Leonardo Alvarez

Leonardo Alvarez
Chemistry
Block 3

Converting Real Things
Table 1
Using the scale, come up with a conversion ratio just by looking at the scale and prove that it works
but converting 10 grams to ounces. (Hint: 7grams and 15 grams)
Table 2
Using a 50 mL graduated cylinder, fill a 600mL beaker 66.7% full.
Table 3
Using the measuring cup determine how many mL are in 4 ounces of water. Make a conversion ratio
you could use to do other conversions.
Table 4
Using a meter stick, measure the back table in inches, feet and yards and convert them into
centimeters and meters.
Table 5
Using the scale, find the mass of a text book on your table in ounces (round to the nearest ounce)
and cover it to grams.
Table 6
Using the ruler, measure the length of a piece of paper in inches and then convert that into meters.
Make a conversion ratio you could use to do other conversions.

Dimensional Analysis
This is a way to convert from one unit of a given substance to
another unit using ratios or conversion units. What this video
www.youtube.com/watch?v=aZ3J60GYo6U
Let’ look at a couple of examples:
1. Convert 2.6 qt to mL.
First we need a ratio or conversion unit so that we can go from quarts to milliliters. 1.00 qt = 946 mL
Next write down what you are starting with
2.6 qt
Then make you conversion tree
2.6 qt
Then fill in the units in your ratio so that you can cancel out the original unit and will be left with the
unit you need for the answer. Cross out units, one at a time that are paired, and one on top one on
the bottom.
2.6 qt mL
qt
Now fill in the values from the ratio.
2.6 qt 946 mL
1.00 qt
Now multiply all numbers on the top and multiply all numbers on the bottom and write them as a
fraction.
2.6 qt 946 mL = 2,459.6 mL
1.00 qt 1.00
Now divide the top number by the bottom number and write that number with the unit that was not
crossed out.

1qt=32 oz 1gal = 4qts 1.00 qt = 946 mL 1L = 1000mL
2. Convert 8135.6 mL to quarts
=
3. Convert 115.2 oz to mL
=
4. Convert 2.3 g to Liters
=
5. Convert 8.42 L to oz
=
Go to http://science.widener.edu/svb/tutorial/ chose #7 “Converting Volume” and do 5 more in the
space provided.
1. Convert _________ to _________
=
2. Convert _________ to _________
=
3. Convert _________ to _________
=
4. Convert _________ to _________
=
5. Convert _________ to _________
=

To use the Stair-Step method, find the prefix the original measurement starts with. (ex. milligram)
If there is no prefix, then you are starting with a base unit.
Find the step which you wish to make the conversion to. (ex. decigram)
Count the number of steps you moved, and determine in which direction you moved (left or right).
The decimal in your original measurement moves the same number of places as steps you moved and in the
same direction. (ex. milligram to decigram is 2 steps to the left, so 40 milligrams = .40 decigrams)
If the number of steps you move is larger than the number you have, you will have to add zeros to hold the
places. (ex. kilometers to meters is three steps to the right, so 10 kilometers would be equal to 10,000 m)
That’s all there is to it! You need to be able to count to 6, and know your left from your right!
1) Write the equivalent
a) 5 dm =_______m b) 4 mL = ______L c) 8 g = _______mg
d) 9 mg =_______g e) 2 mL = ______L f) 6 kg = _____g
g) 4 cm =_______m h) 12 mg = ______ g i) 6.5 cm 3 = _______L
j) 7.02 mL =_____cm 3 k) .03 hg = _______ dg l) 6035 mm _____cm
m) .32 m = _______cm n) 38.2 g = _____kg

2. One cereal bar has a mass of 37 g. What is the mass of 6 cereal bars? Is that more than or less
than 1 kg? Explain your answer.
3. Wanda needs to move 110 kg of rocks. She can carry l0 hg each trip. How many trips must she
make? Explain your answer.
4. Dr. O is playing in her garden again She needs 1 kg of potting soil for her plants. She has 750 g.
How much more does she need? Explain your answer.
5. Weather satellites orbit Earth at an altitude of 1,400,000 meters. What is this altitude in kilometers?
6. Which unit would you use to measure the capacity? Write milliliter or liter.
a) a bucket __________
b) a thimble __________
c) a water storage tank__________
d) a carton of juice__________
7. Circle the more reasonable measure:
a) length of an ant 5mm or 5cm
b) length of an automobile 5 m or 50 m
c) distance from NY to LA 450 km or 4,500 km
d) height of a dining table 75 mm or 75 cm
8. Will a tablecloth that is 155 cm long cover a table that is 1.6 m long? Explain your answer.
9. A dollar bill is 15.6 cm long. If 200 dollar bills were laid end to end, how many meters long would
the line be?
10. The ceiling in Jan’s living room is 2.5 m high. She has a hanging lamp that hangs down 41 cm.
Her husband is exactly 2 m tall. Will he hit his head on the hanging lamp? Why or why not?

Using SI Units
Match the terms in Column II with the descriptions in Column I. Write the letters of the correct term in
the blank on the left.
Column I Column II
_____ 1. distance between two points
a. time
_____ 2. SI unit of length
_____ 3. tool used to measure length
_____ 4. units obtained by combining other units
_____ 5. amount of space occupied by an object
_____ 6. unit used to express volume
_____ 7. SI unit of mass
_____ 8. amount of matter in an object
_____ 9. mass per unit of volume
_____ 10. temperature scale of most laboratory thermometers
_____ 11. instrument used to measure mass
_____ 12. interval between two events
_____ 13. SI unit of temperature
_____ 14. SI unit of time
_____ 15. instrument used to measure temperature
b. volume
c. mass
d. density
e. meter
f. kilogram
g. derived
h. liter
i. second
j. Kelvin
k. length
1. balance
m. meterstick
n. thermometer
o. Celsius
Circle the two terms in each group that are related. Explain how the terms are related.
16. Celsius degree, mass, Kelvin _____________________________________________________
________________________________________________________________________________
17. balance, second, mass __________________________________________________________
________________________________________________________________________________
18. kilogram, liter, cubic centimeter __________________________________________________
________________________________________________________________________________
19. time, second, distance __________________________________________________________
________________________________________________________________________________
20. decimeter, kilometer, Kelvin _____________________________________________________
________________________________________________________________________________

1. How many meters are in one kilometer? __________
2. What part of a liter is one milliliter? __________
3. How many grams are in two dekagrams? __________
4. If one gram of water has a volume of one milliliter, what would the mass of one liter of water be in
kilograms?__________
5. What part of a meter is a decimeter? __________
In the blank at the left, write the term that correctly completes each statement. Choose from the terms
listed below.
Metric SI standard ten
prefixes ten tenth
6. An exact quantity that people agree to use for comparison is a ______________ .
7. The system of measurement used worldwide in science is _______________ .
8. SI is based on units of _______________ .
9. The system of measurement that was based on units of ten was the _______________ system.
10. In SI, _______________ are used with the names of the base unit to indicate the multiple of ten
that is being used with the base unit.
11. The prefix deci- means _______________ .

Standards of Measurement
Fill in the missing information in the table below.
S.I prefixes and their meanings
Prefix
Meaning
0.001
0.01
deci- 0.1
10
hecto- 100
1000
Circle the larger unit in each pair of units.
1. millimeter, kilometer 4. centimeter, millimeter
2. decimeter, dekameter 5. hectogram, kilogram
3. hectogram, decigram
6. In SI, the base unit of length is the meter. Use this information to arrange the following units of
measurement in the correct order from smallest to largest.
Write the number 1 (smallest) through 7 - (largest) in the spaces provided.
_____ a. kilometer
_____ b. centimeter
_____ c. meter
_____ e. hectometer
_____ f. millimeter
_____ g. decimeter
_____ d. dekameter
Use your knowledge of the prefixes used in SI to answer the following questions in the spaces
provided.
7. One part of the Olympic games involves an activity called the decathlon. How many events do you
think make up the decathlon?_____________________________________________________
8. How many years make up a decade? _______________________________________________
9. How many years make up a century? ______________________________________________
10. What part of a second do you think a millisecond is? __________________________________

SCIENTIFIC NOTATION RULES
How to Write Numbers in Scientific Notation
Scientific notation is a standard way of writing very large and very small numbers so that they're
easier to both compare and use in computations. To write in scientific notation, follow the form
N X 10 ᴬ
where N is a number between 1 and 10, but not 10 itself, and A is an integer (positive or negative
number).
RULE #1: Standard Scientific Notation is a number from 1 to 9 followed by a decimal and the
remaining significant figures and an exponent of 10 to hold place value.
Example:
5.43 x 10 2 = 5.43 x 100 = 543
8.65 x 10 – 3 = 8.65 x .001 = 0.00865
****54.3 x 10 1 is not Standard Scientific Notation!!!
RULE #2: When the decimal is moved to the Left the exponent gets Larger, but the value of the
number stays the same. Each place the decimal moves Changes the exponent by one (1). If you
move the decimal to the Right it makes the exponent smaller by one (1) for each place it is moved.
Example:
6000. x 10 0 = 600.0 x 10 1 = 60.00 x 10 2 = 6.000 x 10 3 = 6000
(Note: 10 0 = 1)
All the previous numbers are equal, but only 6.000 x 10 3 is in proper Scientific Notation.

RULE #3: To add/subtract in scientific notation, the exponents must first be the same.
Example:
(3.0 x 10 2 ) + (6.4 x 10 3 ); since 6.4 x 10 3 is equal to 64. x 10 2 . Now add.
(3.0 x 10 2 )
+ (64. x 10 2 )
67.0 x 10 2 = 6.70 x 10 3 = 6.7 x 10 3
67.0 x 10 2 is mathematically correct, but a number in standard scientific notation can only
have one number to the left of the decimal, so the decimal is moved to the left one place and
one is added to the exponent.
Following the rules for significant figures, the answer becomes 6.7 x 10 3 .
RULE #4: To multiply, find the product of the numbers, then add the exponents.
Example:
(2.4 x 10 2 ) (5.5 x 10 –4 ) = ? [2.4 x 5.5 = 13.2]; [2 + -4 = -2], so
(2.4 x 10 2 ) (5.5 x 10 –4 ) = 13.2 x 10 –2 = 1.3 x 10 – 1
RULE #5: To divide, find the quotient of the number and subtract the exponents.
Example:
(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = ? [3.3 / 9.1 = .36]; [-6 – (-8) = 2], so
(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = .36 x 10 2 = 3.6 x 10 1

1. 7,485 6. 1.683
2. 884.2 7. 3.622
3. 0.00002887 8. 0.00001735
4. 0.05893 9. 0.9736
5. 0.006162 10. 0.08558
11. 6.633 X 10−⁴ 16. 1.937 X 10⁴
12. 4.445 X 10−⁴ 17. 3.457 X 10⁴
13. 2.182 X 10−³ 18. 3.948 X 10−⁵
14. 4.695 X 10² 19. 8.945 X 10⁵
15. 7.274 X 10⁵ 20. 6.783 X 10²

Convert each number from Scientific Notation to real numbers:
1. 7.485 X 10³ 6. 1.683 X 10⁰
2. 8.842 X 10² 7. 3.622 10⁰
3. 2.887 X 10−⁵ 8. 1.735 X 10−⁵
4. 5.893 X 10−² 9. 9.736 X 10−¹
5. 6.162 X 10−³ 10. 8.558 X 10−²
Convert each number from a real number to Scientific Notation:
11. 0.0006633 16. 1,937,000
12. 0.0004445 17. 34,570
13. 0.002182 18. 0.00003948
14. 469.5 19. 894,500
15. 727,400 20. 678.3

Significant Figures Rules
There are three rules on determining how many significant figures are in a
number:
1. Non-zero digits are always significant.
2. Any zeros between two significant digits are significant.
3. A final zero or trailing zeros in the DECIMAL PORTION ONLY are
significant.
Please remember that, in science, all numbers are based upon measurements (except for a very few
that are defined). Since all measurements are uncertain, we must only use those numbers that are
meaningful.
Not all of the digits have meaning (significance) and, therefore, should not be written down. In
science, only the numbers that have significance (derived from measurement) are written.
Rule 1: Non-zero digits are always significant.
If you measure something and the device you use (ruler, thermometer, triple-beam, balance, etc.)
returns a number to you, then you have made a measurement decision and that ACT of measuring
gives significance to that particular numeral (or digit) in the overall value you obtain.
Hence a number like 46.78 would have four significant figures and 3.94 would have three.
Rule 2: Any zeros between two significant digits are significant.
Suppose you had a number like 409. By the first rule, the 4 and the 9 are significant. However, to
make a measurement decision on the 4 (in the hundred's place) and the 9 (in the one's place), you
HAD to have made a decision on the ten's place. The measurement scale for this number would have
hundreds, tens, and ones marked.
Like the following example:
These are sometimes called "captured zeros."
If a number has a decimal at the end (after the one’s place) then all digits (numbers) are significant
and will be counted.
In the following example the zeros are significant digits and highlighted in blue.
960.
70050.

Rule 3: A final zero or trailing zeros in the decimal portion ONLY are
significant.
This rule causes the most confusion among students.
In the following example the zeros are significant digits and highlighted in blue.
0.07030
0.00800
Here are two more examples where the significant zeros are highlighted in blue.
When Zeros are Not Significant Digits
4.7 0 x 10−³
6.5 0 0 x 10⁴
Zero Type # 1 : Space holding zeros in numbers less than one.
In the following example the zeros are NOT significant digits and highlighted in red.
0.09060
0.00400
These zeros serve only as space holders. They are there to put the decimal point in its correct
location.
They DO NOT involve measurement decisions.
Zero Type # 2 : Trailing zeros in a whole number.
In the following example the zeros are NOT significant digits and highlighted in red.
200
25000
For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point)
of the numbers ONLY. Here is what to do:
1) Count the number of significant figures in the decimal portion of each number in the problem. (The
digits to the left of the decimal place are not used to determine the number of decimal places in the
final answer.)
2) Add or subtract in the normal fashion.
3) Round the answer to the LEAST number of places in the decimal portion of any number in the
problem
The following rule applies for multiplication and division:
The LEAST number of significant figures in any number of the problem determines the number of
significant figures in the answer.
This means you MUST know how to recognize significant figures in order to use this rule.

How Many Significant Digits for Each Number?
1) 2359 = ______
2) 2.445 x 10−⁵= ______
3) 2.93 x 10⁴= ______
4) 1.30 x 10−⁷= ______
5) 2604 = ______
6) 9160 = ______
7) 0.0800 = ______
8) 0.84 = ______
9) 0.0080 = ______
10) 0.00040 = ______
11) 0.0520 = ______
12) 0.060 = ______
13) 6.90 x 10−¹= ______
14) 7.200 x 10⁵= ______
15) 5.566 x 10−²= ______
16) 3.88 x 10⁸= ______
17) 3004 = ______
18) 0.021 = ______
19) 240 = ______
20) 500 = ______

For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point) of the
numbers ONLY. Here is what to do:
1) Count the number of significant figures in the decimal portion of each number in the problem. (The
digits to the left of the decimal place are not used to determine the number of decimal places in the
final answer.)
2) Add or subtract in the normal fashion.
3) Round the answer to the LEAST number of places in the decimal portion of any number in the
problem.
Solve the Problems and Round Accordingly...
1) 43.287 + 5.79 + 6.284 = _______
2) 87.54 - 3.3 = _______
3) 99.1498 + 6.5397 + 9.7 = _______
4) 5.868 - 5.1 = _______
5) 59.9233 + 86.21 + 99.396 = _______
6) 7.7 + 26.756 = _______
7) 66.8 + 2.3 + 4.8516 = _______
8) 9.7419 + 43.545 = _______
9) 4.8976 + 48.4644 + 1.514 = _______
10) 4.335 + 35.85 = _______
11) 9.448 - 1.7 = _______
12) 75.826 - 8.6555 = _______
13) 57.2 + 23.814 = _______
14) 77.684 - 4.394 = _______
15) 26.4496 + 3.339 = _______
16) 9.6848 + 29.85 = _______
17) 63.11 + 2.5412 + 4.93 = _______
18) 11.2471 + 75.4 = _______
19) 73.745 - 8.755 = _______
20) 6.5238 + 1.7 + 27.79 = _______

The following rule applies for multiplication and division:
The LEAST number of significant figures in any number of the problem determines the number of
significant figures in the answer.
This means you MUST know how to recognize significant figures in order to use this rule.
Solve the Problems and Round Accordingly...
1) 0.6 x 65.0 x 602 = __________
2) 720 ÷ 7.7 = __________
3) 929 x 0.3 = __________
4) 300 ÷ 44.31 = __________
5) 608 ÷ 9.8 = __________
6) 0.06 x 0.079 = __________
7) 0.008 x 72.91 x 7000 = __________
8) 73.94 x 67 x 780 = __________
9) 0.62 x 0.097 x 40 = __________
10) 600 x 10 x 5030 = __________
11) 5200 ÷ 4.46 = __________
12) 0.0052 x 0.4 x 107 = __________
13) 0.099 x 8.8 = __________
14) 0.0095 x 5.2 = __________
15) 8000 ÷ 4.62 = __________
16) 0.6 x 0.8 = __________
17) 2.84 x 0.66 = __________
18) 0.5 x 0.09 = __________
19) 8100 ÷ 34.84 = __________
20) 8.24 x 6.9 x 8100 = __________

Question Sig Figs Question Add & Subtract Question Multiple & Divide
1 4 1 55.36 1 20,000
2 4 2 84.2 2 94
3 3 3 115.4 3 300
4 3 4 0.8 4 7
5 4 5 245.53 5 62
6 3 6 34.5 6 0.005
7 3 7 74.0 7 4,000
8 2 8 53.287 8 3,900,000
9 2 9 54.876 9 2
10 2 10 40.19 10 30,000,000
11 3 11 7.7 11 1,200
12 2 12 67.170 12 0.2
13 3 13 81.0 13 0.87
14 4 14 73.290 14 0.049
15 4 15 29.789 15 2,000
16 3 16 39.53 16 0.5
17 4 17 70.58 17 1.9
18 2 18 86.6 18 0.05
19 2 19 64.990 19 230
20 1 20 36.0 20 460,000

Atoms Are Building Blocks
Atoms are the basis of chemistry. They are the basis for everything in the Universe. You
should start by remembering that matter is composed of atoms. Atoms and the study of
atoms are a world unto themselves. We're going to cover basics like atomic structure
and bonding between atoms.
Smaller Than Atoms?
Are there pieces of matter that are smaller than atoms?
Sure there are. You'll soon be learning that atoms are
composed of pieces like electrons, protons, and neutrons.
But guess what? There are even smaller particles moving
around in atoms. These super-small particles can be found
inside the protons and neutrons. Scientists have many
names for those pieces, but you may have heard of
nucleons and quarks. Nuclear chemists and physicists
work together at particle accelerators to discover the
presence of these tiny, tiny, tiny pieces of matter.
Even though super-tiny atomic particles exist, you only
need to remember the three basic parts of an atom: electrons, protons, and neutrons.
What are electrons, protons, and neutrons? A picture works best to show off the idea.
You have a basic atom. There are three types of pieces in that atom: electrons, protons,
and neutrons. That's all you have to remember. Three things! As you know, there are
almost 120 known elements in the periodic table. Chemists and physicists haven't
stopped there. They are trying to make new ones in labs every day. The thing that
makes each of those elements different is the number of electrons, protons, and
neutrons. The protons and neutrons are always in the center of the atom. Scientists call
the center region of the atom the nucleus. The nucleus in
a cell is a thing. The nucleus in an atom is a place where
you find protons and neutrons. The electrons are always
found whizzing around the center in areas called shells or
orbitals.
You can also see that each piece has either a "+", "-", or a
"0." That symbol refers to the charge of the particle. Have
you ever heard about getting a shock from a socket, static
electricity, or lightning? Those are all different types of
electric charges. Those charges are also found in tiny particles of matter. The electron
always has a "-", or negative, charge. The proton always has a "+", or positive, charge. If
the charge of an entire atom is "0", or neutral, there are equal numbers of positive and
negative pieces. Neutral means there are equal numbers of electrons and protons. The
third particle is the neutron. It has a neutral charge, also known as a charge of zero. All
atoms have equal numbers of protons and electrons so that they are neutral. If there are
more positive protons or negative electrons in an atom, you have a special atom called
an ion.

http://www.learner.org/interactives/periodic/basics_interactive.html

Looking at Ions
We haven’t talked about ions before, so let’s get down to basics. The
atomic number of an element, also called a proton number, tells you the
number of protons or positive particles in an atom. A normal atom has a
neutral charge with equal numbers of positive and negative particles.
That means an atom with a neutral charge is one where the number of
electrons is equal to the atomic number. Ions are atoms with extra
electrons or missing electrons. When you are missing an electron or
two, you have a positive charge. When you have an extra electron
or two, you have a negative charge.
What do you do if you are a sodium (Na) atom? You have eleven
electrons — one too many to have an entire shell filled. You need to
find another element that will take that electron away from you. When you lose that
electron, you will you’ll have full shells. Whenever an atom has full shells, we say it is
"happy." Let's look at chlorine (Cl). Chlorine has seventeen electrons and only needs
one more to fill its third shell and be "happy." Chlorine will take your extra sodium
electron and leave you with 10 electrons inside of two filled shells. You are now a happy
atom too. You are also an ion and missing one electron. That missing electron gives you
a positive charge. You are still the element sodium, but you are now a sodium ion (Na + ).
You have one less electron than your atomic number.
Ion Characteristics
So now you've become a sodium ion. You have ten electrons.
That's the same number of electrons as neon (Ne). But you
aren't neon. Since you're missing an electron, you aren't really
a complete sodium atom either. As an ion you are now
something completely new. Your whole goal as an atom was
to become a "happy atom" with completely filled electron
shells. Now you have those filled shells. You have a lower
energy. You lost an electron and you are "happy." So what
makes you interesting to other atoms? Now that you have
given up the electron, you are quite electrically attractive.
Other electrically charged atoms (ions) of the opposite charge
(negative) are now looking at you and seeing a good partner to
bond with. That's where the chlorine comes in. It's not only chlorine. Almost any ion with
a negative charge will be interested in bonding with you.

Electrovalence
Don't get worried about the big word. Electrovalence is just another word for something
that has given up or taken electrons and become an ion. If you look at the periodic table,
you might notice that elements on the left side usually become positively charged ions
(cations) and elements on the right side get a negative charge (anions). That trend
means that the left side has a positive valence and the right side has a negative
valence. Valence is a measure of how much an atom wants to bond with other atoms. It
is also a measure of how many electrons are excited about bonding with other atoms.
There are two main types of bonding, covalent and electrovalent. You may have heard
of the term "ionic bonds." Ionic bonds are electrovalent bonds. They are just groups of
charged ions held together by electric forces. When in the presence of other ions, the
electrovalent bonds are weaker because of outside electrical forces and attractions.
Sodium and chlorine ions alone have a very strong bond, but as soon as you put those
ions in a solution with H + (Hydrogen ion), OH - (Hydroxide), F - (Fluorine ion) or Mg ++
(Magnesium ion), there are charged distractions that break the Na-Cl bond.
Look at sodium chloride (NaCl) one more time. Salt is a very strong bond when it is
sitting on your table. It would be nearly impossible to break those ionic/electrovalent
bonds. However, if you put that salt into some water (H 2 O), the bonds break very
quickly. It happens easily because of the electrical attraction of the water. Now you have
sodium (Na + ) and chlorine (Cl - ) ions floating around the solution. You should remember
that ionic bonds are normally strong, but they are very weak in water.

Neutron Madness
We have already learned that ions are atoms that are
either missing or have extra electrons. Let's say an atom
is missing a neutron or has an extra neutron. That type of
atom is called an isotope. An atom is still the same
element if it is missing an electron. The same goes for
isotopes. They are still the same element. They are just a
little different from every other atom of the same element.
For example, there are a lot of carbon (C) atoms in the
Universe. The normal ones are carbon-12. Those atoms have 6 neutrons. There are a
few straggler atoms that don't have 6. Those odd ones may have 7 or even 8 neutrons.
As you learn more about chemistry, you will probably hear about carbon-14. Carbon-14
actually has 8 neutrons (2 extra). C-14 is considered an isotope of the element carbon.
Messing with the Mass
If you have looked at a periodic table, you may have noticed that the atomic mass of
an element is rarely an even number. That happens because of the isotopes. If you are
an atom with an extra electron, it's no big deal. Electrons don't have much of a mass
when compared to a neutron or proton.
Atomic masses are calculated by figuring out the
amounts of each type of atom and isotope there are in
the Universe. For carbon, there are a lot of C-12, a
couple of C-13, and a few C-14 atoms. When you
average out all of the masses, you get a number that is a
little bit higher than 12 (the weight of a C-12 atom). The
average atomic mass for the element is actually 12.011.
Since you never really know which carbon atom you are
using in calculations, you should use the average mass
of an atom.
Bromine (Br), at atomic number 35, has a greater variety of isotopes. The atomic mass
of bromine (Br) is 79.90. There are two main isotopes at 79 and 81, which average out
to the 79.90amu value. The 79 has 44 neutrons and the 81 has 46 neutrons. While it
won't change the average atomic mass, scientists have made bromine isotopes with
masses from 68 to 97. It's all about the number of neutrons. As you move to higher
atomic numbers in the periodic table, you will probably find even more isotopes for
each element.

P
N
P
N

P
N
P
N

Electron Configuration
Color the sublevel:
s = Red
d = Green
p = Blue
f = Orange
Write in sublevels
Write period, sublevel and super scripts.
Ctrl Shift =
gives you super scripts

www.youtube.com/watch?v=jtYzEzykFdg
www.youtube.com/watch?
annotation_id=annotation_2076&feature=iv&src_vid=jtYzEzykFdg&v=cOlac8ruD_0
www.youtube.com/watch?
annotation_id=annotation_570977&feature=iv&src_vid=cOlac8ruD_0&v=lR2vqHZWb5A

Electron Configuration
In order to write the electron configuration for an atom you must know the 3 rules of
electron configurations.
1. Aufbau
Notation
nO e
where
n is the energy level
O is the orbital type (s, p, d, or f)
e is the number of electrons in that orbital shell
Principle
electrons will first occupy orbitals of the lowest energy level
2. Hund rule
when electrons occupy orbitals of equal energy, one electron enters each orbital until
all the orbitals contain one electron with the same spin.
3. Pauli exclusion principle
an orbital contains a maximum of 2 electrons and
paired electrons will have opposite spin

In the space below, write the unabbreviated electron configurations of the following elements:
1) sodium ________________________________________________
2) iron ________________________________________________
3) bromine ________________________________________________
4) barium ________________________________________________
5) neptunium ________________________________________________
In the space below, write the abbreviated electron configurations of the following elements:
6) cobalt ________________________________________________
7) silver ________________________________________________
8) tellurium ________________________________________________
9) radium ________________________________________________
10) lawrencium ________________________________________________
Determine what elements are denoted by the following electron configurations:
11) 1s²s²2p⁶3s²3p⁴ ____________________
12) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹ ____________________
13) [Kr] 5s²4d¹⁰5p³ ____________________
14) [Xe] 6s²4f¹⁴5d⁶ ____________________
15) [Rn] 7s²5f¹¹ ____________________
Identify the element or determine that it is not a valid electron configuration:
16) 1s²2s²2p⁶3s²3p⁶4s²4d¹⁰4p⁵ ____________________
17) 1s²2s²2p⁶3s³3d⁵ ____________________
18) [Ra] 7s²5f⁸ ____________________
19) [Kr] 5s²4d¹⁰5p⁵ ____________________
20) [Xe] ____________________
1)sodium 1s 2 2s 2 2p 6 3s 1 2)iron 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6
3)bromine 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 4)barium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2
5)neptunium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 5 6)cobalt [Ar] 4s 2 3d 7
7)silver [Kr] 5s 2 4d 9 8)tellurium[Kr] 5s 2 4d 10 5p 4
9)radium [Rn] 7s 2 10)lawrencium[Rn] 7s 2 5f 14 6d 1
1s 2 2s 2 2p 6 3s 2 3p 4 sulfur 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 rubidium
[Kr] 5s 2 4d 10 5p 3 antimony [Xe] 6s 2 4f 14 5d 6 osmium
[Rn] 7s 2 5f 11 einsteinium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 5 not valid (take a look at “4d”)
1s 2 2s 2 2p 6 3s 3 3d 5 not valid (3p comes after 3s) [Ra] 7s 2 5f 8 not valid (radium isn’t a noble gas)
[Kr] 5s 2 4d 10 5p 5 valid iodine
20)[Xe] not valid (an element can’t be its own electron configuration)

Using Wikipedia define the 8 categories of elements on pages 168 and 169 on the left
page.
Color your periodic table similar to the one on Pages 168—169.
Alkali Metals -
Alkali Earth Metals-
Other Metals-
Metalloids-
Nonmetals-
Noble Gases-
Transions metals-
Inner transion metals-

Atomic Size
Increases
Increases
As you go from the top to boom of the energy level the energy levels of the elements increase
As you go from le to right of the periodic table the mass increases which results in a bigger gravitaonal pull

Ionizaon Energy
Increase
Ionizaon energy is the energy needed to take away an electron away from
an atom
increase

Electronegavity
increases
Electronegavity is the ability for an atom to take an electron
increases

Ion Size
increase
increase
How they compare to atoms to ions
Catons decrease in size because they giving up electrons and go down an energy level

Create groups for these Scientist and explain your groupings
(use the information you got from your research)

Research these Scientist and summarize their contributions to chemistry
Antoine Henri Becquerel
Niels Bohr
Louis de Barogilie
Glenn Seaborg
Hantaro Nagaoka
Democritus
Marie and Pierre Curie
Eugene Goldstein
Dmitri Mendeleev
J.J. Thomson
James Chadwick
Erwin Shrodinger
John Dalton
Lothar Meyer
Robert Millikan
J.W. Dobereiner
Ernest Rutherford

Nuclear radiation notes
Chapter 25.1
Block 3
Leo Alvarez
25.1
Radioactivity
Atoms emit electromagnetic radiation when an electron moves from a higher energy level to a
lower energy level
Antoine made an accidental discovery by studying the ability of uranium salts that have been
exposed to sunlight to fog photographic film plates
In a chemical reaction atoms tend to attain more stable electron configuration by transferring or
sharing electrons
Radio active decay can be done without any input of energy
Radiation is transmitted during radioactive decay
Types of radiation
Some radioactive sources emit helium nuclei and are also called alpha particles
When an atom loses an alpha particle, the atomic number, the atomic number of the product is
lower by two and its mass number is lowered by four

Radioisotopes can cause harm when ingested but do no harm in penetrating objects
When a electron breaks apart from a neutron in an atom its called beta particle
Beta particles are more penetrating than alpha particles are
Gamma rays have no mass or electrical charge but are extremely penetrating
25.2
Nuclear stability and decay
More than 1500 nuclei are known and only 264 are stable and do not decay
When a nucleus is unstable it can undergo spontaneous decay for different reasons like having an
unstable amount of neutrons relative to the number of protons
A beta transmission is when a nucleus increases the amount of protons while decreasing the
amount of neutrons
Half life
Every radioisotope is measured by its half life which is the time required for one half of
the nuclei in a radioisotope sample to decay to products

Notes
The students will learn how ionic compounds from and how metallic bonding affects the
properties of metals
Metals have several qualities that are unique, such as the ability to conduct electricity a low ionization
energy, and a low electro-negativity
Their physical properties include a lustrous (shiny) appearance, and they are malleable
and ductile. Metals have a crystal structure.
The strength of a metallic bond depends on three things:
The number of electrons that become delocalized from the metal
The charge of the cation (metal).
The size of the cation.
A strong metallic bond will be the result of more delocalized electrons
Metallic bonds are strong and require a great deal of energy to break
conduct-the action or manner of managing an activity or organization.
delocalized- detach or remove (something) from a particular place or location, or from local limitations.
ionization-the condition of being dissociated into ions (as by heat or radiation or chemical reaction or
electrical discharge); "the ionization of a gas"
Be 2 N 3-
Mg 3 N 2
magnesium nitride
1-same
2-di
3-tri
4-tetra

Metallic Character
Metals have several qualities that are unique, such as the ability to conduct electricity, a low
ionization energy, and a low electronegativity (so they will give up electrons easily, i.e., they are
cations). Their physical properties include a lustrous (shiny) appearance, and they are malleable
and ductile. Metals have a crystal structure.
Metals that are malleable can be beaten into thin sheets, for example: aluminum foil.
Metals that are ductile can be drawn into wires, for example: copper wire.
Bonding Characteristics
The strength of a metallic bond depends on three things:
The number of electrons that become delocalized from the metal
The charge of the cation (metal).
The size of the cation.
A strong metallic bond will be the result of more delocalized electrons, which causes the
effective nuclear charge on electrons on the cation to increase, in effect making the size of the
cation smaller. Metallic bonds are strong and require a great deal of energy to break, and
therefore metals have high melting and boiling points.
A metallic bonding theory must explain how so much bonding can occur with such few electrons
(since metals are located on the left side of the periodic table and do not have many electrons in
their valence shells). The theory must also account for all of a metal's unique chemical and
physical properties.
http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/General_Principles/Met
allic_Bonding

Notes
Metallic bonding constitutes the electrostatic attractive forces between the delocalized electrons.
metallic bonding is sometimes compared with that of molten salts
conduction electrons divide their density equally over all atoms that function as neutral
Metallic bond is not the only type of chemical bonding a metal can exhibit, even as a simple substance
gallium consists of covalently-bound pairs of atoms in both liquid and solid state—these pairs form a crystal
lattice with metallic bonding between them.
gallium-the chemical element of atomic number 31, a soft silvery-white metal that melts at about 30°C, just above
room temperature.
lattice-the structure of fissionable and non-fissionable materials geometrically arranged within a nuclear reactor.
electrostatic force- the electrostatic interaction between electrically charged particles.

Metallic bonding constitutes the electrostatic attractive forces between the delocalized electrons,
called conduction electrons, gathered in an electron cloud and the positively charged metal ions.
Understood as the sharing of "free" electrons among a lattice of positively charged ions (cations),
metallic bonding is sometimes compared with that of molten salts; however, this simplistic
view[which?] holds true for very few[which?] metals. In a more quantum-mechanical view, the
conduction electrons divide their density equally over all atoms that function as neutral (noncharged)
entities.[citation needed] Metallic bonding accounts for many physical properties of
metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and
luster.[1][2][3][4]
Although the term "metallic bond" is often used in contrast to the term "covalent bond", it is
preferable[by whom?] to use the term metallic bonding, because this type of bonding is
collective in nature and a single "metallic bond" does not exist. Metallic bond is not the only type
of chemical bonding a metal can exhibit, even as a simple substance. For example, elemental
gallium consists of covalently-bound pairs of atoms in both liquid and solid state—these pairs
form a crystal lattice with metallic bonding between them. Another example of a metal–metal
covalent bond is mercurous ion (Hg2+
2).
http://en.wikipedia.org/wiki/Metallic_bond

Ionic bonds
Metal + non metals
Cations + Anion
Cation Non metal (metal)
NaCl= Sodium Chloride = Na Cl
Be F = Beryllium Floride = BeF2
Formula unit is used to show how any electrons the atom will have
Coveland Bond atoms
Sharing electrons
H20
The bigger the mass of an atom, the more electrons it will take
H202 + dihydrogen
Oxygen (O) can disappear sometimes (99 percent of the time)

PCl3
Phosphurus Triclohride\
SBr2

Name Formula Charge
Dichromate Cr₂O₇ 2-
Sulfate SO₄ 2-
Hydrogen Carbonate HCO₃ 1-
Hypochlorite ClO 1-
Phosphate PO₄ 3-
Nitrite NO₂ 1-
Chlorite ClO₂ 1-
Dihydrogen phosphate H₂PO₄ 1-
Chromate CrO₄ 2-
Carbonate CO₃ 2-
Hydroxide OH 1-
Hydrogen phosphate HPO₄ 2-
Ammonium NH₄ 1+
Acetate C₂H₃O₂ 1-
Perchlorate ClO₄ 1-
Permanganate MnO₄ 1-
Chlorate ClO₃ 1-
Hydrogen Sulfate HSO₄ 1-
Phosphite PO₃ 3-
Sulfite SO₃ 2-
Silicate SiO₃ 2-
Nitrate NO₃ 1-
Hydrogen Sulfite HSO₃ 1-
Oxalate C₂O₄ 2-
Cyanide CN 1-
Hydronium H₃O 1+
Thiosulfate S₂O₃ 2-

Orbital Equation Lone Pairs Angle Name

www.youtube.com/watch?v=AsqEkF7hcII
www.youtube.com/watch?v=tEn0N4R2dqA
www.youtube.com/watch?v=Pft2CASl0M0
www.youtube.com/watch?v=rwhJklbK8R0
The Mole

www.youtube.com/watch?v=BTRm8PwcZ3U
www.youtube.com/watch?v=F9NkYSKJifs
www.youtube.com/watch?v=xPdqEX_WMjo
Molar Mass

Mole Conversions

Steps for Mole Conversions

Answer the following questions:
1) How many moles are in 25 grams of water?
Mole Calculation Practice
2) How many grams are in 4.5 moles of Li 2 O?
3) How many molecules are in 23 moles of oxygen?
4) How many moles are in 3.4 x 10 23 molecules of H 2 SO 4 ?
5) How many molecules are in 25 grams of NH 3 ?
6) How many grams are in 8.2 x 10 22 molecules of N 2 I 6 ?

How to Balance Chemical Equations
A chemical equation is a theoretical or written representation of what happens during a chemical
reaction. The law of conservation of mass states that no atoms can be created or destroyed in a
chemical reaction, so the number of atoms that are present in the reactants has to balance the
number of atoms that are present in the products. Follow this guide to learn how to balance chemical
equations.
Step 1
Write down your given equation. For this example, we will use:
C 3 H 8 + O 2 --> H 2 O + CO 2
Step 2
Write down the number of atoms that you have on each side of the equation. Look at the subscripts
next to each atom to find the number of atoms in the equation.
Left side: 3 carbon, 8 hydrogen and 2 oxygen
Right side: 1 carbon, 2 hydrogen and 3 oxygen

Step 3
Always leave hydrogen and oxygen for last. This means that you will need to balance the carbon
atoms first.
Step 4
Add a coefficient to the single carbon atom on the right of the equation to balance it with the 3 carbon
atoms on the left of the equation.
C 3 H 8 + O 2 --> H 2 O + 3CO 2
The coefficient 3 in front of carbon on the right side indicates 3 carbon atoms just as the subscript 3
on the left side indicates 3 carbon atoms.
In a chemical equation, you can change coefficients, but you should never alter the subscripts.

Step 5
Balance the hydrogen atoms next. You have 8 on the left side, so you'll need 8 on the right side.
C 3 H 8 + O 2 --> 4H 2 O + 3CO 2
On the right side, we added a 4 as the coefficient because the subscript showed that we already
had 2 hydrogen atoms.
When you multiply the coefficient 4 times the subscript 2, you end up with 8.
Step 6
Finish by balancing the oxygen atoms.
Because we've added coefficients to the molecules on the right side of the equation, the number of
oxygen atoms has changed. We now have 4 oxygen atoms in the water molecule and 6 oxygen
atoms in the carbon dioxide molecule. That makes a total of 10 oxygen atoms.
Add a coefficient of 5 to the oxygen molecule on the left side of the equation. You now have 10
oxygen molecules on each side.
C 3 H 8 + 5O 2 --> 4H 2 O + 3CO 2.
The carbon, hydrogen and oxygen atoms are balanced. Your equation is complete.

1) ___ NaNO 3 + ___ PbO ___ Pb(NO 3 ) 2 + ___ Na 2 O
2) ___ AgI + ___ Fe 2 (CO 3 ) 3 ___ FeI 3 + ___ Ag 2 CO 3
3) ___ C 2 H 4 O 2 + ___ O 2 ___ CO 2 + ___ H 2 O
4) ___ ZnSO 4 + ___ Li 2 CO 3 ___ ZnCO 3 + ___ Li 2 SO 4
5) ___ V 2 O 5 + ___ CaS ___ CaO + ___ V 2 S 5

6) ___ Mn(NO 2 ) 2 + ___ BeCl 2 ___ Be(NO 2 ) 2 + ___ MnCl 2
7) ___ AgBr + ___ GaPO 4 ___ Ag 3 PO 4 + ___ GaBr 3
8) ___ H 2 SO 4 + ___ B(OH) 3 __ B 2 (SO 4 ) 3 + ___ H 2 O
9) ___ S 8 + ___ O 2 ___ SO 2
10) ___ Fe + ___ AgNO 3 ___ Fe(NO 3 ) 2 + ___ Ag

1) 2 NaNO 3 + PbO Pb(NO 3 ) 2 + Na 2 O
2) 6 AgI + Fe 2 (CO 3 ) 3 2 FeI 3 + 3 Ag 2 CO 3
3) C 2 H 4 O 2 + 2 O 2 2 CO 2 + 2 H 2 O
4) ZnSO 4 + Li 2 CO 3 ZnCO 3 + Li 2 SO 4
5) V 2 O 5 + 5 CaS 5 CaO + V 2 S 5
6) Mn(NO 2 ) 2 + BeCl 2 Be(NO 2 ) 2 + MnCl 2
7) 3 AgBr + GaPO 4 Ag 3 PO 4 + GaBr 3
8) 3 H 2 SO 4 + 2 B(OH) 3 B 2 (SO 4 ) 3 + 6 H 2 O
9) S 8 + 8 O 2 8 SO 2
10) Fe + 2 AgNO 3 Fe(NO 3 ) 2 + 2 Ag
Additional Notes:

Categories of Reactions
All chemical reactions can be placed into one of six categories. Here they are, in no
particular order:
1) Synthesis: A synthesis reaction is when two or more simple compounds combine to form a
more complicated one. These reactions come in the general form of:
A + B ---> AB
One example of a synthesis reaction is the combination of iron and sulfur to form iron (II) sulfide:
8 Fe + S 8 ---> 8 FeS
If two elements or very simple molecules combine with each other, it’s probably a synthesis reaction.
The products will probably be predictable using the octet rule to find charges.
2) Decomposition: A decomposition reaction is the opposite of a synthesis reaction - a
complex molecule breaks down to make simpler ones. These reactions come in the general form:
AB ---> A + B
One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen
gas:
2 H 2O ---> 2 H 2 + O 2
If one compound has an arrow coming off of it, it’s probably a decomposition reaction. The products
will either be a couple of very simple molecules, or some elements, or both.
3) Single displacement: This is when one element trades places with another element in a
compound. These reactions come in the general form of:
A + BC ---> AC + B
One example of a single displacement reaction is when magnesium replaces hydrogen in water to
make magnesium hydroxide and hydrogen gas:
Mg + 2 H 2O ---> Mg(OH) 2 + H 2
If a pure element reacts with another compound (usually, but not always, ionic), it’s probably a single
displacement reaction. The products will be the compounds formed when the pure element switches
places with another element in the other compound.
Important note: these reactions will only occur if the pure element on the reactant side of the equation
is higher on the activity series than the element it replaces.

4) Double displacement: This is when the anions and cations of two different molecules
switch places, forming two entirely different compounds. These reactions are in the general form:
AB + CD ---> AD + CB
One example of a double displacement reaction is the reaction of lead (II) nitrate with potassium
iodide to form lead (II) iodide and potassium nitrate:
Pb(NO 3) 2 + 2 KI ---> PbI 2 + 2 KNO 3
If two ionic compounds combine, it’s probably a double displacement reaction. Switch the cations
and balance out the charges to figure out what will be made.
Important note: These reactions will only occur if both reactants are soluble in water and only one
product is soluble in water.
5) Acid-base: This is a special kind of double displacement reaction that takes place when an
acid and base react with each other. The H + ion in the acid reacts with the OH - ion in the base,
causing the formation of water. Generally, the product of this reaction is some ionic salt and water:
HA + BOH ---> H 2O + BA
One example of an acid-base reaction is the reaction of hydrobromic acid (HBr) with sodium
hydroxide:
HBr + NaOH ---> NaBr + H 2O
If an acid and a base combine, it’s an acid-base reaction. The products will be an ionic compound
and water.
6) Combustion: A combustion reaction is when oxygen combines with another compound to
form water and carbon dioxide. These reactions are exothermic, meaning they produce heat. An
example of this kind of reaction is the burning of napthalene:
C 10H 8 + 12 O 2 ---> 10 CO 2 + 4 H 2O
If something that has carbon and hydrogen reacts with oxygen, it’s probably a combustion reaction.
The products will be CO 2 and H 2 O.
Follow this series of questions. When you can answer "yes" to a question, then
stop!
1) Does your reaction have two (or more) chemicals combining to form one chemical? If yes, then it's
a synthesis reaction
2) Does your reaction have one large molecule falling apart to make several small ones? If yes, then
it's a decomposition reaction
3) Does your reaction have any molecules that contain only one element? If yes, then it's a single
displacement reaction
4) Does your reaction have water as one of the products? If yes, then it's an acid-base reaction
5) Does your reaction have oxygen as one of it's reactants and carbon dioxide and water as
products? If yes, then it's a combustion reaction
6) If you haven't answered "yes" to any of the questions above, then you've got a double
displacement reaction.

1) NaOH + KNO 3 --> NaNO 3 + KOH
2) CH 4 + 2 O 2 --> CO 2 + 2 H 2 O
3) 2 Fe + 6 NaBr --> 2 FeBr 3 + 6 Na
List what type the following reactions are:
4) CaSO 4 + Mg(OH) 2 --> Ca(OH) 2 + MgSO 4
5) NH 4 OH + HBr --> H 2 O + NH 4 Br
6) Pb + O 2 --> PbO 2
7) Na 2 CO 3 --> Na 2 O + CO 2
Predicting Reaction Products
Predict the products of each of the following chemical reactions. If a reaction will not occur, explain
why not:
Category of Reaction
1) __Ag 2 SO 4 + __NaNO 3 →
2) __NaI + __CaSO 4 →
3) __HNO 3 + __Ca(OH) 2 →
4) __CaCO 3 →
5) __AlCl 3 + __(NH 4 )PO 4 →
6) __Pb + __Fe(NO 3 ) 3 →
7) __C 3 H 6 + __O 2 →
8) __Na + __CaSO 4 →
__________________
__________________
__________________
__________________
__________________
__________________
__________________
__________________

1) double displacement
2) combustion
3) single displacement
4) double displacement
5) acid-base
6) synthesis
7) decomposition
List what type the following reactions are: (answers)
Predicting Reaction Products – Answers
Predict the products of each of the following chemical reactions. If a reaction will not occur, explain why not:
1) ____ Ag 2 SO 4 + ____ NaNO 3 → no reaction!
Examining this reaction, it appears that a double displacement reaction will occur. This would lead to the conclusion that
the products would be AgNO3 and Na2SO4. However, for this reaction to occur, both reactants and only one of the
products must be soluble in water. If you look up the solubilities on a chart, you’ll find that Ag2SO3 is partly soluble in
water, and all of the other compounds are totally soluble in water. This tells us that this reaction will not occur.
2) ____ NaI + ____ CaSO 4 → no reaction!
Another double displacement reaction, this time with Na2SO4 and CaI2 as products. Because both products are soluble
in water and CaSO4 is only partially soluble in water, the conditions for a successful double displacement reaction are not
met.
3) 2 HNO 3 + 1 Ca(OH) 2 → 1 Ca(NO 3 ) 2 + 2 H 2 O
It’s an acid-base reaction, and acid-base reactions occur readily whether or not the reactants are both soluble in water.
4) 1 CaCO 3 → 1 CaO + 1 CO 2
It’s a decomposition reaction. If you didn’t guess that these were the products, you should have at least known that it was
a decomposition reaction and predicted that this would have broken into its constituent elements, Ca, C, and O2.
5) 1 AlCl 3 (aq) + 1 (NH 4 ) 3 PO4(aq) → AlPO 4 (s) + 3 NH 4 Cl(aq)
This is a double displacement reaction, except in this case both of the reactants and only one product are soluble in
water. Because the conditions for a successful reaction are met, the reaction does occur!
6) ____ Pb + ____ Fe(NO 3 ) 3 →
Though this is a single displacement reaction, lead is lower on the activity series than the iron it would replace. As a
result, this reaction does not occur.
7) 2 C 3 H 6 + 9 O 2 → 6 CO 2 + 6 H 2 O
The reactants suggest that this is a combustion reaction, meaning that the products must be carbon dioxide and water.
Once you figure this out, the only thing left to do is balance it, as shown.
8) 2 Na + 1 CaSO 4 → 1 Na 2 SO 4 + 1 Ca
This should clearly be a single displacement reaction. Because sodium is higher on the activity series than calcium, this
reaction does occur.

Balancing Equations Practice Problems
From Widener Chemistry Practice
1 page
The students will learn how balanced chemical equations are used in stoichiometric calculations and
how to calculate amounts of reactants and products in a chemical equation.

Stoichiometry Practice Problems
From Widener Chemistry Practice
2 pages
The students will learn how balanced chemical equations are used in stoichiometric calculations and
how to calculate amounts of reactants and products in a chemical equation.

Percent Yield Practice Problems
from Widener Chemistry Practice
1 pages
The students will learn how balanced chemical equations are used in stoichiometric calculations and
how to calculate amounts of reactants and products in a chemical equation.

Learning goal: the students will learn what are the factors that determine and
characteristics that distinguish gases, liquids and solids and how substances change
from one state to another
Gas
The particles in a gas are usually molecules or atoms, they are considered to be small, hard
spheres with an insignificant volume. Between the particles there is empty space no
attractive or repulsive forces exists between the particles. The motion of the particles in a
gas is rapid, constant and random. Gases fill their containers regardless of the shape and
volume. The particles travel in a straight line until they collide with another object.
Gas pressure results from the force exerted by a gas per unit surface area of an object. It is
the result of billions of rapidly moving particles in a gas simultaneously colliding with an
object. The collisions of atoms and molecules in air with objects results in atmospheric
pressure. It decreases when you climb a mountain because earths atmosphere decreases as
the elevation increases.
The average speed of oxygen molecules in air at 20 C is 1700 km/h. At these high speed the
odor from a pizza in Washington D.C, should reach Mexico city in about 115 minutes. This
does not happen because the molecules responsible for the odor are constantly striking
molecules in air and rebounding in other directions. The aimless path the molecules take is
called a random walk.

Liquids
Both the particles in gases and the particles in liquids have kinetic energy. This energy allows
the particles in gases and liquids to flow past one another. The ability of gases and liquids to
flow allows them to conform to the shape of their containers. There are no attractions
between the particles in a gas. The particles in a liquid are attracted to each other, these
intermolecular attractions keep the particles in a liquid close together.
The interplay between the disruptive motions of particles in a liquid and the attractions
among the particles determines the physical properties of a liquid. Intermolecular
attractions reduce the amount of space between the particles in a liquid. Increasing the
pressure of a liquid has hardly any effect on its volume as well as solids. Therefore, liquids
and solids are known as condensed states of matter.
The conversion occurs of a liquid to a gas or vapor is called vaporization. When the
conversion occurs at a surface of a liquid that is not boiling is called evaporation. During
evaporation , only those molecules with a certain minimum kinetic energy can escape
from the surface of a liquid. The rate of evaporation of a liquid from an open container
increases as the liquid is heated. Heating allows a greater number of particles at the
liquids surface to overcome the attractive forces that keep the liquid in state.

Solid
The general properties of solids reflect the orderly arrangement of their particles and the
fixed locations of their particle. In most solids, the atoms, ions, or molecules are packed
tightly together. These solids are dense and not easy to compress because the particles in
solids tend to vibrate about fixed points, solids do not flow.

Iliana Gonzalez, Gaby Cordon, Leo Alvarez, Eddy Isaac

Boyle’s Law
• As volume decreases, Pressure
increases when the temperature stays
the same

Where:
P1 = initial pressure
V1 = initial volume
P2 = final pressure
V2 = final volume
Boyle’s Formula
P1* V1= P2* V2
EXAMPLE:
P1=105 kpa
V1= 2.50 L
P2= 40.5 kpa
V2= ?
105 kpa x 2.50 kpa = 6.48 L
40.5 kpa

http://www.youtube.com/watch?v=27yqJ9vJ5kQ

Charles’ Law
• Volume increases/decreases as
temperature increases/decreases if
the pressure stays the same

Where:
V1 = initial volume
T1 = initial temperature
V2= final volume
T2= final temperature
Charles’ Formula
V1/T1= V2/ T2
EXAMPLE:
V1 = 6.80 L
T1 = 325 ◦C
V2= ?
T2= 25◦C
6.80 L x 25◦C = .523 L
325 ◦C

http://www.youtube.com/watch?v=7JKVtbe-hV8

Gay-Lussac’s Law
• When one connected value goes up
so does the other, meaning that if the
volume stays the same the
temperature increases with pressure.

Gay-Lussac’s Formula
P1/T1= P2/ T2
Where:
P1 = initial pressure
T1 = initial temperature
P2= final pressure
T2= final temperature
EXAMPLE:
P1 = 108 L
T1 = 41 ◦C
P2= ?
T2= 22◦C
108 kpa x 22◦C = .57.9 kpa
41 ◦C

DEMONSTRATION
GAY-LUSSAC’S LAW
WATCH AND LEARN
http://www.youtube.com/watch?v=j1yFvlgHM9Y

Unit 5 & 6 Test Review
Heat of Fusion
q = m ∙ΔH f
Specific Heat
c = ____q____
m ∙ ΔT ᵒC
Heat of Fusion of Water
334 J/g = 80 cal/g
Specific Heat of Water
4.18 J/g ∙ᵒC
Heat of fusion is the amount of heat energy required to change the state of a substance from solid to
liquid. This example problem demonstrates how to calculate the amount of energy required to melt a
sample of water ice.
With Graham's Law, you can find the effusion rates for two gases or the molecular mass of a gas.
This ratio of effusion rates follows the pattern that the gas with the lesser molecular mass has a
greater rate of effusion.
Ideal Gas Law
Combined Gas Law
p∙v=n∙R∙T V 1 ∙P 1 = V 2 ∙P 2
T 1 T 2
R = .082 atm or 8.31kpa
Boyle’s Law
P 1 ∙V 1 = P 2 ∙V 2
Charles’ Law
Gay – Lussac’s Law
V 1 = V 2 P 1 = P 2
T 1 T 2 T 1 T 2
Molarity
M = moles
L
Gas Solubility
S 1 = S 2
P 1 P 2

Unit 5
Chapter 13 States of Matter
Chapter 14 The Behavior of Gases
Chapter 15 Water and Aqueous Systems
Unit 6
Chapter 16 Solutions
Chapter 17 Thermochemistry
Chapter 18 Reaction Rates and Equilibrium
Chapter 19 Acid and Bases
Learning Goals
The students will learn what are the factors that determine and characteristics that distinguish gases
liquids and solids and how substances change from one state to another.
The students will learn how gases respond to changes in pressure, volume, and temperature and why
the ideal gas law is useful even though ideal gases do not exist.
The students will learn how the interactions between water molecules account for the unique
properties of water and how aqueous solutions form.
The student will learn how energy is converted in a chemical or physical process and how to
determine the amount of energy is absorbed or released in that process.