Chem 110 McGill SOS

Chem 110McGill SOS

• Midterm 2

Topics on Midterm II Chapter 10+11+12

Remember, only 15% of the questions are going to be from midterm 1! Don’t overdo it!

Chapter 10 – Bonding I

Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies

3 types of bonding

• Covalent = sharing of electrons

• Ionic = transfer of electrons

• Metallic = delocalization of electrons

Covalent bonding

• pure covalent = same element, equal sharing

• polar covalent = different elements (non-metals), unequal sharing

Polar covalent = think Dipole moment

Arrow always points from the partial positive charge (δ+) to partial negative charge (δ-)

electronegativity differences determine the degree of polarity

Electronegativity a measure of the electron-

attracting power of a bonded atom type of bond determined by

difference in EN remember, 1.7 is arbitrary!

Ionic -greater than 1.7 Polar Covalent - between 0.5-1.7 Pure Covalent - less than 0.5

Lewis theory Valence electrons are most

important in chemical bonding Why do atoms bond? To acquire

a stable octet. Remember, eight is yummy!

Lewis theory -terminology

Octet Rule: Each atom must have 8 valence shell electrons in a lewis structure

Bonding pair vs. lone pair Single, double and triple covalent

bonds

Lewis Symbols

Chemical symbol in the middle

Dots represent valence electrons

Lewis Structures - Ionic Compounds

Lewis Structures - Covalent Compounds

Strategy 1. Count total # valence electrons Add/subtract additional charges

+ve charge, remove e- -ve charge, add e-

2. Identify the central and the terminal atoms (any carbon or hydrogen?)

3. Join atoms through SINGLE covalent bonds. This is your skeletal structure

4. Achieve octet around each atom For remaining valence e-, add to central

atom. Add double and triple bonds when necessary to complete an octet.

Lewis Structures - Tips

Hydrogen is always a terminal atom Carbon is always a central atom Central atom (if not carbon) will be

the least electronegative element Go for a compact, symmetrical

structure

Example: HCN

Step 1: Count total # valence e-

H=1, C=4, N=5

So…1+4+5=10 e-

Example: HCN Step 2: Identify the central and

the terminal atom (any carbon or hydrogen?)

Carbon=central atom Hydrogen= terminal atom therefore Nitrogen= also

terminal atom

Example: HCN Step 3: Join atoms through SINGLE

covalent bonds. This is your skeletal structure.

Example: HCN Step 4: Complete the octet of terminal

atoms. Count the number of valence

electrons remaining. If you have any remaining, use them on the central atom. If your central atom does not have a complete octet, start switching single covalent bonds to double (or triple if necessary)

Formal charges

• helps us draw Lewis structures

• tells us where the e- are located

• F.C. = # valence e- - # lone pair e- - 1/2 # bonding e-

• NOTE: sum of all F.C in molecule must be equal to overall charge!

Formal Charge - Question Types

2 types of questions: 1. what is the formal charge on

each atom?Use FC equation

Example: CHO2(-) Draw structure out, if necessary Use F.C. = # valence e- - # lone pair e-

- 1/2 # bonding e-

F.C. (H) =1-0-0.5(2)=1-0-1=0

F.C. (C) =4-0-0.5(8)=4-0-4=0

F.C. (O)-SB =6-6-0.5(2)=6-6-1=-1 F.C. (O)-DB =6-4-0.5(4)=6-4-2=0 overall F.C. = -1 = charge on

molecule

Formal Charge - Question types

II. Use formal charges to see if the lewis structure is plausible or not?

Rules: Sum of the formal charges must equal 0

exceptions: the molecule is an ion (like in the last example)

Keep structures as small as possible Negative formal charges are usually on the

most electronegative atoms Positive formal charges are usually on the

least electronegative atoms

Example: NH2CN NH2CN has two possible Lewis

structures How do we know which one is

more plausible?

Use formal charges!

Resonance forms Two or more plausible lewis structures

can be written for the same molecule. Note: They will have the same

skeletal structure. It’s only the electron distribution that changes!

Each lewis structure contributes to the resonance hybrid.

increasing resonance = more stable

Exceptions to the Octet Rule

Odd-electron species total # of valence electrons is odd use formal charges to determine where to

place the single electron Incomplete octets in order to avoid breaking a formal charge rule

(i.e. negative formal charges are usually on the most electronegative atoms), you break the octet rule

usually only seen with compounds containing Be, B, Al

Expanded valence shells seen with elements containing d-obitals like S,

P + bonded to a highly electronegative atom like O,F,Cl

VSEPR Theory concerns shapes of molecules Electron pairs (either lone pair or

bond pairs) assume an orientation about the atom that will minimize repulsionstherefore we want to maximize the

distance between electron pairs order of repulsion interactions:

Lone pair-lone pair = most repulsive lone pair-bond pair bond pair-bond pair = least repulsive

VSEPR Theory - Tips

Molecular shape depends on: number of electron pairs type of electron pair (lone pair vs

bonding pair) When determining the molecular

geometry, you must first determine the electron group geometry!

Treat multiple bonds (double, triple) as if it were a single bond

KNOW THIS: Molecular Geometries

2 electron groups▪ 0 lone pairs=linear

3 electron groups▪ 0 lone pairs= trigonal-planar▪ 1 lone pair= bent

4 electron groups▪ 0 lone pairs= tetrahedral▪ 1 lone pair= trigonal-pyramidal▪ 2 lone pairs=bent

5 electron groups▪ 0 lone pairs=trigonal-bipyramidal▪ 1 lone pair=seesaw▪ 2 lone pairs=T-shaped▪ 3 lone pairs=linear

6 electron groups: ▪ 0 lone pair=octahedral▪ 1 lone pair=square-pyramidal▪ 2 lone pair=square-planar

VSEPR Theory Step 1: Count the # of electron

groups (the number of bonded atoms+the number of lone pairs)

Step 2: Determine electron group geometry 2 electron groups: linear 3 electron groups: trigonal-planar 4 electron groups: tetrahedral 5 electron groups: trigonal-

bipyramidal 6 electron groups: octahedral

VSEPR Theory Step 3: Count the number of lone

pairs. If there are no lone pairs, the electron group geometry IS the molecular geometry.

Step 4: Determine molecular group geometry. each additional lone pair

changes the molecular geometry

Example: Sulfur Tetrafluride

5 electron groups Electron geometry is

trigonal bipyramidal However, because 1

of our electron groups is a lone pair, the molecular geometry will be seesaw

Bond engergies, order, and bond

length Bond energy is the energy required to break the bond between atoms single covalent bond= weakest Triple covalent bond = strongest

Bond order Single covalent bond=1 Double covalent bond=2 Triple covalent bond=3

Bond length is the distance between the centers of the two atoms joined by a covalent bond

NOTE: As bond order increases, bond energy increases and bond length decreases.

Chapter 11- Bonding II

Valence Bond Theory Hybridization M.O. Theory Metallic Bonding

Valence Bond Theory

In a molecule, electron probability is highest where the orbitals of the atoms overlap to form covalent bondsoverlap allows electrons greater

freedom across both orbitals thus greater orbital overlap =

more stable (think about single vs double vs triple bonds)

Valence Bond Theory - types of

hybrid bonds σ bond: most of e- density around bond axis formed by overlap of

s-s, s-p, or p-p head-to-head overlap

π bond: e- density above and below the bond axis formed by overlap of

p-p side-to-side overlap

Valence Bond Theory - multiple

bonds Single covalent bond= 1 sigma bond

Double covalent bond=1 sigma bond (σ)+ 1 pi bond (π)

Triple covalent bond=1 sigma bond (σ) + 2 pi bonds (π)

NOTE: add π bonds only after σ bonds! π bonds only for double or triple bonds

Practice QuestionThe resonance structure of PO4(3-) where the

central atom carries a formal charge of zero issurrounded by:

A. 4 sigma bonds B. 4 sigma and 1 pi bonds C. 4 sigma and 2 pi bonds D. 4 sigma and 3 pi bonds E. 4 sigma and 4 pi bonds

Practice question - steps

• draw Lewis structure

• remember to complete octets and use F.C. to determine best structure

• count single and multiple bonds

• single bond = 1 sigma

• multiple bonds = 1 sigma + additional pi

Hybridization Bonded atoms do not have the same

orbitals as non-bonded, isolated atoms. They have hybrid orbitals.VSEPR -> geometryValence bond theory -> bonding

Number of hybrid orbitals= Number of electron groups

NOTE: lone pairs are also hybridized! Double and triple bonds are not!

Hybrid OrbitalsNumber of Electron Groups

Electron Group Geometry

Molecular Group Geometry

Hybrid Orbitals

2 Linear Linear sp+2p

3 Trigonal planar Trigonal PlanarBent

sp^2+p

4 Tetrahedral TetrahedralTrigonal pyramidalBent

sp^3

5 Trigonal-Bipyramidal

Trigonal BipyramidalSeasawT-shapedLinear

sp^3d

6 Octahedral OctahedralSquare PyramidalSquare Planar

sp^3d^2

Hybridization Question: What molecular shapes

can a sp^3d^2 hybridized molecule have?

Hybridization Answer: Octahedral, Square

Pyramidal and Square Planar

Hybridization - tips You need the electron group

geometry to determine the hybrid orbitals

to determine the electron group geometry, you need to draw a lewis structure

Lewis structure-> VSEPR->Hybridization

Example CO2

Steps: Draw lewis structure Determine electron group geometry think: hybridization using chart you

will have memorized!

Molecular Orbital Theory

Compensates for Valence Bond Theory in explaining magnetism

Used for diatomic molecules Excellent for predicting

magnetic properties (diamagnetic and paramagnetic)

M.O. theory -> think: molecular orbitals

Molecular orbitals

Two types of molecular orbitals arise from the combination of atomic orbitals: Bonding Molecular Orbitals Antibonding Molecular Orbitals

• Bonding M.O = increase e- density

• maximum e- density between 2 nuclei

• Anti-bonding M.O = decrease e- density

• no e- density between 2 nuclei

Key principles of the M.O. Theory

# of atomic orbitals combined = # of molecular orbitals formed

Antibonding M.O have higher energy levels than bonding M.O

In ground state electron configuration, electrons enter lowest energy M.O available Use Hund’s Rule + Pauli’s exclusion

principle: They fill up M.O orbitals singly before pairing up

Stabilization energy• differs from molecule to molecule

• presence bonding e- = stabilize molecule

• presence antibonding e- = destabilize molecule

Bond Order Bond Order= 1/2 (# bonding e- -

# antiboding e- ) NOTE: bond order always > 0! For first period elements

(hydrogen and helium), we only use σ1s M.O and σ*1s M.O

For second period elements (lithium, beryllium…), the M.O. listed above are already filled and we concern ourselves with the following M.O.s: σ2p, σ*2p, π2p(x2), π*2p(x2)

Metallic Bonding - Electron Sea Model

Properties of metals

• luster of reflectivity

• high electrical conductivity

• high thermal conductivity

• malleability & ductility

• electronic emission

Band Theory In conductors, the valence

band and the conduction band overlap one another. No energy gap!

In semiconductors, the valence band and the conduction band are seperated by a small energy gap

In insulators, the valence band and the conduction band are seperated by a large energy gap!

Semiconductors Intrinsic semiconductors have a

energy gap with a FIXED size

Extrinsic semiconductors have a energy gap whose size can be controlled via doping

Doping =adding of impurities. It increases conductivity.

Semiconductors

Two types of extrinsic semiconductors: N-type▪ Migration of electrons from donor atoms (i.e.

phosphorus)▪ absorbs blue light ▪ appears yellow

P-type▪ Migration of positive holes from acceptor atoms (i.e.

boron)▪ absorbs red and green light ▪ appears blue

Chapter 12 – Intermolecular Forces

London dispersion forces Dipole-dipole Interactions Hydrogen Bonding Ionic and Network Covalent Solids

Intermolecular forces

forces between molecules intermolecular forces keep molecules close together

explains physical properties of solids, liquids, and gases

THINK: attraction between molecules

Intermolecular forces - physical

properties• when comparing substances of similar nature (elements in same group):

• melting point

• boiling point

• increases with increasing molecular size (weight)

• also affected by molecular shape

• THINK: hydrocarbons

London dispersion forces

dominant force in nonpolar molecules due to small, instantaneous dipole moments

THINK: electrons are always in motion London forces are the weakest of all

intermolecular interactions present in any molecule increases with molecular size

Example Which of the following has a higher boiling

point?

n-pentane

Dipole-dipole Interactions

Dipole-dipole forces occur in polar molecules (unequal sharing of electrons) due to dipole moments

For example, HCl is polarchlorine is partially (-), and hydrogen is

partially (+)

!!!

• NOTE: dipole-dipole interactions do not replace dispersion forces

• they occur in addition to them!

Example 1Which of the following would

have the higher boiling point?a)CH3CNb)CH3CH2CH3

Example 2

• Predict the trend in boiling points of HCl, HBr and HI

Logical reasoning...?

• determine polarity by comparing electronegativity of Cl, Br, and I

• therefore order of increasing b.p: HI < HBr < HCl

Trick question!!

• for large (heavy M.W) molecules, london dispersion forces are stronger than dipole-dipole forces

• therefore we determine increasing b.p by looking at M.W and not polarity!

• correct answer: HCl < HBr < HI

Hydrogen Bonding

occurs with small molecules with Hydrogen bonded to N, O, or F

have much higher melting and boiling points than expected from M.W

why? Hydrogen is attracted to the lone pairs on O, N, and F

Network covalent solids

• THINK: diamonds are forever

• each molecule is held together by strong, covalent bonds

• strong directionality = strong material

Ionic solids

• made up of oppositely charged ions that interact by Coulombic forces between them

• think: salt! Made up of sodium and chloride ions

• crystalline ionic solids: highly organized fashion of ions

Terminology• Unit cell -> lattice -> crystal

• most ionic solids are formed of cubic unit cells

• 6 faces

• 8 corners

• 12 edges

• 3 types of cubic unit cells:

• simple cubic

• body centered

• face centered

Simple cubic unit cell• ions of the same type present at

corners of the cube

• how many ions present per unit cell?

• each corner is shared with 3 other corners (1 -> 1/2 -> 1/4 -> 1/8)

• therefore ion shared in corner = 1/8 inside each unit cell

• total in simple cubic unit cell:

• 8 corners x 1/8 ion = 1 ion per unit cell

Body centered cubic unit cell

• ions of the same type present inside the body and at the corners of the cube


• for corners: 8 corners x 1/8 ion = 1 ion per unit cell

• PLUS 1 ion inside not shared

• total = 2 ions / body centered cubic unit cell

Face centered cubic unit cell

• ions of the same type present at corners and faces of the cube


• corners: 8 corners x 1/8 ion = 1 ion

• faces: each face is shared with 1 other face

• 6 faces x 1/2 ion = 3 ions

• total = 1 ion for corners + 3 ions for faces = 4 ions per face centered cubic unit cell

Example: NaCl

• what type of unit cell is a crystalline ionic solid of NaCl?

NaCl - steps

• look at a single unit cell

• find ions of the same type

• simple, body-centered, or face-centered?

• larger ion occupies the corners, smaller ion comes in to fill the spaces

Coordination number

• How many ions surround those of another type?

• in NaCl, each Cl- is surrounded by 6 Na- ions, and vice versa

• therefore coordianation number = 6

• ratio of cation : anion = 1:1

Chem 110 McGill SOS

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