quantum chemistry (real)

8/14/2019 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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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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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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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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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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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
http://movies/Quantum_Theory.asfhttp://movies/Quantum_Theory.asf


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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
http://movies/Electron_Behavior.asfhttp://movies/Electron_Behavior.asf


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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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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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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


last electron is in sublevel: 4s

1s2 2s2 2p6 3s2 3p6 4s2

1s 2s 2p 3s

3p 4s


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Bromine


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


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



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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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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
http://college.hmco.com/chemistry/resources/shared/molecules/chime/c/ch3cl.htmlhttp://college.hmco.com/chemistry/resources/shared/molecules/chime/c/ch3cl.htmlhttp://college.hmco.com/chemistry/resources/shared/molecules/chime/n/nh3.htmlhttp://college.hmco.com/chemistry/resources/shared/molecules/chime/n/nh3.htmlhttp://college.hmco.com/chemistry/resources/shared/molecules/chime/n/nh3.html


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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
http://college.hmco.com/chemistry/resources/shared/molecules/chime/n/nacl.html


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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

quantum chemistry (real)

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