Chemistry Chemical Bonding Slide 2 The Development of Atomic
Models 1. Dalton solid, indivisible mass Slide 3 2. Thomson
plum-pudding model -Negatively charged e- (raisins) stuck in
positively charged proton dough -No neutrons Slide 4 3. Rutherford
electrons surrounding a dense nucleus Slide 5 4. Bohr model
elctrons arranged in symmetrical orbits around the nuclues
-planetary model -Electrons in a given path have a fixed energy
level Slide 6 5. Quantum mechanical model modern mathematical
description of the atom Sodium atom: Slide 7 Energy level region
around the nucleus where the electron is likely to be moving. An
electron can jump from one level to another by absorbing energy.
Slide 8 Quantum the amount of energy required to move an electron
from its present energy level to the next higher one quantum leap
Slide 9 Quantum mechanical model uses mathematical equations to
describe the location and energy of electrons in an atom Developed
by Erwin Schrodinger Electrons are not in definite paths Their
location is described in terms of probability of being in a certain
region Electron cloud (ceiling fan) Conventionally, the border is
drawn at 90% probability Slide 10 Atomic orbital region in space
that an electron is likely to be in Electrons can be described by a
series of 4 quantum numbers. Slide 11 1. Principle quantum number
(n) Describes the energy level Values of 1, 2, 3, 4, etc. Slide 12
2. Azimuthal quantum number ( l ) Describes the shape of atomic
orbitals Sublevels Values of 0 to n-1 0 = s, 1 = p, 2 = d, 3 = f
Slide 13 s = spherical, p = peanut shape, d&f = more complex
shapes d = daisy f = fancy So if n = 1, then l can be 0 (s) = 1
sublevel n = 4, then l can be 0 (s), 1 (p), 2(d), 3(f), = 4
sublevels Slide 14 s orbitals spherical s orbitals spherical Slide
15 p orbitals dumbbell-shaped Slide 16 d orbitals daisy-shaped
Slide 17 f orbitals fancy shapes Slide 18 3. Magnetic quantum
number (m l ) Orientation of the orbital in space Values of l to +
l So s has 1 orbital p has 3 d has 5 f has 7 Slide 19 4. Spin
quantum number (m s ) Values of + and - Each orbital can hold 2
electrons with opposite spins Since spinning charged objects create
a magnetic field, the electrons must spin opposite directions to
minimize repulsion Slide 20 Slide 21 Ex. How many orbitals are in
the following? A. 3pD. 4p B. 2sE. 3d C. 4fF. 3 rd energy level How
many electrons can be in each of the above? Slide 22 Are these
possible? nlmlml msms 100+1/2 52-5-1/2 21+1/2 Slide 23 Electron
Configuration ways in which electrons are arranged around nuclei of
atoms. Slide 24 Rules that govern filling of atomic orbitals: 1.
Aufbau principle electrons enter orbitals of lowest energy first.
Slide 25 Slide 26 2. Pauli exclusion principle An atomic orbital
can describe at most two electrons. They must have opposite spins.
Slide 27 3. Hunds rule When electrons occupy orbital of equal
energy, one electron enters each orbital until all orbitals contain
one electron with parallel spins. Slide 28 Slide 29 1s 2s 3s 4s 5s
6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 4f Slide 30 Na = 11 1s 2 2s 2 2p 6 3s
1 Cd =48 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 Slide
31 Practice: Write the electron configuration for the following
elements: Li O Sc Slide 32 More practice: Identify each of the
following atoms on the basis of its electron configurations. a) 1s
2 2s 2 2p 6 neon b) 1s 2 2s 2 2p 6 3s 1 sodium c) [Kr] 5s 2 4d 2
zirconium d) [Xe] 6s 2 4f 6 samarium Slide 33 Ground state lowest
energy level for an electron. (Normal, nonexcited state) Slide 34
Exceptional Electron Configurations: Cr:expected: 1s 2 2s 2 2p 6 3s
2 3p 6 4s 2 3d 4 actual: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5
Cu:expected: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 actual: 1s 2 2s 2
2p 6 3s 2 3p 6 4s 1 3d 10 Half-filled energy levels are more stable
than other partially filled energy levels. There are other
exceptions. Slide 35 Light and Atomic Spectra Electromagnetic
radiation a series of energy waves that includes radio waves,
microwaves, visible light, infrared, and ultraviolet light, X-rays,
and gamma rays. Slide 36 Slide 37 Wavelength, crest trough Parts of
a wave: Slide 38 Amplitude height of the wave from the origin to
the crest. Wavelength - - distance between the crests Frequency f
or the number of wave cycles to pass a given point per unit of
time. The units of frequency are 1/s, s -1, or Hertz (Hz) Slide 39
c= f where c = speed of light = 3.00x10 8 m/s or 3.00x10 10 cm/s As
increases, f decreases. Slide 40 As wavelength increases, frequency
decreases. As wavelength decreases, frequency increases. Slide 41
Slide 42 Ex. A certain wavelength of yellow light has a frequency
of 2.73x10 16 s -1. Calculate its wavelength. Convert to nm. C = f
c f 3.00x10 8 m/s 2.73x10 16 s -1 x m Slide 43 Spectrum series of
colors produced when sunlight is separated by being passed through
a prism. ROY G. BIV Red: longest wavelength, lowest frequency
Violet: shortest wavelength, highest frequency Slide 44 Slide 45
Atomic emission spectrum series of lines of colored light produced
by passing light emitted by an excited atom through a prism. This
can be used to identify the element. The atomic emission spectrum
of hydrogen shows three series of lines. The lines in the UV region
(Lyman series) represent electrons falling to n=1, lines in the
visible region (Balmer series) represent electrons falling to n=2
and lines in the IR region (Paschen series) represent electrons
falling to n=3. Slide 46 Slide 47 Slide 48 Max Planck found that
the energy emitted or absorbed by a body changes only in small
discrete units he called quanta. He determined that the amount of
radiant energy, E, absorbed or emitted by a body is proportional to
the frequency of the radiation. E=h f E = energy (J) f = frequency
h = Plancks constant, 6.626x10 -34 Js Slide 49 Einstein studied the
photoelectric effect whereby light of sufficient frequency shining
on a metal causes current to flow. The amplitude of the radiation
was not important, the frequency was. This told him that light must
be in particles, each having a given energy. Einstein proposed that
electromagnetic radiation can be viewed as a stream of particles
called photons: E=h f Slide 50 Photoelectric Effect Electron
(photons) light metal Slide 51 Energy of a photon: E = h f Energy
of a photon: E = h f Slide 52 Example: Calculate the energy of an
individual photon of yellow light having a frequency of 2.73x10 16
s -1. E=h f E = (6.626x10 -34 Js)(2.73x10 16 s -1 ) E = 1.81x10 -17
J Slide 53 Einsteins special theory of relativity: E=mc 2 Matter
and energy are different forms of the same entity. Slide 54 Going
Further: Slide 55 Louis deBroglie suggested that very small
particles like electrons might also display wave particles and he
came up with: deBroglies equation: = h mv m = mass in kg v =
velocity in m/s h = Plancks constant, 6.626x10 -34 Js Slide 56
DeBroglies equation is used to find the wavelength of a particle.
It was determined that matter behaves as through it were moving in
a wave. This is important in small object such as electrons but is
negligible in larger objects such as baseballs. Heavy objects have
very short wavelengths. Slide 57 Example: Calculate the wavelength
of an electron traveling at 1.24x10 7 m/s. The mass of an electron
is 9.11x10 -28 g. = h mv 9.11x10 -28 g 1 kg = 9.11x10 -31 kg 1000g
= (6.626x10 -34 Js) = (9.11x10 -31 kg)(1.24x10 7 m/s) = 5.87x10 -11
m Slide 58 End Going Further. Slide 59 In the photoelectric effect,
electrons (called photoelectrons) are ejected by metals (esp.
alkali) when light of sufficient frequency shines on them. Red
light wont work. Photoelectric cells convert light energy into
electrical energy. They are used in automatically opening doors and
security systems. Slide 60 Heisenbergs Uncertainty Principle it is
impossible to determine accurately both the momentum and the
position of an electron simultaneously. We detect motion by
electromagnetic radiation. This interaction disturbs electrons.
Slide 61 Einstein and Heisenberg! Slide 62 Mendeleev- arranged
elements in order of increasing atomic mass Moseley- arranged
elements in order of increasing atomic number Slide 63 Periodic Law
- When the elements are arranged in order of increasing atomic
number, there is a periodic pattern in their physical and chemical
properties. Slide 64 Be able to locate noble gases, representative
elements, transition metals, inner transition metals. Slide 65
Noble Gases have completely filled shells of electrons similar
electronic structures He1s 2 Ne1s 2 2s 2 2p 6 Ar1s 2 2s 2 2p 6 3s 2
3p 6 Kr1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4p 6 etc. Slide 66
Representative Elements elements in A groups on periodic chart
representative - because they best represent what we know about
elemental structure & periodicity Slide 67 d - Transition
Elements elements in B groups on periodic chart metals have d
electrons transition from metals to nonmetals Slide 68 f -
Transition Elements inner transition metals Slide 69 Electron
configuration using the periodic table: Slide 70 Slide 71 Columns
are called groups or families. Rows are called periods or series.
Slide 72 Shorthand electron configuration: Slide 73 Slide 74 Slide
75 Periodic Trends in atomic size: covalent atomic radius - half
the distance between the nuclei of atoms in a homonuclear diatomic
molecule (like Cl 2 ) Trends: Slide 76 Slide 77 Slide 78 Atomic
size increases going down a group because electrons are added to
higher energy levels that are farther from the nucleus. It
decreases going across a period because as each proton is added to
the nucleus an electron is being added to the same energy level.
This shell of electrons is pulled closer in towards the nucleus.
Slide 79 Size changes little in the transition metals because the
electrons being added are core electrons. Slide 80 Z eff =
effective nuclear charge -actual pull of the nucleus on the valence
electrons. Z eff = Z actual - effect of e - repulsions Slide 81
Trends: Increases from H to He Decreases from He to Li because 1s
electrons shield 2s electrons Increases from Li to Be Decreases
from Be to B because 2s shield 2p Slide 82 Increases from B to C to
N Decreases from N to O because of repulsion due to doubly occupied
orbitals Increases from O to F to Ne Decreases from Ne to Na
because 1s,2s & 2p shield 3s Slide 83 Know exceptions to Z eff
trends and reasons for these. Slide 84 Ionization energy (IE) -
energy required to remove the highest energy electron from a
gaseous atom Li(g) + energy Li + (g) + e - - depends on Z eff and
size Slide 85 Slide 86 Trends: Ionization energy decreases down a a
group because the valence electrons are farther from the nucleus
and are thus held less tightly. Ionization energy increases across
a period because atomic size decreases and valence electrons are
held more tightly. Z eff increases. Slide 87 Slide 88 1 st
ionization energy = energy required to remove the first electron Al
+ energy Al + + e - 2 nd ionization energy = energy required to
remove the second electron Al + + energy Al 2+ + e - 3 rd
ionization energy = energy required to remove the third electron Al
2+ + energy Al 3+ + e - Slide 89 For Na, the 1 st ionization energy
is fairly low but the 2 nd would be high. Slide 90 Know exceptions
to ionization energy trends and reasons for them. Slide 91 Ionic
Size Anions are larger than the atoms from which they were formed.
Know why!!! Cations are smaller than the atoms from which they were
formed. Know why!!!! Slide 92 Sizes of Ions Related to Positions of
the Elements in the Periodic Table Slide 93 Isoelectronic ions- a
group of ions with the same number of electrons The one with the
highest atomic number is the smallest in size (More protons pulling
on the same # of electrons). Slide 94 Na +, Mg 2+, Ne, F -, O 2-, N
3- are isoelectronic. They all have 10 electrons. Mg 2+ is the
smallest because it has 12 protons pulling on 10 electrons. (The
protons win the tug of war) N 3- is the largest because it has 7
protons and 10 electrons. (The electrons win the tug of war) Slide
95 Electron Affinity -energy change that occurs when an electron is
added to a gaseous atom -usually exothermic Cl(g) + e - Cl - (g)+
energy Trends: (but many exceptions) Slide 96 Slide 97
Electronegativity -relative tendency of an atom to attract shared
electrons to itself Trends: FONCl (Phone Call) Slide 98 Slide 99
Elements with high electronegativity (nonmetals ) tend to gain
electrons to form anions. Elements with low electronegativities
(metals) often lose electrons to form cations. Slide 100 Valence
electrons- electrons in the highest occupied energy level of an
atom. Valence electrons are the only electrons involved in the
formation of chemical bonds. Slide 101 Electron dot structures for
atoms: -each dot represents a valence electron p p X s p Electron
dot structures for atoms: -each dot represents a valence electron p
p X s p Slide 102 Examples: N Slide 103 N Slide 104 O O Slide 105 O
Slide 106 Xe Slide 107 Slide 108 Al Al Slide 109 Slide 110 Na Na
Slide 111 Slide 112 I Slide 113 I Slide 114 Si Si Slide 115 Slide
116 One of the major driving forces in nature is the tendency to go
to lower energy. Atoms lose, gain or share electrons to become
lower in energy and thus more stable. Slide 117 Metals lose
electrons easily to become positively charged cations. They will
usually lose their valence electrons to achieve a noble gas
electron configuration. Na Na + + e - [Ne]3s 1 [Ne] Al Al 3+ + 3e -
[Ne]3s 2 3p 1 [Ne] Slide 118 Some transition metals lose their
highest energy level s and p electrons but still have d electrons
remaining. Their electron configuration is not quite that of a
noble gas but is still stable. It is called a pseudo-noble gas
electron configuration. For example, zinc loses its two electrons
in 4s but keeps the ten electrons in 3d. Slide 119 Transition
metals always lose their highest numerical energy level electrons
first. Transition metals in the 4th period lose their 4s and 4p
electrons before losing any from 3d. Metals in groups 3, 4, & 5
do this also. Transition metals always lose their highest numerical
energy level electrons first. Transition metals in the 4 th period
lose their 4s and 4p electrons before losing any from 3d. Metals in
groups 3, 4, & 5 do this also. Slide 120 Example: Fe [Ar]4s 2
3d 6 Fe 2+ [Ar]3d 6 Fe 3+ [Ar]3d 5 Slide 121 Nonmetals tend to gain
electrons to become stable and form negatively charged anions. They
achieve a noble gas electron configuration. Example: Cl + e - [ Cl
] - [Ne]3s 2 3p 5 [Ne]3s 2 3p 6 N + 3e - [ N ] 3- 1s 2 2s 2 2p 3 1s
2 2s 2 2p 6 Slide 122 Ionic Bonding- the attraction of oppositely
charged ions (cations and anions) When the electronegativity
difference between two elements is large, the elements are likely
to form a compound by ionic bonding (transfer of electrons). The
farther apart across the periodic table two Group A elements are,
the more ionic their bonding will be. Slide 123 We can use Lewis
dot formulas to represent the formation of ionic compounds. Na + Cl
Na + [ Cl ] - or NaCl Mg: N Mg 2+ [ N ] 3- Mg: Mg 2+ or Mg 3 N 2
Mg: N Mg 2+ [ N ] 3- Slide 124 Properties of Ionic Compounds: They
are usually crystalline solids with high melting points (>400 o
C) Their molten compounds and aqueous solutions conduct electricity
well because they contain mobile charged particles. Slide 125
Metals Metals form metallic solids that consist of positively
charged metal cations in a sea of loosely held valence electrons.
This arrangement allows metals to have their unique properties.
Slide 126 Slide 127 Metals are ductile (can be pulled into a wire)
and malleable (can be hammered into a thin sheet) because the
valence electrons act as grease, allowing the cations to slide past
each other without colliding with each other and shattering. When
ionic compounds such as NaCl are hammered, like-charged ions
collide causing repulsion and the crystal shatters. Slide 128
Metals can conduct electricity easily. Electricity is a flow of
electrons. As electricity (electrons) enters one end of a piece of
metal, an equal number of electrons exit the other end. Slide 129
Alloys- solutions of solids in solids Slide 130 Substitutional
alloy- atoms of one metal are substituted for atoms of a
similar-sized metal in a metallic crystal. Ex. brass, sterling
silver, pewter Interstitial alloy- smaller metal atoms fit into
holes in the crystal structure of a metal with larger atoms Ex.
steel (carbon in iron) Slide 131 Slide 132 Amalgam- alloy which
contains mercury Slide 133 Covalent Bonding Slide 134 Hydrogen and
nonmetals of Groups 4,5,6 & 7 often become stable and gain
noble gas electron configurations by sharing electrons to form
covalent bonds. Atoms will usually share electrons to follow the
octet rule (eight electrons, like most noble gases) or the duet
rule (2 electrons, like helium). Slide 135 When atoms share one
pair of electrons to form a covalent bond, it is called a single
covalent bond. The electrons shared between the atoms are a shared
pair. A dash can be used instead of two dots to represent the
shared pair. Any other electrons on the atoms are unshared pairs or
lone pairs. Slide 136 Ex. H 2 H-H Cl 2 Cl-Cl HCl H-Cl Ex. H 2 H-H
Cl 2 Cl-Cl HCl H-Cl Slide 137 Atoms must sometimes share more than
one pair of electrons to become stable. When two pair of electrons
are shared between two atoms, it is called a double bond. If three
pair are shared, it is a triple bond. Ex. O 2 N 2 O = O N N Slide
138 Rules for Writing Lewis Structures (electron dot structures):
(Use pencil!) 1.Add up the valence electrons from all the atoms.
Dont worry about keeping track of which electrons come from which
atoms. If you are working with an ion, you must add or subtract
electrons to equal the charge. 2.Use a pair of electrons to form a
bond between each pair of bound atoms. Slide 139 3.Arrange the
remaining electrons to satisfy the duet rule for hydrogen and the
octet rule for everything else. 4.If necessary, change bonds to
double or triple. 5.Remember, we cannot create or destroy
electrons! Slide 140 H 2 O H 2 O 8 electrons H-O-H Slide 141 NH 3
NH 3 8 electrons H-N-H H Slide 142 NH 4 + 9-1 = 8 electrons H +
H-N-H H Slide 143 CO 2 CO 2 16 electrons O-C-O This used 20
electrons! BAD!!! Slide 144 CO 2 O=C=O Slide 145 CCl 4 CCl 4 32
electrons Slide 146 CCl 4 CCl 4 32 electrons Cl Cl-C-Cl Cl Slide
147 CN - 9 + 1 = 10 electrons C-N BAD! BAD! Slide 148 CN - 9 + 1 =
10 electrons C N Slide 149 SO 4 2- SO 4 2- 32 electrons O 2 - O S O
O Slide 150 CO 3 2- 24 electrons CO 3 2- 24 electrons O O C O O 2-
2- Slide 151 Coordinate Covalent Bond- Bond in which both electrons
came from the same atom. This bond is not really any different than
any other single bond. Slide 152 Exceptions to the Octet Rule: A
few compounds are stable with less than an octet. They include
beryllium or boron. These electron deficient compounds are very
reactive. Slide 153 Ex. BF 3 BeCl 2 Ex. BF 3 BeCl 2 24 electrons 16
electrons F Cl-Be-Cl F-B-F Slide 154 Elements in the third period
and below can exceed the octet rule. They can place extra electrons
in empty d orbitals. Elements in the second period can not exceed
the octet rule because there is no 2d orbital for the extra
electrons to go into. If it is necessary to exceed the octet rule,
place the extra electrons on the central atom. Slide 155 Ex. PCl 5
SF 6 Cl Cl F F F P S Cl Cl Cl F F F Slide 156 I3-I3-I3-I3- 22
electrons [ I-I-I ] - Slide 157 More practice: NF 3 Slide 158 F F -
N - F F F - N - F Slide 159 OF 2 20 electrons OF 2 20 electrons
Slide 160 F - O - F Slide 161 KrF 4 36 electrons Slide 162 F F- Kr
- F F F F- Kr - F F Slide 163 BeH 2 4 electrons BeH 2 4 electrons
Slide 164 H - Be - H H - Be - H Slide 165 SO 3 2- 26 electrons
Slide 166 O O - S - O O O - S - O Slide 167 NO 3 - 24 electrons NO
3 - 24 electrons Slide 168 O O - N - O O O - N - O Slide 169 H 2 O
2 14 electrons H 2 O 2 14 electrons Slide 170 H - O - O - H H - O -
O - H Slide 171 Resonance occurs when more than one valid Lewis
structure can be written for a molecule. The actual structure is an
average of all of the resonance structures. Slide 172 Ex. NO 3 - O
O O O - N - O O - N - O O - N - O Slide 173 In nitrate, the
experimental bond length is in-between that of a single bond and a
double bond. It acts like a 1 1/3 bond. Slide 174 Ex. Benzene, C 6
H 6 Slide 175 Slide 176 VSEPR Slide 177 Lewis structures can be
used to determine the shapes of molecules. Their shapes will tell
us a lot about their chemical behavior. Slide 178 The valence shell
electron pair repulsion (VSEPR) theory tells us that valence
electrons on the central atom repel each other. They are arranged
as far apart as possible around the central atom so that repulsions
among them are as small as possible. When we are using VSEPR to
determine molecular shape, we are really looking for regions of
electron density. Double and triple bonds count the same as single
bonds in determining molecular shape. Slide 179 In CO 2, there are
only two regions of electron density (effective electron pairs)
around the central atom. These regions arrange themselves as far
apart as possible, making the bond angle 180 o and the molecular
shape linear. O = C = O Slide 180 In CO 3 2-, there are three
effective electron pairs around the central atom. The bond angle
will be 120 o and the shape will be trigonal planar. 2 - O C O O In
CO 3 2-, there are three effective electron pairs around the
central atom. The bond angle will be 120 o and the shape will be
trigonal planar. 2 - O C O O Slide 181 In CH 4, there are four
effective electron pairs. We might expect the bond angle to be 90
o. Actually, since molecules are three-dimensional, the electron
pairs are 109.5 o apart (further than 90 o ) and take a tetrahedral
arrangement. H H C H H Slide 182 Slide 183 In NH 3, there are four
effective electron pairs. Three are shared but one is unshared.
Unshared pairs of electrons take up more space than shared pair
because they are pulled closer to the nucleus. The presence of the
unshared pair distorts the other bond angles, making them less than
109.5 o and the shape is called trigonal pyramidal. (The bond
angles in ammonia are about 107 o.) H - N - H H Slide 184 Slide 185
In H 2 O, there are four effective electron pairs, also. Two are
shared and two are unshared. Since the unshared pairs repel more
than shared pair, the bond angle is less than 109.5 o (actually
104.5 o for H 2 O) and the shape is bent. H O H Slide 186 Slide 187
Slide 188 In PF 5, there are five effective pairs, all shared. The
bond angles are 90 o and 120 o and the shape is called trigonal
bipyramidal. It is like two trigonal pyramids with their bases
touching. F F F - P - F F Slide 189 Slide 190 In SF 6, there are
six effective pairs, all shared. The bond angles are 90o and the
shape is called octahedral. It is like two square pyramids with
their bases touching. In SF 6, there are six effective pairs, all
shared. The bond angles are 90 o and the shape is called
octahedral. It is like two square pyramids with their bases
touching. Slide 191 Slide 192 Molecules that exceed the octet rule
and have unshared electrons can have more complex shapes such as
T-shaped, see-saw, and square pyramidal. Slide 193 Practice
Determining Molecular Shape: Slide 194 H 2 S H - S - H bent,
