You also want an ePaper? Increase the reach of your titles
YUMPU automatically turns print PDFs into web optimized ePapers that Google loves.
Juan Roque<br />
Chemistry notebook
Chapter 1<br />
Unit 1<br />
Introduction to Chemistry<br />
The students will learn why and how to solve problems using<br />
chemistry.<br />
Identify what is science, what clearly is not science, and what superficially<br />
resembles science (but fails to meet the criteria for science).<br />
Students will identify a phenomenon as science or not science.<br />
Science<br />
Observation<br />
Inference<br />
Hypothesis<br />
Identify which questions can be answered through science and which<br />
questions are outside the boundaries of scientific investigation, such as<br />
questions addressed by other ways of knowing, such as art, philosophy, and<br />
religion.<br />
Students will differentiate between problems and/or phenomenon that can and<br />
those that cannot be explained or answered by science.<br />
Students will differentiate between problems and/or phenomenon that can and<br />
those that cannot be explained or answered by science.<br />
Observation<br />
Inference<br />
Hypothesis<br />
Theory<br />
Controlled experiment<br />
Describe how scientific inferences are drawn from scientific observations<br />
and provide examples from the content being studied.<br />
Students will conduct and record observations.<br />
Students will make inferences.<br />
Students will identify a statement as being either an observation or inference.<br />
Students will pose scientific questions and make predictions based on<br />
inferences.<br />
Inference<br />
Observation<br />
Hypothesis<br />
Controlled experiment<br />
Identify sources of information and assess their reliability according to the<br />
strict standards of scientific investigation.<br />
Students will compare and assess the validity of known scientific information<br />
from a variety of sources:
Print vs. print<br />
Online vs. online<br />
Print vs. online<br />
Students will conduct an experiment using the scientific method and compare<br />
with other groups.<br />
Controlled experiment<br />
Investigation<br />
Peer Review<br />
Accuracy<br />
Precision<br />
Percentage Error<br />
Chapter 2<br />
Matter and Change<br />
The students will learn what properties are used to describe<br />
matter and how matter can change its form.<br />
Differentiate between physical and chemical properties and physical and<br />
chemical changes of matter.<br />
Students will be able to identify physical and chemical properties of various<br />
substances.<br />
Students will be able to identify indicators of physical and chemical changes.<br />
Students will be able to calculate density.<br />
mass<br />
physical property<br />
volume<br />
chemical property<br />
vapor<br />
extensive property<br />
Chapter 3<br />
mixture<br />
intensive property<br />
solution<br />
element<br />
compound<br />
Scientific Measurements<br />
The students will be able to solve conversion problems using<br />
measurements.<br />
Determine appropriate and consistent standards of measurement for the<br />
data to be collected in a survey or experiment.<br />
Students will participate in activities to collect data using standardized<br />
measurement.<br />
Students will be able to manipulate/convert data collected and apply the data<br />
to scientific situations.<br />
Scientific notation<br />
International System of Units (SI)<br />
Significant figures<br />
Accepted value<br />
Experimental value<br />
Percent error<br />
Dimensional analysis
Unit 3<br />
Chapter 25 Nuclear Chemistry<br />
The students will learn what happens when an unstable<br />
nucleus decays and how nuclear chemistry affects their lives.<br />
Explore the theory of electromagnetism by comparing and contrasting the<br />
different parts of the electromagnetic spectrum in terms of wavelength,<br />
frequency, and energy, and relate them to phenomena and applications.<br />
Students will be able to compare and contrast the different parts of the<br />
electromagnetic spectrum.<br />
Students will be able to apply knowledge of the EMS to real world phenomena.<br />
Students will be able to quantitatively compare the relationship between energy,<br />
wavelength, and frequency of the EMS.<br />
amplitude<br />
wavelength<br />
frequency<br />
hertz<br />
electromagnetic radiation<br />
photon<br />
Planck’s constant<br />
Explain and compare nuclear reactions (radioactive decay, fission and<br />
fusion), the energy changes associated with them and their associated<br />
safety issues.<br />
Students will be able to compare and contrast fission and fusion reactions.<br />
Students will be able to complete nuclear decay equations to identify the type of<br />
decay.<br />
Students will participate in activities to calculate half-life.<br />
radioactivity<br />
nuclear radiation<br />
alpha particle<br />
beta particle<br />
gamma ray<br />
positron<br />
½ life<br />
transmutation<br />
fission<br />
fusion<br />
50
Chapter 7<br />
Ionic and Metallic Bonding<br />
The students will learn how ionic compounds form and how<br />
metallic bounding affects the properties of metals.<br />
Compare the magnitude and range of the four fundamental forces<br />
(gravitational, electromagnetic, weak nuclear, strong nuclear).<br />
Students will compare/contrast the characteristics of each fundamental force.<br />
gravity<br />
electromagnetic<br />
strong<br />
weak<br />
Distinguish between bonding forces holding compounds together and other<br />
attractive forces, including hydrogen bonding and van der Waals forces.<br />
Students will be able to compare/contrast traits of ionic and covalent bonds.<br />
Students will be able to compare/contrast basic attractive forces between<br />
molecules.<br />
Students will be able to predict the type of bond or attractive force between<br />
atoms or molecules.<br />
ionic bond<br />
covalent bond<br />
metallic bond<br />
polar covalent bond<br />
hydrogen bond<br />
van der Waals forces<br />
London dispersion forces<br />
Chapter 8<br />
Covalent Bonding<br />
The students will learn how molecular bonding is different<br />
than ionic bonding and electrons affect the shape of a<br />
molecule and its properties.<br />
Interpret formula representations of molecules and compounds in terms of<br />
composition and structure.<br />
Students will be able to interpret chemical formulas in terms of # of atoms.<br />
Students will be able to differentiate between ionic and molecular compounds.<br />
Students will be able to list various VSEPR shapes and identify examples of<br />
each.<br />
Students will be able to predict shapes of various compounds.<br />
Molecule<br />
empirical formula<br />
Atom<br />
Electron<br />
Element<br />
Compound<br />
51
Name ____________________<br />
Juan Roque<br />
Go to the web site www.darvill.clara.net/emag<br />
1. Click on “How the waves fit into the spectrum” and fill in this table:<br />
>: look out for the<br />
RED words on the web site!<br />
Low __________, frequency Long wavelength<br />
High frequency, Short ______________<br />
wavelengths<br />
Radio Waves<br />
microwaves Infra-red Visible Light Ultra-Violet X-rays<br />
Gamma rays<br />
2. Click on “Radio waves”. They are used for _______________________<br />
communication<br />
3. Click on “Microwaves”. They are used for cooking, mobile _________, Wifi _______ speed cameras and _________. radar<br />
4. Click on “Infra-red”. These waves are given off by _____ hot _________. objects They are used for remote controls,<br />
cameras in police ____________ helicopters , and alarm systems.<br />
5. Click on “Visible Light”. This is used in ___ CDplayers and _______ laser printers, and for seeing where we’re going.<br />
6. “UV” stands for “ ________ Ultra ___________”. Violet This can damage the _________ retina in your eyes, and cause<br />
sunburn and even _______ skin cancer. Its uses include detecting forged ______ bank _______. notes<br />
7. X-rays are used to see inside people, and for _________ airport security.<br />
8. Gamma rays are given off by some ________________ radioactive substances. We can use them to kill ________ cancer cells,<br />
which is called R_______________ adiotherapy .<br />
9. My Quiz score is ____%. 94<br />
52
10. Name ________________________________<br />
Go to the web site www.darvill.clara.net/emag<br />
Name How they’re made Uses Dangers<br />
by number of astronomical processesUsed to kill cancer cells<br />
Gamma rays<br />
X-rays<br />
Ultra-violet<br />
Visible light<br />
Infra-red<br />
microwaves<br />
Radio waves<br />
very high-energy electrons are made<br />
When electrons strike a metal target<br />
electrons are liberated from heated<br />
filament and accelerated by high<br />
voltage towards trget.<br />
Given off by the sun in large<br />
quantities<br />
special lamps such as sun beds<br />
white ligth is actually made up of a<br />
whole range of colours, mixed together.<br />
Radiotherapy<br />
Kill microbes and used to sterilise food<br />
To see inside people<br />
airport security<br />
atronomers<br />
Sun tan and detecting forged bank<br />
notes<br />
kill microbes<br />
produces vitamin D<br />
food drug companies sterilise their products<br />
to see things<br />
CD and DVD players<br />
remote controls, video recorders<br />
below the visible red light<br />
heal sprot injuries<br />
anthing else that is given off by hot objects<br />
night sights<br />
extremely high frequency radio waves,<br />
made by various amount of transmitters<br />
Made by various types of transmitters,<br />
depending on the wavelength<br />
mobile phones<br />
wifi<br />
speed cameras<br />
radar<br />
mainly used for communication<br />
Can cause cell damage<br />
cancer<br />
mutations<br />
can cause cell damage<br />
cancer<br />
Large doses can damage<br />
retin in the eyes<br />
can cause sunburn and<br />
skin cancer<br />
can damage the retina<br />
of the eyes<br />
looking at the sun<br />
too much and you can<br />
overheat<br />
ovrexposure can cause<br />
cateracts<br />
large doses can cause<br />
cancer and disorders<br />
_____ Frequency _____ frequency,<br />
Short wavelength ______ Wavelength<br />
53
Learning Goal for this section:<br />
The students will learn what happens when an unstable nucleus decays and how nuclear chemistry affects thier lives.<br />
Notes Section:<br />
Alpha particles He 4 goes through alpha<br />
2 238<br />
92 Uranium ><br />
down by 2<br />
protons><br />
234<br />
90 Thorium<br />
+<br />
4<br />
2 He<br />
Beta Particle<br />
up by 1<br />
negative<br />
electron<br />
6p<br />
8n<br />
14<br />
6 Carbon<br />
converting nuetron to proton<br />
carbon 14 turns into nitrogen<br />
14<br />
Nitrogen + negative electron<br />
7<br />
positive<br />
electron<br />
positron<br />
5p<br />
3n<br />
8<br />
5 Be > 8<br />
Be + e+<br />
4<br />
4p<br />
4n<br />
Gamma Energy<br />
always associatied with alpha particle or beta<br />
Half-LIfe 100g start 20min<br />
50g 20min 50g of something else<br />
25 40min 75g of something else<br />
it decays and becomes something else<br />
12.5 60min 87.5g of something else<br />
still 100g but of something else<br />
0.75 120min<br />
93.75 of something else<br />
54
The Nucleus<br />
A typical model of the atom is called the Bohr Model, in<br />
honor of Niels Bohr who proposed the structure in 1913. The Bohr atom consists of a central nucleus<br />
composed of neutrons and protons, which is surrounded by electrons which “orbit” around the nucleus.<br />
Protons carry a positive charge of one and have a mass of about 1 atomic mass unit or amu (1 amu =1.7x10-<br />
27 kg, a very, very small number). Neutrons are electrically “neutral” and also have a mass of about 1 amu. In<br />
contrast electron carry a negative charge and have mass of only 0.00055 amu. The number of protons in a<br />
nucleus determines the element of the atom. For example, the number of protons in uranium is 92 and the<br />
number in neon is 10. The proton number is often referred to as Z.<br />
Atoms with different numbers of protons are called elements, and are arranged in the periodic table with<br />
increasing Z.<br />
Atoms in nature are electrically neutral so the number of electrons orbiting the nucleus equals the number of<br />
protons in the nucleus.<br />
Neutrons make up the remaining mass of the nucleus and provide a means to “glue” the protons in place.<br />
Without neutrons, the nucleus would split apart because the positive protons would repel each other. Elements<br />
can have nucleii with different numbers of neutrons in them. For example hydrogen, which normally only has<br />
one proton in the nucleus, can have a neutron added to its nucleus to from deuterium, ir have two neutrons<br />
added to create tritium, which is radioactive. Atoms of the same element which vary in neutron number are<br />
called isotopes. Some elements have many stable isotopes (tin has 10) while others have only one or two. We<br />
express isotopes with the nomenclature Neon-20 or 20 Ne 10, with twenty representing the total number of<br />
neutrons and protons in the atom, often referred to as A, and 10 representing the number of protons (Z).<br />
Alpha Particle<br />
Decay<br />
Alpha decay is a radioactive process in which a<br />
particle with two neutrons and two protons is<br />
ejected from the nucleus of a radioactive atom. The particle is identical to the nucleus of a helium atom.<br />
Alpha decay only occurs in very heavy elements such as uranium, thorium and radium. The nuclei of these<br />
atoms are very “neutron rich” (i.e. have a lot more neutrons in their nucleus than they do protons) which makes<br />
emission of the alpha particle possible.<br />
After an atom ejects an alpha particle, a new parent atom is formed which has two less neutrons and two less<br />
protons. Thus, when uranium-238 (which has a Z of 92) decays by alpha emission, thorium-234 is created<br />
(which has a Z of 90).<br />
Because alpha particles contain two protons, they have a positive charge of two. Further, alpha particles are<br />
very heavy and very energetic compared to other common types of radiation. These characteristics allow alpha<br />
particles to interact readily with materials they encounter, including air, causing many ionizations in a very short<br />
distance. Typical alpha particles will travel no more than a few centimeters in air and are stopped by a sheet of<br />
paper.<br />
55
Beta Particle Decay<br />
Beta decay is a radioactive process in which an electron is emitted from the nucleus of a radioactive<br />
atom Because this electron is from the nucleus of the atom, it is called a beta particle to distinguish it<br />
from the electrons which orbit the atom.<br />
Like alpha decay, beta decay occurs in isotopes which are “neutron rich” (i.e. have a lot more<br />
neutrons in their nucleus than they do protons). Atoms which undergo beta decay are located below<br />
the line of stable elements on the chart of the nuclides, and are typically produced in nuclear reactors.<br />
When a nucleus ejects a beta particle, one of the neutrons in the nucleus is transformed into a proton.<br />
Since the number of protons in the nucleus has changed, a new daughter atom is formed which has<br />
one less neutron but one more proton than the parent. For example, when rhenium-187 decays<br />
(which has a Z of 75) by beta decay, osmium-187 is created (which has a Z of 76). Beta particles<br />
have a single negative charge and weigh only a small fraction of a neutron or proton. As a result, beta<br />
particles interact less readily with material than alpha particles. Depending on the beta particles<br />
energy (which depends on the radioactive atom), beta particles will travel up to several meters in air,<br />
and are stopped by thin layers of metal or plastic.<br />
Positron emission or beta plus decay (β+ decay) is a subtype of radioactive decay called beta decay,<br />
in which a proton inside a radionuclide nucleus is converted into a neutron while releasing a positron<br />
and an electron neutrino (νe). Positron emission is mediated by the weak force.<br />
An example of positron emission (β+ decay) is shown with magnesium-23 decaying into sodium-23:<br />
23 Mg12 → 23 Na11 + e +<br />
Because positron emission decreases proton number relative to neutron number, positron decay<br />
happens typically in large "proton-rich" radionuclides. Positron decay results in nuclear transmutation,<br />
changing an atom of one chemical element into an atom of an element with an atomic number that is<br />
less by one unit.<br />
Positron emission should not be confused with electron emission or beta minus decay (β− decay),<br />
which occurs when a neutron turns into a proton and the nucleus emits an electron and an<br />
antineutrino.<br />
56
Gamma<br />
Radiation<br />
After a decay reaction, the nucleus is often in an<br />
“excited” state. This means that the decay has<br />
resulted in producing a nucleus which still has<br />
excess energy to get rid of. Rather than emitting another beta or alpha particle, this energy is lost by<br />
emitting a pulse of electromagnetic radiation called a gamma ray. The gamma ray is identical in<br />
nature to light or microwaves, but of very high energy.<br />
Like all forms of electromagnetic radiation, the gamma ray has no mass and no charge. Gamma rays<br />
interact with material by colliding with the electrons in the shells of atoms. They lose their energy<br />
slowly in material, being able to travel significant distances before stopping. Depending on their initial<br />
energy, gamma rays can travel from 1 to hundreds of meters in air and can easily go right through<br />
people.<br />
It is important to note that most alpha and beta emitters also emit gamma rays as part of their decay<br />
process. However, their is no such thing as a “pure” gamma emitter. Important gamma emitters<br />
including technetium-99m which is used in nuclear medicine, and cesium-137 which is used for<br />
calibration of nuclear instruments.<br />
Half Life<br />
Half-life is the time required for the quantity of a<br />
radioactive material to be reduced to one-half its<br />
original value.<br />
All radionuclides have a particular half-life, some<br />
of which a very long, while other are extremely<br />
short. For example, uranium-238 has such a<br />
long half life, 4.5x109 years, that only a small fraction has decayed since the earth was formed. In<br />
contrast, carbon-11 has a half-life of only 20 minutes. Since this nuclide has medical applications, it<br />
has to be created where it is being used so that enough will be present to conduct medical studies.<br />
57
The Learning Goal for this assignment is:<br />
Distinguish between bonding forces holding compounds together and other attractive forces, including hydrogen bonding<br />
and Van der Waals forces.<br />
Introduction to Ionic Compounds<br />
Those molecules that consist of charged ions with opposite charges are called IONIC. These ionic<br />
compounds are generally solids with high melting points and conduct electrical current. Ionic<br />
compounds are generally formed from metal and a non-metal elements. See Ionic Bonding below.<br />
Ionic Compound Example<br />
For example, you are familiar with the fairly benign unspectacular behavior of common white<br />
crystalline table salt (NaCl). Salt consists of positive sodium ions (Na + ) & negative chloride ions (Cl - ).<br />
On the other hand the element sodium is a silvery gray metal composed of neutral atoms which react<br />
vigorously with water or air. Chlorine as an element is a neutral greenish-yellow, poisonous, diatomic<br />
gas (Cl2).<br />
The main principle to remember is that ions are completely different in physical and chemical<br />
properties from the neutral atoms of the elements.<br />
The notation of the + and - charges on ions is very important as it conveys a definite meaning.<br />
Whereas elements are neutral in charge, IONS have either a positive or negative charge depending<br />
upon whether there is an excess of protons (positive ion) or excess of electrons (negative ion).<br />
Formation of Positive Ions<br />
Metals usually have 1-4 electrons in the outer energy level. The electron arrangement of a rare gas is<br />
most easily achieved by losing the few electrons in the newly started energy level. The number of<br />
electrons lost must bring the electron number "down to" that of a prior rare gas.<br />
How will sodium complete its octet?<br />
First examine the electron arrangement of the atom. The atomic number is eleven, therefore, there<br />
are eleven electrons and eleven protons on the neutral sodium atom. Here is the Bohr diagram and<br />
Lewis symbol for sodium:<br />
58
This analysis shows that sodium has only one electron in its outer level. The nearest rare gas is neon<br />
with 8 electron in the outer energy level. Therefore, this electron is lost so that there are now eight<br />
electrons in the outer energy level, and the Bohr diagrams and Lewis symbols for sodium ion and<br />
neon are identical. The octet rule is satisfied.<br />
Ion Charge?<br />
What is the charge on sodium ion as a result of losing one electron? A comparison of the atom and<br />
the ion will yield this answer.<br />
Sodium Atom<br />
Sodium Ion<br />
11 p+ to revert to 11 p + Protons are identical in<br />
12 n an octet 12 n<br />
the atom and ion.<br />
Positive charge is<br />
11 e- lose 1 electron 10 e-<br />
caused by lack of<br />
0 charge + 1 charge<br />
electrons.<br />
Formation of Negative Ions<br />
How will fluorine complete its octet?<br />
First examine the electron arrangement of the atom. The atomic number is nine, therefore, there are<br />
nine electrons and nine protons on the neutral fluorine atom. Here is the Bohr diagram and Lewis<br />
symbol for fluorine:<br />
This analysis shows that fluorine already has seven electrons in its outer level. The nearest rare gas<br />
is neon with 8 electron in the outer energy level. Therefore only one additional electron is needed to<br />
complete the octet in the fluorine atom to make the fluoride ion. If the one electron is added, the Bohr<br />
diagrams and Lewis symbols for fluorine and neon are identical. The octet rule is satisfied.<br />
59
Ion Charge?<br />
What is the charge on fluorine as a result of adding one electron? A comparison of the atom and the<br />
ion will yield this answer.<br />
Fluorine Atom Fluoride Ion *<br />
9 p+ to complete 9 p + Protons are identical in<br />
10 n octet 10 n<br />
9 e- add 1 electron 10 e-<br />
0 charge - 1 charge<br />
the atom and ion.<br />
Negative charge is<br />
caused by excess<br />
electrons<br />
* The "ide" ending in the name signifies a simple negative ion.<br />
Summary Principle of Ionic Compounds<br />
An ionic compound is formed by the complete transfer of electrons from a metal to a nonmetal and<br />
the resulting ions have achieved an octet. The protons do not change. Metal atoms in Groups 1-3<br />
lose electrons to non-metal atoms with 5-7 electrons missing in the outer level. Non-metals gain 1-4<br />
electrons to complete an octet.<br />
Octet Rule<br />
Elemental atoms generally lose, gain, or share electrons with other atoms in order to achieve the<br />
same electron structure as the nearest rare gas with eight electrons in the outer level.<br />
The proper application of the Octet Rule provides valuable assistance in predicting and explaining<br />
various aspects of chemical formulas.<br />
Introduction to Ionic Bonding<br />
Ionic bonding is best treated using a simple<br />
electrostatic model. The electrostatic model<br />
is simply an application of the charge<br />
principles that opposite charges attract and<br />
similar charges repel. An ionic compound<br />
results from the interaction of a positive and<br />
negative ion, such as sodium and chloride in<br />
common salt.<br />
The IONIC BOND results as a balance<br />
between the force of attraction between<br />
opposite plus and minus charges of the ions<br />
and the force of repulsion between similar<br />
negative charges in the electron clouds. In<br />
crystalline compounds this net balance of<br />
forces is called the LATTICE ENERGY.<br />
Lattice energy is the energy released in the<br />
formation of an ionic compound.<br />
DEFINITION: The formation of an IONIC<br />
BOND is the result of the transfer of one or<br />
more electrons from a metal onto a nonmetal.<br />
60
Metals, with only a few electrons in the outer energy level, tend to lose electrons most readily. The<br />
energy required to remove an electron from a neutral atom is called the IONIZATION POTENTIAL.<br />
Energy + Metal Atom ---> Metal (+) ion + e-<br />
Non-metals, which lack only one or two electrons in the outer energy level have little tendency to lose<br />
electrons - the ionization potential would be very high. Instead non-metals have a tendency to gain<br />
electrons. The ELECTRON AFFINITY is the energy given off by an atom when it gains electrons.<br />
Non-metal Atom + e- --- Non-metal (-) ion + energy<br />
The energy required to produce positive ions (ionization potential) is roughly balanced by the energy<br />
given off to produce negative ions (electron affinity). The energy released by the net force of<br />
attraction by the ions provides the overall stabilizing energy of the compound.<br />
Notes Section:<br />
61
The Learning Goal for this assignment is:<br />
Introduction to Covalent Bonding:<br />
Bonding between non-metals consists of two electrons shared between two atoms. Using the Wave<br />
Theory, the covalent bond involves an overlap of the electron clouds from each atom. The electrons<br />
are concentrated in the region between the two atoms. In covalent bonding, the two electrons shared<br />
by the atoms are attracted to the nucleus of both atoms. Neither atom completely loses or gains<br />
electrons as in ionic bonding.<br />
There are two types of covalent bonding:<br />
1. Non-polar bonding with an equal sharing of electrons.<br />
2. Polar bonding with an unequal sharing of electrons. The number of shared electrons depends on<br />
the number of electrons needed to complete the octet.<br />
NON-POLAR BONDING results when two identical non-metals equally share electrons between<br />
them. One well known exception to the identical atom rule is the combination of carbon and hydrogen<br />
in all organic compounds.<br />
Hydrogen<br />
The simplest non-polar covalent molecule is hydrogen. Each hydrogen<br />
atom has one electron and needs two to complete its first energy level.<br />
Since both hydrogen atoms are identical, neither atom will be able to<br />
dominate in the control of the electrons. The electrons are therefore<br />
shared equally. The hydrogen covalent bond can be represented in a<br />
variety of ways as shown here:<br />
The "octet" for hydrogen is only 2 electrons since the nearest rare gas is<br />
He. The diatomic molecule is formed because individual hydrogen atoms<br />
containing only a single electron are unstable. Since both atoms are<br />
identical a complete transfer of electrons as in ionic bonding is<br />
impossible.<br />
Instead the two hydrogen atoms SHARE both electrons equally.<br />
Oxygen<br />
Molecules of oxygen, present in about 20% concentration in air are<br />
also covalent molecules. See the graphic on the left of the Lewis Dot<br />
Structure.<br />
There are 6 electrons in the outer shell, therefore, 2 electrons are<br />
needed to complete the octet. The two oxygen atoms share a total of<br />
four electrons in two separate bonds, called double bonds.<br />
The two oxygen atoms equally share the four electrons.<br />
62
POLAR BONDING results when two different non-metals unequally share electrons between them.<br />
One well known exception to the identical atom rule is the combination of carbon and hydrogen in all<br />
organic compounds.<br />
The non-metal closer to fluorine in the Periodic Table has a greater tendency to keep its own electron<br />
and also draw away the other atom's electron. It is NOT completely successful. As a result, only<br />
partial charges are established. One atom becomes partially positive since it has lost control of its<br />
electron some of the time. The other atom becomes partially negative since it gains electron some of<br />
the time.<br />
Hydrogen Chloride<br />
Hydrogen Chloride forms a polar covalent molecule. The graphic<br />
on the left shows that chlorine has 7 electrons in the outer shell.<br />
Hydrogen has one electron in its outer energy shell. Since 8<br />
electrons are needed for an octet, they share the electrons.<br />
However, chlorine gets an unequal share of the two electrons,<br />
although the electrons are still shared (not transferred as in ionic<br />
bonding), the sharing is unequal. The electrons spends more of the<br />
time closer to chlorine. As a result, the chlorine acquires a "partial"<br />
negative charge. At the same time, since hydrogen loses the<br />
electron most - but not all of the time, it acquires a "partial" charge.<br />
The partial charge is denoted with a small Greek symbol for delta.<br />
Water<br />
Water, the most universal compound on all of the earth, has the property of<br />
being a polar molecule. As a result of this property, the physical and<br />
chemical properties of the compound are fairly unique.<br />
Dihydrogen Oxide or water forms a polar covalent molecule. The graphic on<br />
the left shows that oxygen has 6 electrons in the outer shell. Hydrogen has<br />
one electron in its outer energy shell. Since 8 electrons are needed for an<br />
octet, they share the electrons.<br />
Notes Section:<br />
metal and nonmetal make a ionic bond<br />
metal and metal make a covalent bond<br />
left side metal right side nonmetal<br />
C^2H^6O<br />
flourine most electronegative and hydrogen is right above it.<br />
1) C2 x 4 8<br />
H^6 x 1 6<br />
O1 x 6 8 = 25<br />
5)<br />
2) C^2 x 8 16<br />
H^6 x 2 12<br />
O^1 x 8 8 = 36<br />
3) 36-20=16e-<br />
4) 2/ 16= 8 bonds<br />
63
C 2 H 6 O Ethanol CH 3 CH 2 O<br />
Step 1<br />
Find valence e- for all atoms. Add them together.<br />
C: 4 x 2 = 8<br />
H: 1 x 6 = 6<br />
O: 6<br />
Total = 20<br />
Step 2<br />
Find octet e- for each atom and add them together.<br />
C: 8 x 2 = 16<br />
H: 2 x 6 = 12<br />
O: 8<br />
Total = 36<br />
Step 3<br />
Subtract Step 1 total from Step 2.<br />
Gives you bonding e-.<br />
36 – 20 = 16e-<br />
Step 4<br />
Find number of bonds by diving the number in step 3 by 2<br />
(because each bond is made of 2 e-)<br />
16e- / 2 = 8 bond pairs<br />
These can be single, double or triple bonds.<br />
Step 5<br />
Determine which is the central atom<br />
Find the one that is the least electronegative.<br />
Use the periodic table and find the one farthest<br />
away from Fluorine or<br />
The one that only has 1 atom.<br />
64
Step 6<br />
Put the atoms in the structure that you think it will<br />
have and bond them together.<br />
Put Single bonds between atoms.<br />
Step 7<br />
Find the number of nonbonding (lone pairs) e-.<br />
Subtract step 3 number from step 1.<br />
20 – 16 = 4e- = 2 lone pairs<br />
Step 8<br />
Complete the Octet Rule by adding the lone<br />
pairs.<br />
Add any left over bonds to make double or triple<br />
bonds.<br />
Then, if needed, use any lone pairs to make<br />
double or triple bonds so that all atoms meet<br />
the Octet Rule.<br />
See Step 4 for total number of bonds.<br />
Step 9<br />
Find the formal charges for the atoms in the compound.<br />
Arrange atoms so that all formal charges<br />
are as close to 0 as possible.<br />
Some central atoms do not meet the octet rule.<br />
Boron can sometimes have only 6 electrons and<br />
some elements in Periods 3—7 may exceed the<br />
octet rule.<br />
65
Linear<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp AX 2 none 180<br />
BeCl 2<br />
Berylium<br />
dicloride<br />
Cl<br />
Be<br />
Cl<br />
element bond lone pair<br />
C<br />
66
Trigonal planal<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 2 AX 3 none 60<br />
BF 3<br />
F<br />
B<br />
Boron triflouride<br />
F<br />
F<br />
element bond lone pair<br />
C<br />
67
Trigonal planar<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 3 AX 2 E 2 2 104<br />
OF 2<br />
Oxygen diflouride<br />
O<br />
F<br />
F<br />
element bond lone pair<br />
C<br />
68
Tetrahedral<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 3 AX 4 none 109.5<br />
PO 4<br />
3-<br />
Phosphate<br />
O<br />
O<br />
P<br />
5-5=0<br />
O<br />
6-6=0<br />
O<br />
6-7=1-<br />
element bond lone pair<br />
C<br />
69
Trigonal Pyramidal<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 3 AX 3 E 2 104.5<br />
PH 3<br />
Phosphorous trihydride<br />
H<br />
P<br />
H<br />
H<br />
element bond lone pair<br />
C<br />
70
linear<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp AX 2 2 180<br />
O 3<br />
trioxide<br />
O<br />
O<br />
O<br />
element bond lone pair<br />
C<br />
71
Trigonal Bi pyramidal<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 3 d AX 5 none 120/90<br />
PCl 5<br />
Phosphorous<br />
pentacloride<br />
Cl<br />
Cl<br />
Cl<br />
P<br />
Cl<br />
Cl<br />
element bond lone pair<br />
C<br />
C<br />
72
T-shaped<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 3 d AX 3 E 2 2 90<br />
ClF 3<br />
Chlorine trifloride<br />
F<br />
Cl<br />
F<br />
F<br />
element bond lone pair<br />
C<br />
73
Octahedral<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 3 d 2 AX 6 none 90<br />
SF 6<br />
Sulfur hexafluoride<br />
F<br />
F<br />
F<br />
S<br />
F<br />
F<br />
F<br />
element bond lone pair<br />
74
Square Planar<br />
Molecular Geometry<br />
Orbital Equation Lone Pairs Angle<br />
sp 3 d 2 AX 4 E 2 2 90<br />
ICl 4<br />
-<br />
Iodine tetrafluoride<br />
ion<br />
Cl<br />
Cl<br />
I<br />
Cl<br />
Cl<br />
element bond lone pair<br />
C<br />
75
Orbitals Equation Lone Pairs<br />
Angle<br />
Name<br />
Sp 3 d AX5 none 120/90 Phosphorous Penta Chloride<br />
Sp 3 d 2 AX6 none 90<br />
Sulfur hexafluoride<br />
Sp 3 d AX3E2 2 90<br />
Chlorine trifluoride<br />
sp AX2 2 180<br />
Trioxide<br />
Sp 3 AX3E 2 104.5<br />
Phosphorous trihydride<br />
Sp 3 AX4 none 109.5<br />
Phosphorous tetroxide<br />
Sp 3 d 2 AX4E2 2 90<br />
Iodine tetrafluoride ion<br />
Sp 3 AX2E2 2 104<br />
Oxygen difluoride<br />
Sp 2 AX3 none 60<br />
Boron trifluoride<br />
sp AX2 none 180<br />
Beryllium dichloride<br />
78
Name Formula Charge<br />
Dichromate Cr₂O₇ 2-<br />
Sulfate SO₄ 2-<br />
Hydrogen Carbonate HCO₃ 1-<br />
Hypochlorite ClO 1-<br />
Phosphate PO₄ 3-<br />
Nitrite NO₂ 1-<br />
Chlorite ClO₂ 1-<br />
Dihydrogen phosphate H₂PO₄ 1-<br />
Chromate CrO₄ 2-<br />
Carbonate CO₃ 2-<br />
Hydroxide OH 1-<br />
Hydrogen phosphate HPO₄ 2-<br />
Ammonium NH₄ 1+<br />
Acetate C₂H₃O₂ 1-<br />
Perchlorate ClO₄ 1-<br />
Permanganate MnO₄ 1-<br />
Chlorate ClO₃ 1-<br />
Hydrogen Sulfate HSO₄ 1-<br />
Phosphite PO₃ 3-<br />
Sulfite SO₃ 2-<br />
Silicate SiO₃ 2-<br />
Nitrate NO₃ 1-<br />
Hydrogen Sulfite HSO₃ 1-<br />
Oxalate C₂O₄ 2-<br />
Cyanide CN 1-<br />
Hydronium H₃O 1+<br />
Thiosulfate S₂O₃ 2-<br />
77
Chapter 9<br />
Unit 4<br />
Chemical Names and Formulas<br />
The students will learn how the periodic table helps them<br />
determine the names and formulas of ions and compounds.<br />
Chapter 22 Hydrocarbon Compounds<br />
The student will learn how Hydrocarbons are named and the<br />
general properties of Hydrocarbons.<br />
Describe how different natural resources are produced and how their rates<br />
of use and renewal limit availability.<br />
Students will explore local, national, and global renewable and nonrenewable<br />
resources.<br />
Students will explain the environmental costs of the use of renewable and<br />
nonrenewable resources.<br />
Students will explain the benefits of renewable and nonrenewable resources.<br />
Nuclear reactors<br />
Natural gas<br />
Petroleum<br />
Refining<br />
Coal<br />
78
Chapter 23 Functional Groups<br />
The student will learn what effects functional groups have on<br />
organic compounds and how chemical reactions are used in<br />
organic compounds.<br />
Describe the properties of the carbon atom that make the diversity of carbon<br />
compounds possible.<br />
Identify selected functional groups and relate how they contribute to<br />
properties of carbon compounds.<br />
Students will identify examples of important carbon based molecules.<br />
Students will create 2D or 3D models of carbon molecules and explain why this<br />
molecule is important to life.<br />
covalent bond<br />
single bond<br />
double bond<br />
triple bond<br />
monomer<br />
polymer<br />
79
http://www.bbc.co.uk/education/guides/zm9hvcw/revision<br />
students will learn what effects functional groups have on<br />
organic compounds and how chemical reactinos are used in<br />
organic compounds.<br />
A homologous series is a family of hydrocarbons with similar<br />
chemical properties who share the same genral formula.<br />
The general formula of the alkanes is CnH2n+2<br />
methane- (natural gases) cooking, heating<br />
propane- used in gas cylinders for BBQ etc<br />
octane- used in petrol for cars.<br />
The longest unbranched chain containing the functional group is th<br />
parent molecule, or simply the longest unbranched chain for alkane<br />
rememeber that, the longest chain can go round a bend.<br />
80
5 with<br />
double bond<br />
2 without<br />
81