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Chapters 7, 8, 9 notes
________________ gave us a rough idea of what an atom looks like, protons and ____________ in a
nucleus, and ___________________ orbiting the nucleus. But what does an atom look like? Where are the
electrons located in an atom? This is an important topic of discussion because where electrons are
determines the _________________ properties of matter. Amazingly, the answer to the above questions
comes from studying ________________.
Around 1670, Isaac ______________ said that light is made up of distinct _____________. He used a
_______ to show that light breaks up into ________________________, which can reform into ________
light if a second prism is used. Christiaan _______________, a Dutch scientist, said that light is not made
up of particles (Newton called them corpuscles), but is a series of ________________ in 1680. Huygens
was able to show that Newton’s prism experiment was really just wave refraction. Huygens showed that
light is a wave because light __________________. Diffraction is the _______________ out of waves after
they encounter an __________________________________. [picture]
Light comes in many different forms. The shortest wavelength to longest wavelengths include
____________ rays, ________,
ultraviolet rays, visible light (which can be broken into ___________), infrared light, microwaves, radio
waves, TV waves. All together, these forms of light make up the _______________________________,
which can be found on p. 293 in the textbook. Wavelength is defined as the __________________ at which
a wave repeats its cycle. The symbol we use is the Greek letter, lambda, . The units for lambda can be any
length unit. In equations, we have to use meters, m, but we could use cm, mm, km, nm, or Ao
[ recall 1x1010 Ao = 1 m]. Different forms of light are different because they have different ___________.
Waves also have a __________________, which we define as the _______________ of waves that pass a
fixed point per second. The symbol for frequency is the Greek letter, nu, . The units for frequency
include: 1/sec, sec-1, hertz (symbol Hz), MHz, KHz, or cycles per second, cps. In equations, we use sec -1.
A third property of waves is speed. In a vacuum, light travels at 2.9979 x 10 8 m/s. The symbol for speed is a
lower case, c. So for any wave, c =  . Note: wavelength and frequency are _______________
proportional, since they multiply to equal a constant. Thus, the electromagnetic spectrum could also be
separated by ________________ as well as by wavelength. Gamma rays have a ______ frequency, TV
waves have a _____ frequency.
TTP 1] Find the frequency of red light that has a wavelength of 6200 Ao.
2] Find the wavelength of a radio signal that is 99.1 MHz
3] Find the frequency of an X-ray with a wavelength of 0.75 nm.
4] Find the wavelength in angstroms if the frequency is 2.00 x 10 18 Hz.
So for hundreds of years, light was thought of as a wave. One problem did persist. How was it that light can
travel through a ___________? Scientists of the era said that the universe had an ether that allowed light to
travel through a vacuum. It would seem that all energy is a _________ since it can be divided into smaller
pieces, but matter was made up of ______________. But in 1900, Max ____________, a German scientist,
showed that energy comes in discreet bundles called _____________. He is the founder of
______________ theory, which says that energy comes to us in discreet chunks or packets called
____________ (singular:________)
The size of one packet can be computed by the equation:
where E is energy in joules (joules are a kilogram per m2-s2 - the hyphen separates units of multiplied
variables);  is the frequency in s-1. h is Planck’s constant, and its value is: __________.
Note, energy and frequency are ______________ proportional. X-rays have a ______ frequency and a
________ amount of energy. More vocabulary: a particle of light is called a ____________.
An _______________ is a mole of photons. A mole is 6.022 x 10 23 particles. This big number is called
Avogadro’s number.
TTP 1] Find the energy of a radio wave with a frequency of 1400 kHz?
2] Find the frequency of a photon if its energy is 4.56 x 10 -19 joules.
3] Find the energy of one photon of blue light. The wavelength is 4600 angstroms. Find the energy of one
mole of these blue photons.
4] Find the wavelength of light in angstroms if its energy is 4.77 x 10 -19 joules in one photon. Is the light
visible, infrared, or ultraviolet?
5] Find the wavelength of light in angstroms if its energy is 6.6 x 10 5 kj/mol of photons.
In 1905, _______________ explained the ________________ effect. The ___________________ effect is
when light of sufficiently short ________________ hits a ____________ surface, electrons are ejected off
the surface of the ______________. Einstein showed that this proves that light is a _______________.
Light is a wave because it ________________. Particles do not diffract. But light is a particle because they
can hit electrons, a particle, with sufficient force to make them eject from the nucleus of an atom. If light
were only a wave, the wave energy would build up and any type of light could cause the ______________
effect. Thus light is particle, wave, or wavicle.
In 1923, a Frenchman named Louis _____________ used three equations, c =  ; E = h; and from
Einstein’s theory of relativity, E = mc2, to show that matter is both a wave and a particle just like light. The
derivation is as follows.
From this we have the deBroglie wave equation:
where lambda,  is the wavelength in meters, h is Planck’s constant, 6.626 x 10-34 j-s; m is the mass in
kilograms, and v is the velocity in meters per second.
TTP 1] Find the wavelength of an electron traveling at 10.0% of the speed of light in angstroms. The mass
of an electron is 9.11 x 10-31 kg.
2] Find the wavelength of a ball, mass 0.10 kg traveling at 45 m/sec.
What the above problems show is that quantum effects for particles are only significant for very low _____
which means masses of molecules or smaller, or for particles that move near the speed of ____________.
This also means that energy and matter really are equivalent. Energy is a dilute form of __________ and
matter is a concentrated form of ______________.
In 1923, two Americans named _______________ and ______________ showed that matter does diffract
according to the deBroglie wave equation. They did this by shooting electrons into a ________________.
The photographic plate on the opposite side showed a ____________ pattern, exactly as deBroglie would
predict.
In 1913, a Danish physicist named ____________ finally connected all these ideas into an atomic structure.
Since energized atoms give off light of specific colors, the wavelengths of light they give off are ________.
Specific wavelengths are specific frequencies (see _________). Specific energies means the electrons that
are energized when energy is put into them have to be in specific ____________. The normal resting orbit
is called the _____________ state. Energized orbits, which are any above the ground state, are called the
______________ state. Light is given off when an electron drops from an _____________ state down to
the ______________ state. The bigger the physical drop, the more energy, so the higher the __________,
and the shorter the ________________ of light. So for Bohr, electrons orbit the ____________ like planets
orbit the sun in our solar system. Bohr successfully wrote an equation that describes the energy of light that
will be given off when an electron drops between orbits for a hydrogen atom. It reads:
where R = the ionization energy of hydrogen, 2.178 x 10 -18 joules; nu is the upper level orbit the electron
drops from, and nl is the lower orbit the electron drops to. Note: for an electron to go from the ground state
to an __________, energy must be _____________. The same equation can be used to predict the energy
added to move an electron from a lower orbit to a higher orbit. When electrons drop levels, the energy is
emitted in the form of __________. When energy is given off, we note this with a negative sign. There is no
such thing as negative energy, but we use the negative sign to indicate whether energy is absorbed (it’s
positive) or emitted (then it’s negative).
For any one electron state, the Bohr equation becomes:
The Bohr model was a significant discovery. It married ____________ theory to atomic structure. We call
this _____________ mechanics. Unfortunately, his equations didn’t work for any other element than
hydrogen. There were other problems with his model as well. So a more sophisticated model had to be
developed. In the mid-1920’s Werner _______________ in Germany, and Erwin ___________________
in Switzerland developed a better model of the atom. Heisenberg’s model used matrix calculus and was
very complicated. Schrodinger’s equation was a bit simpler. With this equation (do not attempt to
memorize), we can predict the probability of locating an electron in space:
the Greek letter psi, , is the wave operator function. It has no physical significance. But 2 is the
probability of locating an electron in space.
In the meantime, we can do problems like the following:
1. Find the energy, frequency and wavelength of a photon of light emitted when an electron in a hydrogen
atom drops from the sixth to the second energy level. Determine if the light is visible, uv or infrared.
2.
Find the energy, frequency, and wavelength of a photon of light emitted when an electron in a hydrogen
atom drops from the fifth to the third energy level. Determine if the light is visible, uv, or infrared.
3.
The wavelength of light is 486.4 nm. If the light came from energized hydrogen, and the electron
dropped to the second energy level, what was the energy level from which the electron began?
Electrons are waves (because they ______________), and they are particles because they have mass. If
electrons are waves, what type of waves are they? The answer is they are _________________ waves, a
series of _________________ waves that reinforce one another. There is an important aspect to this. It was
stated first by Werner _____________ in 1927. This is called the _____________________________.
What this means is that you cannot precisely say exactly where an electron is at a given moment in time. We
simply cannot state how the electron is moving around the nucleus because it is a _____ and a ________ at
the same time. We can give the most probabilistic region of space over which the electron can exist. We call
this the _______________. The mathematics of the Heisenberg Uncertainty Principle is as follows:
where x = the extent to which I can determine the location of a particles location measured in meters
v = the extent to which I can determine the velocity of a particle measured in m/s
h = Planck’s constant, 6.626 x 10-34 j-s; m is the mass in kilograms
The Heisenberg Uncertainty principle has some important corollaries: 1] the more you try to define the
____ of a particle, the more you change its _____________. 2] you cannot make a measurement to an
infinite number of ______________________. 3] absolute certitude is impossible
Please appreciate that measurements are still possible. Knowledge is real. But there are definable limits to
our measurements, and limits to our knowledge. The Heisenberg Uncertainty Principle does not say I’m
uncertain about everything.
Sample problems: 1] Calculate the uncertainty in location of an electron if v = 0.200 m/s. The mass of an
electron is 9.11 x 10-31 kg. Compare this to the size of an atom.
2] Calculate the uncertainty in location of a tennis ball (mass is 56.7 g) if v = 0.200 m/s. Compare this to
the size of a tennis ball.
Returning back to our standing waves, there is a point in a standing wave where there is no lateral motion.
This point is called the ___________. The more nodes, the more __________ the standing wave has. The
more energy an electron has, the ____________ away it can be from the nucleus. Because we can solve the
Schrodinger equation, we can state the most probable location for an electron. What we find is that
electrons make four distinct _____________ called electron orbitals. They are called:
Writing electron configurations in standard notation:
1H
2He
8O
9F
3Li
4Be
10Ne
5B
11Na
6C
7N
+3
13Al
7N
-3
Since writing electron configurations can take a very long time, we have a way to make the process a bit
shorter, called the ____________________ shorthand method. In brackets, you write the inert gas that has
the closest, but less than, atomic number to the element, then the rest of the electron configuration.
There are some factors to consider when writing electron configurations. The s sublevel begins with 1. The
p sublevel begins with _____. The d sublevel begins with _____. The f sublevel begins with _____. For
elements 58-71 (the __________________) and elements 90-103 (the _________________) collectively,
the lanthanides and the actinides are called the _______________ elements, you cover the d 1 before you
write the f electrons. But for elements 72-86, and 104-118, you skip the d1, and go from the s to the f’s, and
then to the d’s and/or p’s. Follow the sample problems below:
Write the electron configuration in standard notation for: (always use core notation unless I direct you to
write the entire electron configuration)
13Al
19K
79Au
26Fe
86Rn
63Eu
106Sg
50Sn
56Ba
44Ru
98Cf
17Cl
40Zr
104Rf
70Yb
There is a second way to describe electron configuration. It’s called _________________ notation, or circle
notation. The advantage of this notation is that it can identify the number of _____________ electrons,
which are important because they are involved in bonding, and determine the _____________ properties of
an element. The disadvantage is that they are cumbersome to write. The key to writing electron
configurations in orbital notation is that one circle is used for the s sublevel, _____ circles are used for p,
______ circles are used for d, and _____ circles are used for the f sublevel. Each circle represents an
__________, a region of space over which there is a probability of locating an electron. There are three
orbitals of p because I can be a dumbbell shape in three directions, px, py, and pz. There are 5 ways to be a
cloverleaf, d. There are 7 different orientations of f’s. But a sphere (s), can only have one orientation.
Arrows represent ___________. Up arrows indicate one ________, and a down arrows represents the
opposite ________. Only two electrons can occupy an _____________, and they must have opposite ___.
That’s because spinning electrons create a ___________ field. The opposite spins create opposite magnetic
fields which are stronger than the ___________________ repulsion that two electrons create. For multiorbital sublevels (p,d and/or f), electrons fill in with the same ______ in separate orbitals before doubling
up with the opposite spin. This is called __________ rule.
Write the electron configuration in orbital notation for:
8O
20Ca
23V
46Pd
61Pm
54Xe
For the above elements, identify the number of unpaired electrons.
There is a third way to describe electron configuration. We can use _____________________. There are
_______ quantum numbers. The first is called the principal quantum number, ____, which represents the
number of ___________ in the standing wave, relative energy, and relative distance from the _________.
n can equal 1, 2,3, to  . The periodic table stops for n = 7, but there is no real upper limit. The second
quantum number is _______ which is the sublevel, as it indicates the shape.
l = 0 is s;
l = 1 is p;
l = 2 is d;
l = 3 is f; and l = 4 is g (there is no g sublevel on the periodic
table, but it does exist in the _____ state. The book makes the point that the quantum numbers characterize
possible properties only for the _____________ atom. They could also characterize the highest energy
electron in the _____________ state for any possible element. The possible values of l depend on n. For a
given value of n, l = n-1. l is called the angular momentum number. The third quantum number, ml,
indicates the __________ in space of the orbital. ml is called the magnetic quantum number.
If l = 0 (s), then ml = 0; If l = 1 (p), then ml = -1,0,+1. These correspond to px, py and pz.
If l = 2(d), then ml = -2, -1, 0, +1, +2. These correspond to the 5 d orbitals.
If l = 3 (f), then ml = -3, -2, -1, 0, +1, +2, +3. Note that ml   l. The maximum number of orbitals for a
given sublevel is 2 l + 1. The fourth quantum number is ms is the electron’s _________. The  electron has
an ms value of +1/2; the  electron has an ms value of ________. There is something called the ________
Exclusion Principle which says that each electron has a ____________ set of the 4 quantum numbers. The
maximum number of orbitals for a given energy level = __________. Since two electrons can fit into each
orbital, the maximum number of electrons that can fit in a given energy level is _______.
Sample problems:
Which set of quantum numbers are not possible? (The text asks which sets of quantum numbers are not
possible for a hydrogen atom - it’s the same question)
a] n = 3; l = 2; ml = 0; ms = +1/2
b] n = 2; l = 2; ml = -1; ms = -1/2 c] n = 4; l = 1; ml = 2; ms = +1/2
d] n = 1; l = 5; ml = 7; ms = +1/2 e] n = 5; l = 0; ml = 0; ms = -1/2
f] n = 5; l = 2; ml = -2; ms = -1
Write the quantum numbers for all of the electrons in 5B
What are the maximum number of orbitals that can have the designation:
1] 3p
2] 4dz2
3] n = 3
4] 3f
5] 5f
6] 4fxyz
7] 2s
What are the maximum number of electrons that can have the designation:
1] 2p
2] 3py
3] n = 2
4] n = 4
5] 2d
6] 3dxy
7] 5s
How many electrons can have the quantum numbers:
1] n = 3; ml = +1
2] n = 2; ms = +1/2
3] n = 2
5] n = 4; l = 3
6] n = 3; l = 2; ml = -2; ms = -1/2
4] n = 2; l = 3
7] n = 3; ms = -1/2
Write the values for the four quantum numbers for the highest energy, ground state electron for
1] 22Ti
2] 11Na
3] 52Te
4] 60Nd
5] 87Fr
6] 18Ar
What neutral ground state element has its highest energy electron described by the following four quantum
numbers?
1] n = 3; l = 2; ml = -1; ms = -1/2
2] n = 4; l = 1; ml = +1; ms = +1/2
3] n = 5; l = 3; ml = 0; ms = +1/2 4] n = 2; l = 0; ml = 0; ms = -1/2
5] n = 1; l = 1; ml = 0; ms = -1/2
On a test, I only want the expected electron configuration. However, there are elements that do not fit the
expected pattern. They are found in the d and f sublevels. For instance, 24Cr actual electron configuration is:
[Ar]4s13d5. The actual electron configuration for 29Cu is [Ar]4s13d10. Why does this happen? Because the
difference in ______________ between the 4s and 3d is small, the expected 4s2 electron can easily move
into the 3d. In the case of chromium, this causes the electron from 4s to move into the 3d so the sublevel is
half-filled. Half-filled sublevels are __________, and matter wants to be as stable as possible. In the case of
copper, moving the 4s electron into the 3d’s ________________the 3d’s which is stable.
Positive ions for the transition metals:
More definitions: valence electrons: electrons in the _________________ principal quantum level of an
atom.
Core electrons: inner electrons, or all electrons that are not _________electrons.
Alkali metals:
alkali earth metals:
halogens:
noble gases:
transition metals:
lanthanides:
actinides:
metalloids:
periodicity:
periods:
groups:
Important name: Dmitri __________________: put elements in sequence on the ___________ table by
chemical and physical properties.
Periodic trends in atomic properties:
1] ionization energy: the amount of energy needed to
Trend: generally increases going across the periodic table from left to right
always decreases going down the periodic table from top to bottom
explanations: period increase due to the fact that the electrons in the same period all come in at the same
__________. The number of protons _____________. Thus the electrons are held tighter as you move
across a period. When they become noble gases, they complete the __________, and become very
__________, thus causing the ionization energy to be very large. There are places where the trend reverses
within a period. Going from 3Be to 4B, or from 7N to 8O, we find such exception. You should be able to
explain why.
The ionization energy always decreases going down the periodic table due to the ____________________.
The _________________________ is the _____________________ outer electrons experience due to
_____ electrons. The more core electrons, the more it repels the outer valence electrons, making it easier to
__________ them.
Successive ionization energies refers to the energy to remove the next most outer electron. They are
indicated as I1, I2, I3, and so on for the first ionization energy, second ionization energy, third ionization
energy, and so on. In all cases, successive ionization energies always _______________. You should be
able to explain why. They do not go up evenly. At certain points, there is a giant increase in ionization
energies.
Sample problem: On a galaxy far, far away, the very reactive element tedonium, Td, is a period 3 element
with the following successive ionization energies: I1 = 45 zorks; I2 = 68 zorks; I3 = 622 zorks; I4 = 689
zorks; I5 = 765 zorks. Find the atomic number of Td, and predict the formula for the oxide of Td.
New problem: On a galaxy far, far away, the less reactive element woodsium, Wd, is a period 4 element
with the following successive ionization energies: I1 = 125 mims; I2 = 1004 mims; I3 = 1114 mims;
I4 = 1201 mims; I5 = 1308 mims. Find the atomic number of Td, and predict the formula for the nitride of
Wd.
Periodic trend #2: electron affinity: the _____________________ associated with the addition of an _____
in the gaseous atom. The more negative the value, the greater the amount of _____________ released when
the atom takes in an electron. The _____________ are elements that have high electron affinities.
Within a period, the elements have a stronger electron affinity as you move from left to right. That means
the energy values become more negative. This is because elements on the right side of the periodic table are
_____________, and nonmetals want to _________ electrons. Noble gases do not have an electron affinity.
Within a group, the electron affinities ____________ as you move down the periodic table. That is because
it is easier to add an electron when there are fewer ___________ electrons because core electrons repel
outer electrons (this is called the _________________________).
Periodic trend 3: atomic radius (how big the atom/ion is). As you move down the periodic table within a
______ or family, the atomic radius always __________________. This is because you start a new ______
with each step down. As you move across within a period from left to right, the atomic radius generally
________________. This is because of the increased ________________________. As you go across,
there are more ___________ which can attract the electrons that are coming in at the same level, so the
atom ____________. For ions, the more electrons that an atom has in becoming an _________________,
the larger the atom. The more electrons the atom ___________ in becoming a __________, the more that
atom/ion _______________. You can also compare different atoms with different charges if the number of
protons stays the same. ______________________ refers to different atoms and/or ions that have the same
electron configuration. If two different atoms and or ions have the same number of electrons, the one with
the _____________ number of protons will be smaller than the one with fewer protons in its nucleus.
Sample problems:
Place in order of predicted increasing atomic radius (no comparison is a possible answer)
1] P, Al, S
2] S, O, Te
3] P, Si O
4] S, N, P
5] Br, Br-1, Br+1
6] S, S-1, S-2
7] Ar, S-2, Ca+2, Cl-1, K+1
8] Na+1, Rb+1, K+1
Place in order of predicted increasing ionization energy (no comparison is a possible answer)
1] Si, Ge, C
2] Al, S, P
3] Al, C, Si
4] Kr, Cl, F
Types of Bonds
chemical bonding the result of sharing unpaired _______________ from different atoms. There are many
ways to classify bonds. We will first designate bonds as either covalent or ionic. Covalent means ______;
while ionic means _________________. But how do we determine this? It’s with a concept called
___________________, the ability of an atom to pull in _______________ electrons. The scale was
developed by Linus _________________, with the highest electronegativity going to ______________ at
4.0. The electronegativity of noble gases is essentially zero because they don’t want to gain electrons. [see
handout] The number ______ was developed as the cutoff point for ionic versus covalent bonding. If the
electronegativity difference is _____ or greater, the bond is ionic. Less than _____, it’s covalent. But within
covalent bonding, some electrons are more shared (or transferred) than others. So we have slightly
transferred covalent bonding electrons as a sub-category of covalent bonding. If the electronegativity
difference is between ____________, the bond is classified as ___________________. Polar refers to any
difference in electronegativity, and thus in charge. The negative end of a polar bond is the end of a bond
with the ____________ electronegativity. The positive end of a polar bond is the end with the ________
electronegativity. The bigger the difference in electronegativity, the more _________ the bond.
True or false: 1]All ionic bonds are polar.
3] Pure covalent bonds do not exist
2]All polar bonds are ionic.
4] Pure ionic bonds do not exist
Note the trend of electronegativity: Within a period, electronegativity increases going across (have to throw
out the _______________, however. Within a group, the electronegativity decreases going ________ the
periodic table. You should be able to explain why.
More vocabulary: A molecule that has a positive and negative centers of charge are called __________.
The __________________ is the property of a molecule whose charge distribution can be represented by a
center of positive charge and negative charge. We use the Greek letter, lower case, delta to indicate this
partial charge. It looks like: .
Sample problems.
For the following bonds, classify them as either polar covalent, pure covalent, or ionic. List the bonds from
least to most polar. Indicate the direction of the polarity for those bonds that are polar.
H - Cl
C-H
O-O
F - Al
K-F
There is another way to classify bonds: If the two atoms share ____ electrons, the bond is a single bond. If
the two atoms share _____ electrons, the bond is a _ _ _______ bond. If the two atoms share _____
electrons, the bond is a triple bond. There cannot be bonds made up of more than 6 electrons. To determine
if a bond is single, double or triple, we can either show the bonding using orbital notation, or write the __
____ formula, also known as the __ __________________________ .
Show the bonding for the following molecules, and write their structural formulas:
1.
NaCl
2.
H 2O
3.
O2
4.
N2
5.
NH3
6.
H 2O 2
7.
OF2
8.
MgCl2
Note that #8 has a problem:
We know Mg forms bonds. Therefore, we must have left off something. We did. It’s called _ ___________.
We know that the ________________________ column elements hybridize. They do so by doing the
following:
Now show the bonding using orbital notation for:
9.
KCN
10. HAlO
11. H2CO
12. C2H2
To write a Lewis formula, you determine the total number of ________________ electrons in the molecule.
You draw the skeletal molecule. Then put in lone pairs around atoms so that they will obey the ________
rule. Note: 1A’s end up with ____ electrons; 2A’s end up with _____ electrons; 3A’s end up with _____
electrons in the final Lewis formula for most stable, neutral molecules.
Write the Lewis formula for:
1] KBr
2] H2O
3] H2CO
4] NH4+1
5] Al F3
6] C2H4
7] NO-1
8] NaCN
In coordinate __________ bonds, the bonding electrons come from a __________ atom. Note the example
of: ammonia, _____ and ______.
Sometimes, more than one Lewis formula can be written for the same molecule. This is called __________.
An example:
The way to determine if a Lewis formula is correct is to determine the _________ charge on the atoms
within the molecule. The formal _________ is determined by the formula:
The sum of the formal charges must ________ the overall charge on that species.
If ______________ Lewis structures exist for a species, those with formal charges closest to ______ and
with any negative formal charges on the most ______________________ atoms are considered best to
describe the bonding in the atom or ion.
Molecular Structure: The VSEPR model (stands for ___________ shell electron _______ repulsion)
The basic premise is that the molecular structure is determined by minimizing electron-pair ___________.
In this model, single, double and triple bonds are considered to be one bonding pair. The number of
bonding pairs and lone pairs on the central atom determines the overall molecular shape. The central atom
is generally the one with the lowest electronegativity, or the atom in the molecule that has only one atom.
The lone pairs are the “double dots” around the central atom.
Linear molecules are AX2 (A is the central atom – X are the atoms around the central atom – E are the lone
pairs). The central atom in AX2 is sp hybridized.
Trigonal planar atoms are AX3. AX2E are bent. The central atom in AX3 and AX2E is sp2 hybridized.
Tetrahedral molecules are AX4. Trigonal pyramidal molecules are AX3E. AX2E2 is bent. The central atom
in these three molecules is sp3hyrbized
Trigonal bipyramidal molecules are AX5. The lone pairs go around the equatorial position. AX4E is seesaw.
AX3E2 is T-shaped. AX2E3 is linear. The central atom’s hybridization in these 4 molecules is dsp3
Octahedral molecules are AX6. Square pyramidal are AX5E. Square planar molecules are AX4E2. The
central atom’s hybridization in these molecules is d2sp3
See the handout.
Predict the shape of: HCN
HBO
H2CO
MgO
AsF5
TeCl6
CO2
NH4+
HNO
BF4-1
See handout. This method is based on the number of bonding atoms and the number of electron pairs. This
method can explain such interesting shapes as T-shaped (AX3E2), Seesaw (AX4E), square planar (AX4E2),
and square pyramid (AX5E). The pyramidal shape (trigonal pyramidal) is AX3E, and the bent shape is
AX2E2 or AX2E. You can be a linear shape if your formula is AX2E3.
Molecular Polarity
___________ can be classified as polar or nonpolar based on their electronegativity. Molecules can also be
classified as __________ or __________. This is important because polar molecules dissolve in ________,
or other polar solvents, but nonpolar molecules dissolve in _________ or ________, or other nonpolar
solvents. Nonpolar molecules have either 1] _____________ bonds, or 2] symmetry. Molecules with
nonpolar bonds include: ________________, molecules with the same two atoms, and _____________,
molecules that contain only hydrogen and _____________. Symmetric molecules have a ___________
atom that has hybridized, and all the _________________ are the same. Examples include:
Polar molecules have both 1] _____________ bonds, and 2] a lack of symmetry. All _________ and
________ shapes lack symmetry. If a molecule contains just ____ atoms that are different, the molecule
is ____________.
Determine the molecular polarity of
F2
C3H8
CH2F2
KrF6
CO2
AlF3
OF2
NH3
MgO
CF4
SbF5
CO2 is an example of a molecule with ___________ bonds, but is a __________ molecule.
A single bond consists of one _________ bond. A ______ bond is symmetric. It has the shape of a:
A double bond consists of one ___________ and one ______ bond. A ______ bond has the shape of a
_____ orbital. Pi bonds are ____________ to the plane of the bonding atoms. Double bonds are ________,
and ____________ than single bonds. A triple bond contains one _______ and two ________ bonds. The
two ______ bonds are perpendicular to each other. Thus, all pi bonds are __________. Triple bonds are
even ___________ and shorter than ____________ bonds.
Molecular orbital theory (MO theory)
All of the above discussion treated the electron as a ________________, but it is also a ________. The
concept of MO theory treats bonds as the interaction between electron ________.
There are two important equations regarding MO theory. Bond order:
Type of magnetism:
For n = 1
For n = 2
TTP. Predict the bond order and type of magnetism for:
1] H2 2] He2
3] N2
6] CN-1
4] O2
5]CO
7] NO+1
Isomer notes
If you bond C2H6O, you’ll find that there are two different ways to form this molecule. Each molecule is
completely different. One is _______________________ and the other is _____________________.
An alcohol has a hydrocarbon chain and an _______
An ether has two hydrocarbon chains about an ______________ atom.
Recall CH4 is___________; C2H6 is _____________; C3H8 is ________________; C4H10 is ___________.
When these are substituents, you drop the –ane suffix and replace with –yl.
Different molecules with the same molecular formula are called ________________________.
There are 3 different types of isomers. They are_____________________________________.
Structural isomers have a different ___________________________________________. Examples
include the ethyl alcohol and dimethyl ether identified above. Other important examples include:
Geometric isomers differ in the orientation of the atoms about a _______________ point. They come in two
types: _________________
Examples include:
Note: C2H2Cl2 can be written 3 ways. Note which is what type of isomer.
Generally, you can only have geometric isomers about a double bond, and in the octahedral shape. Note that
the molecular polarity of cis and trans isomers.
The third type of isomer is the ___________ isomer. They differ in the direction they can bend a plane of
____________________ light. These were first discovered by ______________________. They are
divided into two groups: those where you have to turn the second polarizer clockwise – they are
called_______________. And those where you have to turn the second polarizer _______________- and
they are called ______________________.
These isomers exist around _______________________ where the 4 ligands are all ________________.
These are important biologically. ______________________ are compounds that make _______________,
which make enzymes, and these make possible the biological reactions that keep us going. There are 22 of
these in nature. All of them are _________________________. Why?
Murchison meteorite and its significance.
It is also possible to have optical isomers around octahedral shapes.