quantum chemistry (real)
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Ch. 4 Arrangement of
Electrons in Atoms
4.1 The Development of a
New Atomic Model
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Light
Before 1900, scientists thought thatlight behaved only as wave
discovered that also has particle-likecharacteristics
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Light as a Wave
electromagnetic radiation: form of energy that acts as a wave as it
travels
includes: visible light, X rays, ultravioletand infrared light, microwaves, and radiowaves
All forms are combined to formelectromagnetic spectrum
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Light as a Wave
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Light as a Wave
all form of EM radiation travel at aspeed of 3.0 x 108 m/s in a vacuum
it has a repetitive motion
wavelength: () distance betweenpoints on adjacent waves; in nm(109nm = 1m)
frequency: () number of waves thatpasses a point in a second, inwaves/second
Inversely proportional!
=c
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Photoelectric Effect
when light is shone on a piece ofmetal, electrons can be emitted
no electrons were emitted if thelights frequency was below a certainvalue
scientists could not explain this withtheir classical theories of light
Ex: coin-operated sift drink machine
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Photoelectric Effect
Max Planck: a German physicist
suggested that an object emits
energy in the form of small packets ofenergy called quanta
quantum- the minimum amount ofenergy that can be gained or lost byan atom
Plancks constant (h): 6.626 x 10
-34
J*s
hE=
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Photoelectric Effect
Einstein added on to Plancks theoryin 1905
suggested that light can be viewed
as stream of particles photon- particle of EM radiation
having no mass and carrying one
quantum of energy energy of photon depends on
frequency
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Photoelectric Effect EM radiation can only be absorbed by
matter in whole numbers of photons
when metal is hit by light, anelectron must absorb a certain
minimum amount of energy to knockthe electron loose
this minimum energy is created by a
minimum frequency since electrons in different metal
atoms are bound more or less tightly,
then they require more or less
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H Line-Emission Spectrum
ground state- lowest energy state ofan atom
excited state- when an atom has
higher potential energy than it has atground state
line-emission spectrum- series of
wavelengths of light created whenvisible portion of light from excited
atoms is shined through a prism
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H Line-Emission Spectrum
scientists using classical theoryexpected atoms to be excited bywhatever energy they absorbed
continuous spectrum- emission ofcontinuous range of frequencies of EM
radiation
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H Line-Emission Spectrum
Why had hydrogen atoms only givenoff specific frequencies of light?
current Quantum Theory attempts to
explain this using a new theory of atom
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H Line-Emission Spectrum
when an excited atom falls back toground state, it emits photon ofradiation
the photon is equal to the differencein energy of the original and finalstates of atom
since only certain frequencies areemitted, the differences between thestates must be constant
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Bohr Model
created by Niels Bohr
(Danish physicist)
in 1913 linked atoms electron with emission
spectrum
electron can circle nucleus in certainpaths, in which it has a certainamount of energy
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Bohr Model
Can gain energy bymoving to a higherrung on ladder
Can lose energy bymoving to lower rungon ladder
Cannot gain or losewhile on same rung ofladder
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Bohr Model
a photon isreleased thathas an energy
equal to thedifferencebetween the
initial and finalenergy orbits
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Bohr Model
problems:
did not work for other atoms
did not explain chemical behaviorof atoms
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Introduction to Quantum Th
Quantum Theory-
describes mathematically the waveproperties of electrons
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Electrons as Waves In 1924, Louis de Broglie
(French scientist) suggested the way quantized
electrons orbit the nucleus is similar to
behavior of wave electrons can be seen as waves
confined to the space around a nucleus
waves could only be certainfrequencies since electrons can onlyhave certain amounts of energy
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Electrons as Waves
hvE=
vc =cv =
hcE=
2mcE=
2mchc =vm
h=
shows that anything with both mass andvelocity has a corresponding wavelength
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Uncertainty Principle
In 1927 by Werner Heisenberg(German theoretical physicist)
electrons can only be detected bytheir interaction with photons
any attempt to locate a specificelectron with a photon knocks theelectron off course
Heisenberg Uncertainty Principle- it isimpossible to know both the position
and velocity of an electron
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Schrdinger WaveEquation
In 1926, Erwin Schrdinger
(Austrian physicist)
his equation proved thatelectron energies are quantized
only waves of specific energies
provided solutions to his equation solutions to his equation are called
wave functions
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Schrdinger WaveEquation
wave functions give only theprobability of finding an electron in acertain location
orbital- 3D area around a nucleusthat has a high probability ofcontaining an electron
orbitals have different shapes andsizes
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Quantum Numbers
specify the properties of atomicorbitals and of electrons in orbitals
the first three numbers come fromthe Schrdinger equation anddescribe: main energy level
shape
orientation
4th describes state of electron
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1st Quantum Number
Principal Quantum Number: n main energy level occupied by
electron
values are all positive integers(1,2,3,)
As n increases, the electrons energyand its average distance from the
nucleus increase multiple electrons are in each level
so have the same n value
the total number of orbitals in a level
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1st Quantum Number
Energy
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2nd Quantum Number
Angular Momentum QuantumNumber: l
indicates the shape of the orbital
(sublevel) for a certain energy level, the number
of possible shapes is equal to n
the possible values ofl are 0 and allpositive integers less than or equal to n-1
each atomic orbital is designated by therinci al uantum number followed b
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2nd Quantum Number
s orbitals:
spherical l value of 0
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2nd Quantum Number
p orbitals:
dumbbell-shaped
l value of 1
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2nd Quantum Number
d orbitals:
various shapes
l value of 2
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2nd Quantum Number
f orbitals:
various shapes l value of 3
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2nd Quantum Number
Level
Sublevels
Sublevels
0 1 23
0 1 2
0 1
0
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3rd Quantum Number
Magnetic Quantum Number: ml
indicates the orientation of an orbital
around the nucleus has values from +l -l
specifies the exact orbital that the
electron is contained in each orbital holds maximum of 2
electrons
EnergyEnergy SublevelSublevel ## Total #Total #
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EnergyEnergy
LevelLevel
(n)(n)
SublevelSublevel
s in Levels in Level##
OrbitalsOrbitals
inin
SublevelSublevel
Total #Total #
ofof
OrbitalsOrbitals
in Levelin Level11 ss 11 11
22 ss 11 44
pp 33
33 ss 11 99
pp 33
dd 55
44 ss 11 1616
pp 33
dd 55
ff 77
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4th Quantum Number
Spin Quantum Number: ms
indicates the spin state of the
electron only 2 possible directions
only 2 possible values: + and -
paired electrons musthave opposite spins
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Energy Level 1
nn ll mmll
mmss
11 00 00 -,+-,+
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Energy Level 2
nn ll mmll
mmss
22 00 00 -,+-,+
11 -1-1 -,+-,+
00 -,+-,+
+1+1 -,+-,+
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Energy Level 3
nn ll mmll
mmss
33 00 00 -,+-,+
11 -1-1 -,+-,+
00 -,+-,+
+1+1 -,+-,+
22 -2-2 -,+-,+
-1-1 -,+-,+
00 -,+-,+
+1+1 -,+-,+
+2+2 -,+-,+
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Energy Level 4nn ll mm
llmm
ss
44 00 00 -,+-,+
11 -1-1 -,+-,+
00 -,+-,+
+1+1 -,+-,+
22 -2-2 -,+-,+
-1-1 -,+-,+
00 -,+-,+
+1+1 -,+-,+
+2+2 -,+-,+
ll mmll
mmss
33 -3-3 -,+-,+
-2-2 -,+-,+
-1-1 -,+-,+
00 -,+-,+
+1+1 -,+-,++2+2 -,+-,+
+3+3 -,+-,+
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Electron Configurations
the arrangement of electrons in anatom
each type of atom has a uniqueelectron configuration
electrons tend to assume positions thatcreate the lowest possible energy for
atom
ground state electron configuration-lowest energy arrangement of
electrons
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Rules for Arrangements
Aufbau Principle- anelectron occupies thelowest-energy orbital
that can receive it
Beginning in the 3rdenergy level, theenergies of thesublevels in differentenergy levels begin to
overlap
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Rules for Arrangements
Pauli Exclusion Principle- no twoelectrons in the same atom can havethe same set of 4 quantum numbers
Hunds Rule- orbitals of equal energyare each occupied by one electronbefore any orbital is occupied by asecond
all unpaired electrons must have thesame spin
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Rules for Arrangements
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Writing Configurations
Orbital Notation: an orbital is written as a line
each orbital has a name written below it
electrons are drawn as arrows (up anddown)
Electron Configuration Notation number of electrons in sublevel is added
as a superscripthttp://www.cowtownproductions.com/vining/Sims/atomic_el
O d f illi S bl l
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Order for Filling Sublevels
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Writing Configurations
Start by finding the number of electrons inthe atom
Identify the sublevel that the last electronadded is in by looking at the location inperiodic table
Draw out lines for each orbital beginningwith 1s and ending with the sublevelidentified
Add arrows individually to the orbitals untilall electrons have been drawn
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Silicon
number of electrons: 14 last electron is in sublevel: 3p
1s 2s 2p 3s3p
Valence Electrons- the electrons in the
outermost energy level
Chl i
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Chlorine
number of electrons: 17 last electron is in sublevel: 3p
2p 3s 3p1s 2s
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Sodium
number of electrons: 11
last electron is in sublevel: 3s
1s2 2s2 2p63s1
1s 2s 2p 3s
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Calcium
number of electrons: 20
last electron is in sublevel: 4s
1s2 2s2 2p6 3s2 3p6 4s2
1s 2s 2p 3s
3p 4s
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Bromine
number of electrons: 35
last electron is in sublevel: 4p
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
1s 2s 2p 3s 3p
4s 3d 4p
1s 2s 2p 3s 3p
4s 3d 4p
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Argon
number of electrons: 18
last electron is in sublevel: 3p
1s2 2s2 2p63s2 3p6
1s 2s 2p 3s 3p
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Noble Gas Notation
short hand for larger atoms
configuration for the last noble gas isabbreviated by the noble gass symbol in
brackets
l C fi i i
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Electron Configuration Exceptions
Copper
EXPECT: [Ar] 4s2
3d9
ACTUALLY: [Ar] 4s1 3d10
Copper gains stability with a fulld-sublevel.
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Electron Configuration Exceptions
Chromium
EXPECT: [Ar] 4s2
3d4
ACTUALLY: [Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel.
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1
2
3
4
5
6
7
Full sublevel (s, p, d, f)
Half-full sublevel
Stability
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Bell Ringer
What do you already know abouthow bonds are formed? Are theredifferent types?
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Bonding
Introduction to ChemicalBonding
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Chemical Bonds
atoms rarely exist alone
when atoms are bonded together,they have less potential energy and
are more stable
What is potential energy?
chemical bond mutual electricalattraction between the nuclei andvalence electrons of different atomsthat binds the atoms together
Ionic Bonds results from
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Ionic Bonds results fromelectricalattractionbetweenlargenumbers of
cations andanions
atomsdonate oracceptelectronsfrom each
other
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Covalent Bonds
results from sharing ofelectron pairsbetween two atoms
the electrons sharedbelong to both atoms
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Covalent Bonds
Polar Covalent
when electrons are
shared unevenly
Nonpolar Covalent
when electrons areshared evenly
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Ionic vs. Covalent
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Ionic vs. Covalent bonding usually does not fall in one
category or the other, but somewhere inbetween
type of bond depends on the elements
differences in electronegativities
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Ionic vs. Covalent
Difference inDifference inelectronegativitieelectronegativitie
ss
Percent IonicPercent IonicCharacterCharacter
IonicIonic > 1.7> 1.7 > 50 %> 50 %
PolarPolar
CovalentCovalent0.3 1.70.3 1.7 5 50 %5 50 %
NonpolarNonpolar
CovalentCovalent0 0.30 0.3 0 5 %0 5 %
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Polarity
Polar- unevendistribution of charge
Show partial charges
on structure by using (lowercase delta)
Practice
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Practice
Determine whether each of the followingbonds will be:
ionic, polar covalent, OR nonpolarcovalent
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Practice
S and H2.5-2.1=0.4
polar covalent
S and Cs2.5-0.7=1.8
ionic
C and Cl
3.0-2.5=0.5
polar covalent
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Practice
Cl and Ca3.0-1.0=2.0
ionic
Cl and O3.5-3.0=0.5
polar covalent
Cl and Br
3.0-2.8
nonpolar covalent
B ll Ri
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Bell Ringer How do you determine whether a
compound is molecular or ionic?Give an example of each.
Write the formula for thecompound made from:Mg and O
Ca and Br
Li and N
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Covalent Bonding
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Molecular Compounds
molecule: neutral group of atoms heldtogether by covalent bonds
molecular compound: compound whose
simplest unit is a molecule
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Formulas
chemical formula: tells the number ofeach type of atom in a compound
molecular formula: tells the numberof each type of atom in a molecularcompound
ex. H2O, Cl2, C6H12O2
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Molecular Compounds
diatomic molecule: a molecules containingonly 2 atoms
usually refers to 2 of the same atoms
ex: O2, Br2, F2, etc.
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Formation of Covalent Bond
Formation of Covalent Bond
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Formation of Covalent Bond
approaching nucleiand electron cloudsare attracted toeach other tocreate a decrease
in Potential Energy(PE)
two nuclei and two
electron cloudsrepel each othercreating anincrease in PE
i f C l d
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Formation of Covalent Bond
a distance between the nuclei isreached in which repulsion and attraction forces are equal
potential energy is at the lowest pointpossible
at the bottom of the curve on PE graph
C l B d
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Covalent Bonds
Bond Length distance between two bonded atoms at
their lowest PE
average distance since there are somevibrations
measured in pm (1012 pm = 1 m)
stronger the bond, shorter the bond
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C l t B d
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Covalent Bonds
Bond Energy energy is released when atoms become
because they have lower PE
the same amount of energy must beused to break the bond and form neutralisolated atoms
stronger bond, higher bond energy
average since varies a small amountbased on atoms in entire molecule
in kJ/mol
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10/28 St t
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10/28 Starter
Which elements naturally exist asdiatomic molecules? Remember, the 7 + 1 rule
How many valence electrons do eachof the halogens have?
Show or describe how two bromineatoms would form a covalent bond.
O t t R l
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Octet Rule
representative elements can filltheir outer energy level by sharingelectrons in covalent bonds
Octet Rule- a compound tends toform so that each atom has an octet
(8) of electrons in its highest energylevel by gaining, losing or sharingelectrons
Duet Rule- applies to H and He
O t t R l
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Octet Rule Less than 8:
Boron: 6 in outer energy level
More than 8:
anything in 3
rd
period or heavier because may use the empty d orbital
ex: S, P, I
El t D t Di
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Electron Dot Diagrams
a way to show electronconfiguration
identifies the number and pairing of
valence electrons to show howbonding will occur
3. write the noble gas notation
4. identify the number of valence5. identify how many are paired and
how many are alone
6. do not go by Figure 6-10
E l
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Example
Nitrogen 1s2 2s2 2p3
5 valence
2 are paired 3 are alone
Sulfur 1s2 2s2 2p6 3s2 3p4
6 valence
4 paired (2 pairs) 2 are alone
N
L i St t
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Lewis Structures
like dot diagrams but for entiremolecules
atomic symbols represent nucleus
and core electrons and dots ordashes represent valence electrons unshared electrons: (lone pairs) pair of
electrons not involved in bondingwritten around only one symbol
bonding electrons: written in between 2atoms as a dash
T f B d
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Types of Bonds
single- sharing of one pair ofelectrons weakest, longest
double- sharing of 2 pairs ofelectrons stronger and shorter
triple- sharing of 3 pairs of electrons strongest and shortest
multiple bonds include double and
tri le bonds
Drawing Lewis Structures
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Drawing Lewis Structures
find the number of valenceelectrons in each atom and addthem up
draw the atoms next to each otherin the way they will bond
add one bonding pair between each
connected atoms add the rest of the electrons until
all have 8 (consider exceptions to octet
rule)
Example 1 CH
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H
H C Cl
H
Example 1 CH3Cl methyl chloride
C: 4 x 1 = 4
H: 1 x 3 = 3
Cl: 7 x 1 = 7
total = 14 electrons
carbon is central H
H C Cl
H
duet
duet
duet
octet
octet
Example 2 NH3
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Example 2NH3 ammonia
N: 5 x 1 = 5
H: 1 x 3 = 3 total = 8
N is central
H N H
H
Example 3
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Example 3
N2 nitrogen gas
N: 5 x 2 = 10
10 electrons
N N
N N
Example 4
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Example 4
CH2O formaldehyde
C: 4 x 1 = 4
H: 1 x 2 = 2
O: 1 x 6 = 6
total = 12
C is central
H C H
O
Polyatomic Ions
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Polyatomic Ions
charged group of covalently bondedatoms
Example: CN-
NH + : ammonium ion
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NH4+ : ammonium ion
SO42- : sulfate ion
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4
5 x 6 = 30
total = 30 + 2 = 32
OH- : hydroxide ion
6 + 1 + 1 = 8 total
SO
O
OO
O H
Example 5
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Example 5
O3 ozone
O: 6 x 3 = 18
two completelyequal
arrangements
the real structureis an average of these two
where each bond is sharing 3 electronsinstead of 4 or 2
O O O
O O O
Resonance Structures
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Resonance Structures
resonance bonding between atomsthat cannot be represented in onLewis structure
show all possible structures withdouble-ended arrow in between toshow that electrons are delocalized
O O O O O O
Example 6 NO3
1-
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Example 63
N: 5 x 1 = 5
O: 6 x 3 = 18
total = 23 + 1 = 24
Covalent Network Bonding
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Covalent Network Bonding
a different type of covalent bonding not specific molecules
lots of nonmetal atoms covalently
bonded together in a network in alldirections
example:
diamond silicon dioxide
graphite
Bell Ringer
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Bell Ringer
Draw the Lewis Structure for
XeF4
I3-
PCl5
RnCl2
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Ionic Bonding
Ionic Compounds
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Ionic Compounds
ionic bonds do NOT form molecules chemical formulas for ionic
compounds represent the simplest
ratio of ion types made of anions and cations
Ionic Compounds
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Ionic Compounds
combined so that amount of positiveand negative charge is equal
usually crystalline solid
formula of ionic compound dependsof the charges of the ions combined
Formation
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Formation
attractive forces: oppositely charged ions
nuclei and electron clouds of adjacent ions
repulsive forces: like-charged ions
electrons of adjacent ions
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Formation
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specific lattice patterncreated depends on: charges of ions
size of ions
Calcium Bromide:
each Ca2+ is surrounded by 8 F-
each F- is surrounded by 4 Ca2+
Sodium Chloride
each Na+ is surrounded by 6 Cl-
each Cl- is surrounded by 6 Na+
Lattice Energy
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Lattice Energy
energy released whenseparate gaseous ionbond to form ionic solid
the larger the amountof energy released, the
stronger the bond since it is released, the
value is negative
NaClNaCl -787.5-787.5
NaBrNaBr -751.4-751.4
CaFCaF22 -2634.-2634.77
CaOCaO -3385-3385
LiClLiCl -861.3-861.3
MgOMgO -3760-3760
KClKCl -715-715
Ionic vs Molecular
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Ionic vs. Molecular
ionic bonds and molecular bonds areboth strong ionic bonds connect all ions together
molecules are more easily pulled apartbecause intermolecular forces are weak
Ionic vs Molecular
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Ionic vs. Molecular
Molecular Compounds: low melting and boiling points
many are gases at room temperature
Because the intermolecular forces of themolecules are weak so they are easily
separated
Ionic vs Molecular
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Ionic vs. Molecular
Ionic Compounds: higher melting and boiling points
all are solid at room temperature
hard: Because of the strong forces, it isdifficult for one layer of ions to movepast another
brittle: if one layer is moved, the layerscome apart completely
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Bell Ringer
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Bell Ringer
Why is waters structurebent and not linear?
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VSEPR Theory and
Molecular Shapes
VSEPR Theory
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VSEPR Theory
V alence
S hell
E lectron
P air
R epulsion
repulsion betweenpairs of electronsaround an atom cause
them to be as far apartas possible
used to predict thegeometry of molecules
Molecular Shapes
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Molecular Shapes
diatomic molecules will always belinear
all other molecules can have
different shapes based on thenumber of charge clouds around thecentral atom
charge clouds include: bonding pairs
lone pairs
2 Charge Clouds
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2 Charge Clouds
no lone pairs:
linear CO2
3 Charge Clouds
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no lone pairs:
trigonal planar
CH2O
1 lone pair: bent
SO2
4 Charge Clouds no lone pairs: CH4tetrahedral
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1 lone pair: NH3
trigonal
pyramidal
2 lone pairs: H2O
bent
5 Charge Clouds
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5 Charge Clouds
no lone pairs:trigonal
bipyramidal PCl5
1 lone pair:
seesaw
SF4