official bonding notecards

MX2 and MX2E3 (180° bond angle)

Linear (formula and bond angle)

MX3 (120° bond angle)

Trigonal Planar (formula and bond angle)

MX2E (118) and

MX2E2(104.5º)

Bent (formula and bond angle)

MX4 (109.5)

Tetrahedral (formula and bond angle)

MX3E (107)

Trigonal Pyramidal (formula and bond angle)

Bonding Rules of atoms that have no formal charge

Substitutional (similar size atoms)

Interstitial (different size atoms)

Types of Alloys

Substitution, interstitial

__ is on the right, _ is on the left.

Ionic Bonds

Transfer of electrons, metal-nonmetal,

2nd strongest type of bond. Can be

“cleaved”

F 4.0, O 3.5, N 3.0

C 2.5, H 2.1

Electronegativities that

need to be memorized

Also good to know:

Cl 3.0, S 2.5

Internuclear distance

graph (Bond length

graph)

Strength and Length is always

the “low point in the curve”

Hybridization examples

sp3

These are

sp2

These are

sp1

These are

= Valence electrons on the atom–

(number of bonds) – (number of

lone-pair electrons)

The more electronegative element

must have a negative formal if

there is formal charge created.

Less formal charge means a more

correct Lewis structure

Formal Charge formula

and rules

Double and triple bonds count

as single bond, lone pairs

count. Each letter means two

electrons around central.

s,p,p,p.

Hybridization

A single bond -

Sigma bonds (alpha level symbol) bonds

Pie bonds \pi

Double and triple

bonds.

Example: a triple

bond has one sigma

bond and two

pi bonds

= sum of the bond

energies for bonds broken

(absorbed energy) – sum of the

bond energies for bonds formed

(released energy)

Don’t use this one below!!

Finding bond energy

(DeltaH

rxn)

Formula must be

memorized for AP exam

Resonance

Same atomic linkages, different bonding

Example:

Nonpolar Covalent bonds

Electrons shared evenly.

Electronegativity

Diff. 0.0

Ex bonds/ Br-Br

Polar Covalent Compounds

Electrons shared unevenly

Electronegativity

Diff. of 0.1 and up

Enthalpy of fusion

The amount of heat

required in kJ/mol

to melt

Enthalpy of vaporization

The amount of heat

required in kJ/mol

to boil

low melting points, brittle,

nonconducting as a solid and in solution.

Molecular –

Variable melting points, Hard and

brittle, conducting

Metallic –

High to very high melting points, Hard

and brittle (can undergo clevage), nonconducting

solid, conducting liquid (molten), conducts when

dissolved (must be soluble)

Ionic –

Very high melting points,

Very hard, usually nonconducting solid, Diamond

(C) and Quartz (SiO2) are examples

Covalent Network –

Interstitial

If it has C or B in it, it’s

Force of attraction, Found in

nonpolar molecules esp.

More electrons = more

polarizing of the electron

cloud = stronger London

dispersion forces!

Example/ ALL molecules

have them!!!

Note: examples of molecules with London

only He, all noble gases, CH4 (non-polar

molecule therefore no dipole-dipole!)

Intermolecular Force:

London Dispersion

forces

Found in polar molecules.

Results from diff in E.N.

usually stronger than

London forces unless lots

of e- are present

Note: molecule must have overall dipole due

to geometry! Example CH3Cl or HCl

Intermolecular Force:

Dipole-Dipole force

Results from H bonded with

O,F,N. Strongest of three

forces usually. NOT

ACTUALLY A “BOND”.

Watch out for C-H tricks!

Intermolecular Force:

Hydrogen Bonding

Stronger forces, lower vapor

pressure at a given temp.

Example: H2O is less than

C8H18

Note: H2O - – London dispersion, dipole -dipole.

and hydrogen bonding

C8H18 – London dispersion only (non-polar

molecule)– high vapor pressure – becomes vapor

easily

Effect of intermolecular

forces on Vapor

Pressure

As molar mass increases in

hydrocarbons, B.P increases. --

Reason – Stronger London

Dispersion Forces due to more

atoms=more electrons, which

can be more polarized!

CH4 = -1640C, C8H18 = 1250C

***hydrocarbons are

nonpolar**

Effect of intermolecular

forces on Boiling Point

Covalent Network Solid

Strongest type of bond

Unique Covalent Bonding –

VERY HIGH M.P.

Examples/ Diamond, graphite,

SiO2

Molecular can contain polar bonds but be

nonpolar molecules: A General Rule: Molecules

with no lone pairs of electrons on central atom M,

as well as MX2E3 (linear) and MX4E2 (sq. planar):

Are nonpolar because bond polarities cancel out!

Examples – CO2 contains polar bonds but due to

linear geometry the polar bonds cancel out.

Polar Bonds but a

nonpolar molecule?

As the number of electrons around the

central atom increases the bond angles

decrease due to increased electron

repulsion. Lone pairs exhibit more

repulsion than bonding electrons.

Example: CH4 (no lone pairs) – 109.5,

NH3 – (1 lone pair) - 107, H2O (two lone

pairs) – 104.5

Effect of lone pairs on

bond angle

When a substance melts or boils

intermolecular forces are broken, NOT

intramolecular forces. Example: When

water boils, the hydrogen bonds (an

intermolecular force) are broken

between water molecules but the

covalent bonds (an intramolecular force)

within the water molecule between O

and H are NOT broken. Otherwise

heating water would result in H2 gas and

O2 instead of “steam/water vapor”

Understanding phase

changes! In terms of

intra and inter

molecular forces.

SUPER IMPORTANT!!!

A force within the

molecule. The bond

between H-H (in H2)

is one

Intramolecular force

A force between atoms and

molecules. The London

dispersion forces of one H2

molecule effecting another H2

molecule

Intermolecular force

As the number of electrons

around the central atom

increases the bond angles

decrease due to increased

electron repulsion and

because of decreased space!

Since electrons have the same

charge they repel as far away

from one another as possible.

Effect of bonding

electron pairs on bond

angle

Single Bond

Longest and weakest

Example: C-C SINGLE bond

length=154pm, bond energy = 348

kJ/mol

Triple Bond

Shortest and longest

C-C TRIPLE bond length=120pm,

bond energy = 839 kJ/mol

Left carbon – tetrahedral, 109.50 bond angles, sp3,

4 sigma bonds

Right carbon – trigonal planar, 120o, sp2 (double

bond counts as single for hybridization), 3 sigma

bonds, 1 pi bond.

Rightmost oxygen – Bent, 104.50, sp3 (lone pairs

count for hybridization), 2 sigma bonds

Determination of

geometry/bond

angles/hybridization/t

ypes of bonds with

respect to larger

molecules for

individual atoms.

When dry ice sublimes no

covalent C=O bonds are

broken (intramolecular forces)

but intermolecular forces are

broken (London dispersion

forces only since CO2 is

nonpolar due to geometry)

Sublimation of CO2 with

respect to intra and

inter molecular forces

Hydrogen Bonding in

water. The dashed line

is the hydrogen bond

1

2

3

Bond orders are: Single bond =

Double bond =

Triple bond =

1.5

Nonwhole number bond orders are possible in

resonance.

All the O---O bonds have a bond order of ___ in

the resonance structures of O3