a solvent
(usually water).
◦ A compound that is soluble will dissolve to a significant extent.
◦ Compounds that are insoluble will not dissolve, remaining solid in solution.
• A reaction which produces an insoluble product can be described as a precipitation reaction;
the product “falls out” of the solution like rain precipitates.
5

Chemistry 120
Chapter Eight Notes
Solubility
• It is important to know whether or not some common chemicals are soluble or not in order to
predict how a reaction will proceed.
• You should know the following solubility rules (in water) by memory:
◦ All nitrates and acetates are soluble.
◦ All salts of Group I cations (Li + , Na + , etc.) and ammonium are soluble.
◦ All chlorides, bromides, and iodides are soluble, except those of Ag + , Pb 2+ , and Hg 2 2+ .
◦ All hydroxides are insoluble except those of Group I, NH 4 + , Ba 2+ , Sr 2+ , and Ca 2+ .
• Ca(OH) 2 is only slightly soluble.
Unstable Products of Reactions
• When produced in a reaction, some compounds will immediately decompose into other
products.
◦ Carbonic acid (H 2 CO 3 ) will decompose into CO 2(g) and H 2 O (l) .
◦ Ammonium hydroxide (NH 4 OH) will decompose into NH 3(aq) and H 2 O (l) .
◦ Sulfurous acid (H 2 SO 3 ) will decompose into SO 2(g) and H 2 O (l) .
• Notice that each produces water and a compound formed by the atoms left after water has
been removed from the starting formula.
• If you produce any of these three compounds in a reaction, cancel it out and replace it with
the decomposition products.
Electrolytes
• When an ionic compound dissolves in water, it dissociates (separates) into ions.
• A solution containing ions is capable of conducting an electric current.
• The term electrolytes is applied to these ions.
• A solution formed from a soluble ionic compound conducts the current very strongly; the
compound is called a strong electrolyte.
• A solution formed from an only slightly soluble ionic compound conducts the current weakly;
the compound is called a weak electrolyte.
• A solution derived from a non-ionic (i.e. covalent) compound will not conduct an electric
current; the compound is called a nonelectrolyte.
Acids and Bases
• You have heard the term acid applied to several compounds.
◦ All of these compounds mentioned so far contain the H + , called a proton.
• We will look more in depth at the definitions of acids and bases later, but for now we will use
the following:
◦ Acids are compounds which give up H + ions.
◦ Bases are compounds which accept H + ions.
• For now, we will limit our study of bases to some hydroxide-containing compounds
and NH 3 .
Strong Acids
• Strong acids are strong electrolytes, meaning they completely break up into ions in solution.
◦ So, HCl, which is a strong acid, breaks up into H + ions and Cl - ions in solution.
There are six strong acids you must know:
•
HCl HBr HI
HNO 3 H 2 SO 4 HClO 4
• Note that H 2 SO 4 dissociates into H + and HSO - 4 .
6

Chemistry 120
Chapter Eight Notes
Example: If you were to “look” into a beaker of the following solutions, what ions, solids, or
molecules would you expect to see (besides water)
a. HNO 3
b. HBr
c. HClO 4
Weak Acids
• Weak acids are weak electrolytes, meaning that relatively few acid molecules will break up
into separate ions.
• Some samples of weak acids are
H 3 PO 4 H 2 CO 3
HCH 3 CO 2
Strong and Weak Bases
• Similar definitions apply to strong bases.
• The strong bases are the Group I hydroxides (LiOH, NaOH, KOH, etc.) and some Group II
hydroxides: Ba(OH) 2 , Sr(OH) 2 , and Ca(OH) 2 .
• The weak bases include Mg(OH) 2 and NH 3(aq) (which can be treated as
NH 4 OH (aq) for now).
Neutralization Reactions
• Neutralization reactions are a subclass of double substitution reactions.
• The general form of this reaction is
acid + base water + salt
where the salt is any ionic compound.
• This can be further simplified as follows:
H + + OH - H 2 O
• Strong acids and strong bases completely neutralize each other.
• A strong acid will neutralize a weak base, and a strong base will neutralize a weak acid.
• The reactions of weak acids with weak bases is somewhat more complicated and will not be
considered in this class.
Examples
HBr (aq) + KOH (aq) H 2 O (l) + KBr (aq)
2HNO 3(aq) + Ba(OH) 2(aq) 2H 2 O (l) + Ba(NO 3 ) 2(aq)
7

Chemistry 120
Chapter Eight Notes
Write the balanced equations showing:
a) the neutralization of lithium hydroxide by perchloric acid.
b) the reaction of sulfuric acid with strontium hydroxide.
c) an ammonia solution is treated with hydrobromic acid.
Combination Reactions
• In a combination (or synthesis) reaction, two chemicals combine into one new chemical.
• These reactions have the general form
A + B C
• It will not always be possible to predict the products of combination reactions at this level of
preparation, so we will study a few general cases.
Oxide Formation
• Metals often react with oxygen to form a metal oxide.
Ex. 4Na + O 2 2Na 2 O
• Nonmetals often react with oxygen to form a nonmetal oxide.
Ex. C + O 2 CO 2
◦ It is often difficult to predict the products in this case, as CO was another possible oxide
you might have considered.
Reactions of Oxides
• Metal oxides often react with water to form metal hydroxides.
Ex. Na 2 O + H 2 O 2NaOH
• Nonmetal oxides often react with water to form oxyacids.
Ex. CO 2 + H 2 O H 2 CO 3
P 4 O 10 + 6H 2 O 4H 3 PO 4
SO 3 + H 2 O H 2 SO 4
N 2 O 5 + H 2 O 2HNO 3
• Metal oxides and nonmetal oxides often combine to form a salt.
Na 2 O + CO 2 Na 2 CO 3
CaO + SO 3 CaSO 4
Decomposition Reactions
• Decomposition reactions are simply the reverse of combination reactions.
• Their general form is
Z X + Y
• Like combination reactions, predicting the products of these reactions is often difficult.
• For now, simply decompose a given compound into two products which could have produced
it from a method we discussed earlier.
◦ This is very oversimplified, and not always correct, but it is the best we can do at this
level.
8

Chemistry 120
Chapter Eight Notes
Examples
2HgO 2Hg + O 2
Na 2 CO 3 Na 2 O + CO 2
Li 2 SO 4 Li 2 O + SO 3
Heat in Chemical Reactions
•Virtually every chemical reaction results in either a gain or loss of heat to the reaction system
•A reaction which produces more heat than it consumes is called an exothermic reaction
•A reaction which consumes more heat than it produces is called an endothermic reaction
•The chemical reaction written below describes an exothermic reaction:
CH 4(g) + 2 O 2(g) CO 2(g) + 2 H 2 O (g) + 890 Kj
•For an endothermic reaction, the heat would appear with the reactants in the chemical
equation:
PCl 5 + 92.9 kJ PCl 3 + Cl 2
•Although this method is convenient for indicating heat changes, it often leads to confusion in
balancing equations
◦A better method for indicating energy changes will be presented in Chem 130
Heat and Chemical Bonds
•Energy is always consumed in the breaking of chemical bonds.
•The opposite is also true: energy is always released when new chemical bonds are formed.
•Ultimately, whether a reaction is exothermic or endothermic therefore depends on the number and type of
bonds being broken and formed in a chemical reaction
Activation Energy and Energy Diagrams
•In order for a chemical reaction to begin, some amount of energy specific to that reaction must be
supplied to start the reaction process
•The amount of this initial investment of energy is called the activation energy
•The activation energy, as well as the energy changes in chemical reactions, may be illustrated by energy
diagrams
9

Activity Series of Metals
Most Li Displaces hydrogen from
Active Rb cold water, steam, and acids
K
Cs
Ba
Sr
Ca
Na
Mg Displaces hydrogen from
Al steam and acids
Mn
Zn
Cr
Fe
Cd Displaces hydrogen from
Co acids
Ni
Sn
Pb
H 2
Cu Cannot displace hydrogen from
Hg cold water, steam, or acids
Ag
Least Pt
Active Au
Solubilities of Selected Ionic Compounds in Water
In general, all ionic compounds containing ammonium (NH 4 + ) and Group IA cations (Li + , Na + , K + , etc.) are
soluble.
nitrates
acetates
chlorates
chlorides
bromides
iodides
sulfates
chromates
oxides
hydroxides
sulfides
carbonates
phosphates
sulfites
all are soluble
all are soluble except Ag + , Pb 2+ , & Hg 2
2+
all are soluble except Ag + , Pb 2+ , & Hg 2 2+ , Hg 2+ ,
Ba 2+ , Sr 2+ , & Ca 2+
note: CaCrO 4 is soluble, CuCrO 4 is not
none soluble except Group IA, NH 4 + ,
Ba 2+ , Sr 2+ , & Ca 2+
none soluble except Group IA, NH 4 + , and
Group IIA
none soluble except Group IA & NH 4
+

Activity Series of Metals—Test Version
Most Li Displaces hydrogen from
Active Rb cold water, steam, and acids
K
Cs
Ba
Sr
Ca
Na
Mg Displaces hydrogen from
Al steam and acids
Mn
Zn
Cr
Fe
Cd Displaces hydrogen from
Co acids
Ni
Sn
Pb
H 2
Cu Cannot displace hydrogen from
Hg cold water, steam, or acids
Ag
Least Pt
Active Au
Solubilities of Selected Ionic Compounds in Water
Note that the solubility rules for those compounds which you were to have memorized have been deleted
from the table.
chlorates
sulfates
chromates
sulfides
carbonates
phosphates
sulfites
all are soluble
all are soluble except Ag + , Pb 2+ , & Hg 2 2+ , Hg 2+ ,
Ba 2+ , Sr 2+ , & Ca 2+
note: CaCrO 4 is soluble, CuCrO 4 is not
none soluble except Group IA, NH 4 + , and
Group IIA
none soluble except Group IA & NH 4
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Chemistry 120
Net Ionic Equation Notes
Net Ionic Equations
Revisiting “aqueous”
• The term aqueous is mainly a visual one; when you add a chemical to water and it appears to
dissolve, we say an aqueous solution has been formed.
• However, we must be more specific:
◦ For ionic compounds, including strong acids, aqueous means that the compound has
dissociated into ions.
◦ For weak acids and other covalent compounds, aqueous means that the compound has
dissolved, but not been broken down into smaller parts. It is merely “surrounded” by
water molecules.
Looking Back
• Consider the following reaction:
NaCl (aq) +KNO 3(aq)
• Since both the products of this reaction are aqueous, and since no molecular compound is
formed, we say that there is no reaction.
• However, since we did bring these two solutions together, why is it incorrect to say that the
products are KCl (aq) and NaNO 3(aq)
• Earlier, we defined a chemical change as one that involves transforming one or more
chemical compounds into one or more different products.
• By combining two solutions which do not make an insoluble product or a molecular
compound, we fail to create a new product, as we shall now prove.
Total Ionic Equations
• Total molecular equations show all species in solution for a given reaction.
◦ The term species refers to the specific ions and/or compounds in the solution.
• Examples of species:
For NaCl (aq) , Na + (aq) and Cl - (aq).
For HNO 3(aq) , H + (aq) and NO 3
-(aq).
For H 3 PO 4(aq) , it remains H 3 PO 4(aq) .
• Some simple guidelines for writing down the species.
◦ Strong electrolytes, including strong acids and bases, dissociate into their ions.
◦ Weak electrolytes, including weak acids and bases, generally do not dissociate, so the
species is just the original compound.
◦ Nonelectrolytes do not dissociate, so the species is again the same as the original
compound.
Total Ionic Equations
Example:
NaCl (aq) + AgNO 3(aq) AgCl (s) + NaNO 3(aq)
Total Molecular Equation:
Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3
-
(aq) AgCl (s) + Na + (aq) + NO 3
-
(aq)
1

Chemistry 120
Net Ionic Equation Notes
Spectator Ions
• A spectator of sports is someone who watches the game from the sidelines, but does not
participate.
• Similarly, in chemical reactions, spectator ions “hang out” in a solution but do not actively
participate in the reaction itself.
◦ In other words, any ion which is both on the reactants and products side of a reaction is a
spectator ion, for it has not undergone a chemical change.
• The ions’ main purpose is to maintain constant charge in the solution.
Net Ionic Equations
• Net ionic equations only show those chemicals which participate in the reaction. Spectator
ions are not included.
• To write a net ionic equation, first write down the total ionic equation.
• Then, cancel anything which appears identically on both sides of the reaction.
• Consider the total ionic equation
Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3
-
(aq) AgCl (s) + Na + (aq) + NO 3
-(aq)
Now, factor out the spectator ions
•
Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3
-
(aq) AgCl (s) + Na + (aq) + NO 3
-(aq)
• The net ionic equation is left over.
Ag + (aq) + Cl - (aq) AgCl (s)
Examples
Write chemical equations, total ionic equations, and net ionic equations for the following:
a. a solution of barium chloride reacts with a sodium sulfate solution.
b. sodium hydroxide is neutralized by hydrochloric acid.
c. phosphoric acid is reacted with potassium hydroxide.
d. solutions of sodium nitrate and ammonium chloride are brought together.
e. a piece of sodium metal is dropped into cold water.
2

Type of Substance Type of Electrolyte Ionizes in Water Examples
How to Write in Net
Ionic Equations
NaCl (aq)
Na + (aq) + Cl - (aq)
soluble ionic
compounds
strong
completely
KNO 3(aq)
K + (aq) + NO 3
-
(aq)
(NH 4 ) 2 SO 4(aq)
2 NH 4
+
(aq) + SO 4
2-
(aq)
AgCl (s)
AgCl (s)
insoluble ionic
compounds
nonelectrolyte
PbI 2(s)
PbI 2(s)
covalent/molecular
compounds
(not all compounds in these
categories are nonelectrolytes,
but you should
not worry about exceptions)
not at all
BaSO 4(s)
H 2 O (l)
C 6 H 12 O 6(aq)
BaSO 4(s)
H 2 O (l)
C 6 H 12 O 6(aq)
H 2(g)
H 2(g)
strong acids
and
strong bases
strong
completely
HCl (aq)
HNO 3(aq)
NaOH (aq)
H + (aq) + Cl - (aq)
H + (aq) + NO 3
-
(aq)
Na + (aq) + OH - (aq)
weak acids
and
weak bases
weak
only slightly
H 3 PO 4(aq)
HC 2 H 3 O 2(aq)
Mg(OH) 2(s)
H 3 PO 4(aq)
HC 2 H 3 O 2(aq)
Mg(OH) 2(s)

Net Ionic Equations
Write (1) full formula, (2) total ionic, and (3) net ionic equations for each of the following chemical
combinations. Write “no rxn.” To the right of the arrow if you expect no reaction to occur.
1. A piece of potassium is dropped in water.
(1) Full:
(2) Total
(3) Net Ionic
2. Aqueous solutions of iron (III) sulfate and ammonia are mixed.
(1) Full:
(2) Total
(3) Net Ionic
3. Silver is added to a cupric bromide solution.
(1) Full:
(2) Total
(3) Net Ionic
4. A sodium fluoride solution is mixed with nitric acid.
(1) Full:
(2) Total
(3) Net Ionic

5. Aqueous solutions of potassium phosphate and cobalt (II) chloride are mixed.
(1) Full:
(2) Total
(3) Net Ionic
6. Aqueous ammonia is mixed with hydoiodic acid.
(1) Full:
(2) Total
(3) Net Ionic
7. Phosphoric acid is mixed with an aqueous solution of lithium hydroxide.
(1) Full:
(2) Total
(3) Net Ionic
8. Carbon dioxide is passed into aqueous barium hydroxide.
(1) Full:
(2) Total
(3) Net Ionic
9. Hydrobromic acid is mixed with a sodium carbonate solution.
(1) Full:
(2) Total
(3) Net Ionic

10. Nitric acid is mixed with aqueous potassium bicarbonate.
(1) Full:
(2) Total
(3) Net Ionic
11. Rubidium oxide is added to water.
(1) Full:
(2) Total
(3) Net Ionic
12. Aqueous solutions of lithium acetate and tin (II) chloride are mixed.
(1) Full:
(2) Total
(3) Net Ionic
13. Hydrochloric acid is added to a lithium sulfite solution.
(1) Full:
(2) Total
(3) Net Ionic

Chemistry 120
Chapter Nine Notes
Chapter Nine: Stoichiometry
Conservation
• Recall that, in any chemical reaction, that matter is always conserved.
• Consider the reaction of zinc with sulfur:
Zn (s) + S (s) ZnS (s)
• From this reaction we can imply that, assuming the reaction works perfectly, every one atom
of zinc reacts with one atom of sulfur to produce one unit of zinc sulfide.
• More conveniently, we can say that one mole of zinc atoms (1 mol Zn) reacts with one mole
of sulfur atoms (1 mol S), producing one mole of zinc sulfide (1 mol ZnS)
Common Misconceptions
• Is it true that if I start a reaction with 5.0 g of zinc and 5.0 g of sulfur that I will end the reaction
with 10.0 g of matter
• Is it true that if I start a reaction with 5.0 g of zinc and 5.0 g of sulfur that I will end the reaction
with 10.0 g of zinc sulfide
• In predicting the quantities of products created during a reaction, it is ultimately the number of
atoms/molecules reacting that we must consider, not the mass.
• The most convenient way to do this is to consider the moles of reactants participating in a
reaction.
Coefficients of Reactions
• The coefficients in chemical equations are useful in this context because they tell us how
many moles of reactants are required to produce a certain number of moles of products.
• We can say that one mole of zinc reacts with one mole of sulfur, producing 1 mole of zinc
sulfide. The three are equivalent for the purpose of this reaction.
1 mol Zn1mol S1 mol ZnS
• We can convert between moles of reactants and products using dimensional analysis, the
same technique we used for unit conversions.
• To convert between moles of zinc and moles of zinc sulfide, we use the following ratio:
Basic Problems
Let’s consider the following reaction:
2H 2(g) + O 2(g) 2H 2 O (l)
To produce 2 mol of water we need to combine_____mol H 2 with_____mol O 2 .
How many moles of each reactant are required to produce 10 moles of water
1

Chemistry 120
Chapter Nine Notes
How many moles of water can be produced at most from 13 moles of hydrogen How many
moles of oxygen would I need to accomplish this
In the reaction of aqueous solutions of hydrochloric acid and barium hydroxide:
a. How many moles of each reactant do you need to produce 5.75 mol water
b. What is the maximum number of moles of water you can produce with 18.0 mol HCl How
many moles of barium hydroxide would you need to accomplish this
2

Chemistry 120
Chapter Nine Notes
Using the Mass-Mole Relationship
• Recall that the molar mass relates the mass of a compound (or element) to the number of
moles.
• Therefore, if we know the mass of a given reactant, we can easily determine the number of
moles of the reactant, and from there the number of moles of product produced.
• Remember, you must use the mole-mole relationship to carry out a conversion; the mass
alone is not sufficient!
Mass-Mole Problems
Consider the reaction of sodium metal with chlorine to make sodium chloride.
a. How many moles of sodium are contained in 15.50 g of sodium
b. Assuming we have excess (that is, more than enough) chlorine, how many moles of
sodium chloride could be produced
c. How many grams of sodium chloride is this
When 39.75 g of sodium is reacted with excess chlorine, how many grams of sodium chloride can
be produced
3

Chemistry 120
Chapter Nine Notes
When 146.7 g of chlorine is reacted with excess sodium, how many grams of sodium chloride can
be produced
How many grams of each reactant do you need to produce 100.00 g of sodium chloride
Another Example:
Ethanol has molecular formula CH 3 CH 2 OH.
a. How many grams of carbon dioxide and water can be produced if 55.0 mL of ethanol is
combusted
Note: The density of ethanol is 0.789 g/mL.
4

Chemistry 120
Chapter Nine Notes
b. How many oxygen molecules are required to react with this amount of ethanol
The Ice Cream Sundae Problem
• Suppose we are going to make ice cream sundaes.
• We decide that each of our sundaes must contain
◦ 2 scoops of ice cream
◦ 50 mL of chocolate syrup
◦ 1 cherry
• How many sundaes can I make if I have 8 scoops of ice cream, 13 cherries, and 300 mL of
chocolate syrup available
5

Chemistry 120
Chapter Nine Notes
Limiting Reactants
• The problem we had with our ice cream sundaes is the same type of problem we face in
chemical reactions.
• In virtually all reactions, we will have a disproportionate amount of reactants.
• The reactant which will produce less product is called the limiting reactant; in all cases, it
predicts the amount of product formed.
• The other reactants are said to be in excess, meaning there is more than enough of it.
• The key to identifying a limiting reagent problem is that it will give you the amounts being
reacted of more than one reactant.
• Suppose you with to synthesize zinc sulfide.
You combine 50. grams of zinc with 50. grams of sulfur.
• What mass of zinc sulfide is predicted to be produced from the zinc
• What mass of zinc sulfide is predicted to be produced from the sulfur
• Which reactant is the limiting reactant
What mass of zinc sulfide is produced by the reaction of 75.8 g of zinc with 60.5 g of sulfur
6

Chemistry 120
Chapter Nine Notes
Using the Tabulation Method
• The tabulation method is an alternative method for solving some stoichiometry problems
• While most problems do not require using this method, it is essential for more complex
problems, especially those involving equilibrium calculations
• This method is most helpful when you need to know the amounts of each chemical present at
the end of the reaction
Use the tabulation method to solve this problem:
3.0 mols of N 2(g) react with 8.0 mols of H 2(g) to produce NH 3(g) . If the reaction goes to
completion (that is, until you run out of the limiting reactant), how many moles of each
substance will be present in the flask at the end of the reaction
Consider the hydrogenation of propene(C 3 H 6 ):
C 3 H 6 + H 2 C 3 H 8
78.24 g of propene is treated with 25.0 g of hydrogen.
a. What mass of propane (C 3 H 8 ) is produced
7

Chemistry 120
Chapter Nine Notes
b. Which reactant was in excess, and what mass of it remains after the reaction is
completed
Percent Yield
• So far, we have assumed that all reactions will produce exactly the amount of products that
our calculations will predict.
• The theoretical yield is the amount of a product you expect to get at the end of a reaction.
• In reality, we never actually get 100% of this amount.
◦ Why might this be so
• The amount of a product that is really produced at the end of a reaction is called the actual
yield.
• The percent yield gives us an idea of how much of a product we actually produced compared
to the amount that should have been produced.
actual yield
% yield =
× 100%
theoretical yield
• In general, the closer the percent yield is to 100% percent, the better the reaction
Example
• 8.0 g of hydrogen is combined with excess oxygen, producing 58.5 g of water. What is the
percent yield of this reaction
8

Chemistry 120
Chapter Nine Notes
Example
Suppose that you wish to produce 15.00 grams of barium sulfate by reacting excess sulfuric acid
with barium hydroxide. If this reaction is known to usually give an 87.5% yield, how many
grams of barium hydroxide should you use
Stoichiometry and Net Ionic Equations
• It is often more practical to use net ionic equations in stoichiometry calculations, especially for
problems dealing with solutions
• There is no significant difference in solving problems using these equations than with the
methods used so far
9

Chemistry 120
Chapter Nine Notes
Example
12.0 grams of strontium chloride is dissolved in approximately 500 mL of water. To this is added
5.00 grams of silver nitrate. Determine
a. The total ionic equation and net ionic equation describing this process.
b. The mass of solid precipitate theoretically produced, and
c. The moles of each ion remaining in solution at the end of the reaction.
10

Chemistry 120
Molarity Notes
Molarity
Concentration and Molarity
• The concentration of a solution is a measurement of how much of a solute it contains (the
solute is the substance dissolved within it.)
• The most useful term for concentration relates the number of moles of solute in a liter of
solution; it is called the molarity, and abbreviated “M”:
mol solute
molarity(M ) =
L solution
• We often use square brackets around a compound or ion to indicate molarity
• For example [Na + ] means “molarity of sodium ion”
• For example, to prepare a 1.0 M NaCl solution, we take 1.0 mol of NaCl (58.44 g) and add
enough water to it to bring the total volume of the solution to 1.0 L.
• We cannot simply say “add 1.0 L of water,” because the presence of the salt will effect the
final volume of the solution.
Example
Explain how to prepare 2.00 L of a 0.550 M solution of potassium chloride.
Mole-Volume Problems
• Since the molarity of a solution relates the volume to the number of moles of a solute, we can
use it to help us calculate the amount of product formed in a reaction.
• First, we must convert the volume of a solution to moles of solute, using the molarity.
• From there, we can use the mole ratios of the chemical equation to figure out how much
product is produced.
1

Chemistry 120
Molarity Notes
Examples
25.00 mL of a 0.45 M AgNO 3 solution is treated with excess NaCl solution. What solid precipitate
is formed, and what should its mass be in grams
What volume of 0.224 M Ba(OH) 2 solution is required to completely neutralize 15.00 mL of a
0.855 M HNO 3 solution
2

Chemistry 120
Molarity Notes
25.00 mL of a 0.425 M potassium iodide solution is added to about 90 mL of water. To this is
added 25.00 mL of a 0.724 M lead(II) nitrate solution. What precipitate is formed, and what is
its mass
Millimoles
• Since we often carry out volume measurements in the lab in milliliters, it is often convenient to
use millimoles (mmol) rather than moles in calculations
• A millimole is defined in the same way other metric units are
◦ 1 mmol = 10 -3 mol, or equivalently
◦ 1000 mmol = 1 mol
• We can define molarity as
mol solute mmol solute
molarity(M ) =
=
L solution mL solution
and use the form which is most convenient
Example
Let’s redo an earlier problem, using millimoles instead of moles in the calculation:
What volume of 0.224 M Ba(OH) 2 solution is required to completely neutralize 15.00 mL of a
0.855 M HNO 3 solution
3

Chemistry 120
Molarity Notes
Titrations
• A titration is usually carried out to determine the molarity of an unknown solution.
• A specific amount of a compound whose concentration is well known, called the primary
standard is measured out.
• An indicator, a compound with a different color under different chemical environments, is
added to the primary standard.
• The unknown solution is added until the indicator changes color, signaling the end of the
titration (called the end point).
• From the collected data it is possible to determine the molarity of the unknown solution.
• At the end point, we assume that there are “equivalent” amounts of each chemical present
◦ This is not precisely true, but this definition is adequate for now
• By equivalent, we mean that the ratio of the moles of each reactant added equals the ratio in
the balanced equation
Examples
HCl + NaOH H 2 O + NaCl
Ratio of acid to base: 1 to 1, so we need 1
equivalent of acid for every 1
equivalent of base
H 2 SO 4 + 2 NaOH 2 H 2 O + Na 2 SO 4
Ratio of acid to base: 1 to 2, so we need 1 equivalent of acid for every 2
equivalents of base
Examples
10.00 mL of an HCl solution is pipetted into a beaker containing about 25 mL of deionized water.
This solution is then titrated with 0.4575 M NaOH. It is found that 17.85 mL of base is
required to reach the end point. What is the molarity of the acid solution
4

Chemistry 120
Molarity Notes
4.2554 grams of solid oxalic acid (H 2 C 2 O 4 ) is dissolved in about 50 mL of deionized water. It is
found that 37.55 mL of a KOH solution is required to titrate this acid to the end point. What is
the concentration of the base solution
Ion Concentration
• For many chemical calculations, it is essential that we know the concentration of each ion in
the solution
• For example, consider a 1.00 M CuCl 2 solution
• What is the molarity of Cu 2+ in this solution
• What is the molarity of Cl -
Examples
10.00 mL of 2.45 M NaCl solution is combined with 10.00 mL of 1.50 M BaCl 2 solution. What is
the concentration of each ion in solution after the two have been combined
5

Chemistry 120
Molarity Notes
How could we prepare 500. mL of a solution with [Ni 2+ ] = 0.100 M We will use nickel (II) nitrate
as the source of the nickel ion. (Why is it not possible to add only Ni 2+ ion to the solution
without adding any other ion)
15.00 mL of a 0.1555 M HCl solution is combined with 10.00 mL of a 0.1250 M Sr(OH) 2 solution.
What is the concentration of each ion in solution after the reaction has occured
6

Chemistry 120
Chapter Twelve Notes
Chapter Twelve: Gases
Pressure
• An important parameter which must be taken into account when studying gases is pressure.
• Pressure is the amount of force applied to an area:
force
pressure =
area
• Consider laying on a bed of nails. Do you want there to be many nails or few nails in the
bed
Pressures of Gases
• The pressure exerted by a gas in a sealed container results from collisions of many individual
gas particles with the walls of its container.
◦ When we say a gas is at high pressure, its individual particles are colliding frequently with
the sides of its container.
• Later we will see that many other parameters can have an effect on the pressure of a gas
sample, such as the volume of the enclosing container.
Air Pressure
• The air in the atmosphere exerts a substantial amount of pressure on everything on the
surface of the earth.
• Since air is composed of matter (nitrogen, oxygen, etc.), it has mass.
• Air reaches up to at least 20 miles above the surface of Earth.
• The weight of the air pushes down on everything on the surface, including us.
• In a similar fashion, the weight of ocean water exerts tremendous pressures at the lower
depths of the oceans.
Measuring Pressure
• A device used for measuring pressure is called a manometer; a device specifically designed
to measure air pressure is called a barometer.
• To build a simple barometer, we completely fill a thin cylinder with mercury
◦ Mercury is very dense and is therefore ideal for this application
• Then, we partially fill a deep dish with mercury and invert the mercury tube in this pool of
mercury.
• Some of the mercury will descend out of the tube as gravity pulls it down.
• However, some of the mercury will stay in the tube, as the air pressure keeps this mercury
supported.
◦ The weight of the air on the surface of the mercury pool opposes the force of gravity as it
tries to pull the mercury down out of the column.
• On an “average” day, the height of the mercury column will be close to 760 mm.
• What would the height of the column be on average at a higher altitude: greater than 760
mm, or less
• The “standard” air pressure is set at this value of 760 mm, as we shall shortly see.
1

Chemistry 120
Chapter Twelve Notes
Measuring Pressure
• Measuring the pressure of the gas in a container involves a similar principal.
• One method of measuring this pressure involves attaching the sample to a manometer, which
has a U-shaped tube partially filled with mercury.
• The gas exerts pressure on its side of the U-tube, and the atmosphere exerts pressure on the
other open end.
• The difference in the heights of the columns can tell use the difference in the pressures of the
gases.
• The side of the U-tube with the lower height indicates that that side is at a higher pressure,
since it is ultimately pushing back the force of the other gas.
• To determine the pressure of the gas, the difference in the height is measured and
◦ added to the atmospheric pressure value if the height is lower on the gas sample side.
◦ subtracted from the atmospheric pressure value if the height is higher on the gas sample
side.
Units of Pressure
• Unfortunately, many different units are used to measure pressure.
• The units we will employ most often include
◦ Atmospheres (atm)
• Though not the SI unit of pressure, this is a very common unit. The value of the air
pressure at sea level on an average day is close to 1 atm.
◦ Millimeters of Mercury (mm Hg)
• This unit is derived from the way one takes a measurement with a barometer or
manometer. Recall that on an average day, the height of a mercury column in a
barometer is about 760 mm Hg.
• 1 atm = 760 mm Hg (approximately, though treat this as exact).
◦ Torr (torr)
• This unit is essentially the same as the mm Hg, so
1 atm = 760 torr (exactly).
• You must know these units and conversions.
• Other units of pressure you may encounter, but need not memorize the conversion for,
include
◦ Pascals (Pa) (the SI unit of pressure) and kilopascals (kPa)
◦ Pounds per square inch (psi)
◦ Bars (bar) and millibars (mbar)
1 atm = 101,325 Pa = 14.7 psi = 1013 mbar
2

Chemistry 120
Chapter Twelve Notes
Kinetic-Molecular Theory
• Recall that our only previous description of gases stated that gases completely fill and take
the shape of their containers.
• The Kinetic-Molecular Theory goes much farther in describing how gases ideally behave.
• There are five principle statements that comprise it, which should be understood thoroughly
as well as memorized.
◦ Note that some textbooks give slight modifications of these
The Statements of the Kinetic-Molecular Theory
• The particles of a gas are so small compared to the distances between them that the volume
of the individual gas particles is considered negligible (zero).
• Gas particles are in constant, rapid, and random motion.
• Gas particles exert neither attractive nor repulsive forces on each other.
• The average kinetic energy of a collection of gas particles is directly proportional to the
temperature (in Kelvin) of the gas.
◦ From this we can state that all gases at the same temperature have the same average
kinetic energy.
• Collisions between gas particles are perfectly elastic
◦ By this, we mean that the gas molecules do not lose energy when they collide
The Gas Laws
• In considering gases, there are three important measured values which we will often
consider.
◦ Pressure (P)
◦ Volume (V)
◦ Temperature (T)
• Note that the units of temperature will always be expressed in Kelvin in the gas laws.
• The gas laws describe the mathematical relationship between these parameters
Boyle’s Law
• Suppose we have a gas in a sealed piston (a container whose volume can change), and we
know its pressure, volume, and temperature.
• Keeping the temperature the same, we decrease the volume of the container.
• What will happen to the pressure
◦ Hint: Will the gas molecules collide with the sides of the container more or less
frequently
• Pressure and volume are inversely proportional.
◦ That is, if one increases, the other must decrease.
◦ Remember that the temperature must remain constant for this to be true.
• We can use this fact to derive Boyle’s Law:
P 1 V 1 = P 2 V 2 (at constant T)
where the 1 stands for the beginning values and the 2 the final values.
• Derivation:
3

Chemistry 120
Chapter Twelve Notes
Example
A gas occupies 29.5 L, and exerts a pressure of 640. mm Hg. If the container’s volume is
expanded to 40.0 L (keeping the temperature constant), what will be the pressure of the gas
in mm Hg and in atm
Charles’ Law
• Suppose we have a gas in an expandable container, and, without changing its pressure, we
raise its temperature.
• How will the behavior of the gas molecules change
• In order to keep the pressure constant, what will happen to the volume of the gas
• Volume and temperature are directly proportional.
◦ That is, if one increases the other increases.
◦ Pressure must remain constant for the process.
• Mathematically, this is stated as
V
1 =
T
1
V
T
2
2
(constant P)
Example
A gas occupies 35.8 L at 1.13 atm and 42.3 °C. What would be the temperature of this gas, in
Celcius, if the volume were decreased to 25.0 L and the pressure was left unchanged
4

Chemistry 120
Chapter Twelve Notes
Gay-Lussac’s Law
• Suppose a gas is kept in a rigid container whose volume cannot change.
• If the gas is heated, how will the pressure be effected
• Pressure and temperature are directly proportional, assuming the volume remains
unchanged. Mathematically, this is stated:
T
1
1
P
T
2
2
(constant V)
P =
2
The Combined Gas Law
• The three previous laws can all be combined into one convenient form, the combined gas
law.
P1
V
T
1
1
P2
V
=
T
2
• You may use this law in place of the other three if you wish, since the constant values will
always divide out.
Example
72.6 L of a gas at 35.7 °C and 0.872 atm is heated to 50.0 °C and allowed to expand to 87.5 L.
What is the pressure of the gas under these new conditions
Avogadro’s Hypothesis
• Amadeo Avogadro is credited with stating one of the most fundamental statements describing
gases:
“Equal amounts (moles) of gases occupy equal volumes under the same conditions”
• This greatly simplifies our treatment of gases, as we can conclude that, regardless of what
composes the gas, if we know how many moles of it we have and its conditions, we can find
its volume.
5

Chemistry 120
Chapter Twelve Notes
Avogadro’s Law
• Mathematically we can use Avogadro’s hypothesis to derive the mathematical relationship
V
1 =
n1
V
n
2
2
(constant P, T)
• n is the number of moles of the gas, regardless of what the gas actually is.
The Ideal Gas Law
• All the previous laws allow us to compare a gas under one set of conditions with its values
under other conditions.
• The Ideal Gas Law summarizes the previous laws in a fundamentally different way: its power
is that it shows how the four parameters (P, V, T, and n) can be used to determine one
another.
• The Ideal Gas Law is, mathematically,
PV = nRT
where R is a constant, called the universal gas constant.
L ⋅atm
R = 0.0821
mol⋅K
• If you know any three parameters from the ideal gas law, the fourth can be determined by
substitution into this equation.
Examples
What volume is one mole of a gas expected to occupy at exactly 1 atm and exactly 0 °C
Note: These pressures and temperatures are called STP (standard temperature and
pressure).
If a balloon containing methane has a volume of 7.25 L at 754 mm Hg and 296.4 K, then how
many moles of methane does the balloon contain How many grams is this
6

Chemistry 120
Chapter Twelve Notes
A 1.45 g piece of dry ice (solid CO 2 ) is put in a balloon and warmed to 298 K, causing all the CO 2
to sublimate (turn to gas). The volume of the balloon after the sublimation is completed is
measured to be 3.29 L. What is the pressure exerted by the gas inside the balloon
What is the density of pure carbon dioxide gas at 1.05 atm and 300.0 K Report your answer in
grams per liter.
Dalton’s Law of Partial Pressures
• In a mixture of gases (like air) each substance in the mixture exerts its own indivual pressure,
called the partial pressure of that gas.
• The total pressure of the mixture of these gases is simply the sum of these partial pressures.
◦ So, air pressure is derived from the partial pressures of all the gases making up the air,
including N 2 , O 2 , Ar, H 2 O, CO 2 , and others.
• Mathematically, this is stated as
P total = P gas a + P gas b + P gas c + …
◦ There are as many terms on the right side of the equation as there are gases in the
mixture.
7

Chemistry 120
Chapter Twelve Notes
Example
Argon, neon, and helium gases are all combined in a 40.0 L container. The total pressure of the
three gases is measured to be 1.45 atm. Argon and neon each exert a partial pressure of
0.65 atm. What is the partial pressure of helium in this mixture
Collecting a Gas Over Water
• One common way to collect a gas produced by an experiment is over water.
• From the reaction vessel, the gas enters a hose which leads to a container filled with water.
• The gas then bubbles up through the water into a tube filled with water that penetrates the
waters surface.
• The gas displaces the water in the tube, and is collected there.
• However the gas is “wet”, meaning it contains some water vapor.
• To find the pressure of the “dry” (i.e. without water) sample, we must subtract out the
pressure of the water vapor mixed with it.
◦ The pressure exerted by the water pressure can be found in a table, so long as we know
the temperature of the water.
• If the level of the water in the tube is the same as the level of the water in the container the
tube is in, the following relationship exists:
P air = P gas + P water vapor
Example
A sample of argon is collected over water at 23 °C. After collecting the gas, the height of the gas
in the collecting tube is lowered to that of the water container. According to a barometer, the
air pressure is 763 mm Hg. What is the pressure exerted by argon in atmospheres
8

Chemistry 120
Chapter Twelve Notes
Gas Stoichiometry
• The Ideal Gas Law aids us in predicting the quantity relationships when reactants or products
are gases.
• An important value to keep in mind is that 1 mole of a gas at STP will occupy 22.4 liters.
• We will use the Ideal Gas Law for similar calculations under different conditions.
Examples
54.8 L of hydrogen gas is combined with excess oxygen at STP. What mass of water should be
produced
72.9 L of hydrogen (at STP) is reacted with 28.0 L of nitrogen (at STP), producing ammonia.
What volume will the ammonia produced occupy at 1.25 atm and 325 K
9

Chemistry 120
Chapter Twelve Notes
65.0 grams of nitrogen gas and 15.0 grams of hydrogen gas are combined in a 10.0 L container.
The gases react, producing ammonia gas. Assume that the reaction goes to completion.
a. What is the partial pressure of each gas in the container at the end of the reaction,
assuming the total pressure in the container is 2.50 atm
b. What is the temperature of the gas sample in the flask under the conditions in part a
c. Suppose that all the ammonia produced in this reaction is dissolved into water to make 555
mL of solution. How many milliliters of 3.00 M H 2 SO 4 solution would be required to
completely neutralize the ammonia solution
10