Lewis Structures: 5 steps1. Count valence e- available

If an ANION, add charge to # valence e- If a CATION, subtract charge from # valence e-

2. Draw skeletonLeast electronegative atom is the central

atom(usually first element in formula, never H)Place surrounding atoms around the central

atomadd dashes to show bonds from central

atom to surrounding atoms

Lewis Structures-cont’d

3. Satisfy the octet rule for each surrounding atom using valence electrons. Note that H can only have 2 e-

4. Any remaining valence e- go on the central atom. 5. Check final structure

If the octet rule is not satisfied for the central atom, borrow e- pairs to make double or triple bonds.

It is possible for the central atom to have an expanded octet, (eg. SF6) or a contracted octet, (eg. BF3). Theses are exceptions to the octet rule.

Lewis Structures– H2O

1. # val e- available = 2 H = 2(1) = 2

1 O = 1(6) = 6= 8 e- available

2. Skeleton (O is central…)H O H

Lewis Structures– H2O

3. Add dots so that each has 8 e-Each dash is 2 e-

4. Count e- shown : 2 dashes = 4 e-

4 dots = 4 e-

= 8 e- total

5. Check:

8 e- shown = 8 e- available

So OK!!

Lewis Structure, C2H4

1. # e- available = 2 C = 2(4) = 8

4 H = 4(1) = 4= 12 e- available

2. Skeleton (C---C is central…)

Lewis Structure, C2H4, cont’d

3. Add dots so that each has 8 e-

Remember each dash = 2 e-

4. Count e- shown: 4 C-H bonds = 4 (2) = 8 e-

1 C-C bond = 1 (2) = 2 e-

4 dots = 4 e-

total = 14 e-

Lewis Structure, C2H4, cont’d

5. Compare # e- shown to # e- available14 e- shown >12 e- available

So must make a double bond:

• Remove 1 electron from each C

•Combine remaining 2 e- in a covalent bond, forming a double bond

Lewis Structure, C2H4, cont’d

Check: # e-shown = # e- available

12 = 12 OK!!

4 C-H bonds = 4 (2) = 8 e-

1 C=C bond = 1 (4) = 4 e-

total = 12 e-

(or count # lines = 6

6 lines *2e- each = 12 e-)

Lewis Structures Exercise

1. Get in pairs2. Molecules assigned3. Give “model kits”4. Each group use model to make

Lewis Structure5. Show model, draw LS on board

Lewis Structures Exercise- compounds

CH4 N2O NO3 –

NO2 PCl5 NH4 +

SO2 SF6

CO2 SO4 2-

NH3 SO3 2-

XeF4 NO2 –

Molecular Shape – VSEPR Theory

• Can relate Lewis Structure to 3-D shape of a molecule

ValenceShellElectronPairRepulsion Theory

e- pairs arrange themselves around a nucleus to minimize -/- repulsions

e- pairs get as far away from each other as possible and as close to the nucleus as possible

VSEPR Theory

• Use “AXnEm” designation

A = central atomX = atoms bonded to An = # atoms bonded to E = unshared e- PAIRS on A (lone pairs)m = # unshared e- pairs on A• Get from Lewis Structure!!

AXE designation

Ex— water, H2OA = “Oxygen

AX2E2

X = H

n = 2

E = pairs of dots on O

m = 2 pairs

Use AXE to give shapen+m AXE example shape Bond angle

2 AX2E0 CO2 linear 180o

3 AX3E0 BF3Triangular

planar120o

3 AX2E1 SO2 bent <120o

Use AXE to give shapen+m AXE example shape Bond angle

4 AX4E0 CH4 Tetrahedral 109.5o

4 AX3E1 NH3Trigonal

pyramidal< 109.5o

4AX2E2

H2O Bent < 109.5o

Use AXE to give shape

n+m AXE example shape Bond angle

5 AX5E0 PF5

Triangular bipyramida

l

180o, 90o, 120o

5 AX4E1 SF4 See-saw90o, 180o,

<120o

5 AX3E2 ClF3 T-shape 90o, 180o

5 AX2E3 XeF2 linear 180o

n+m = 5 family

Use AXE to give shape

n+m AXE example shapeBond angle

6 AX6E0 SF6 Octahedral 180o, 90o

6 AX5E1 BrF5Square

pyramidal180o, 90o

6 AX4E2 XeF4Square planar

180o, 90o

n+m = 6 family

Be able to…

1. Draw a Lewis structure for any assigned molecule

2. Based on the Lewis Structure, give the AXE designation.

3. Determine molecule shape, bond angles, and polarity.

Molecule Polarity

• Tell if molecule has one “side” that the electrons like to congregate…

• Based on molecule shape and bond polarity

Bond Polarity

• Even though electrons are shared between two nuclei in covalent bonds, often the sharing is NOT – One atom often has greater affinity for e-

than other

• Look at differences in electronegativity (EN)

• The more EN the atom, the more it “hogs” the shared e-

Bond PolarityWhich is more EN? F or H?

FMake the “bond” an arrow

pointing toward the more EN atom

Put a “+” across the tail (other end) of the arrow

So e- spend more time around F

So F has a “partial negative” charge d-

H has a “partial positive” charge, d+

F H

F H

F H

F H- +

Bond Polarity

• The 2 shared e- between H and F tend to spend more time around F

Molecule Polarity

• Consider both bond polarity AND shape!

Is Water a polar molecule?1. Draw the Lewis Structure

2. Get shape: AX2E2…so BENT

Is Water a polar molecule?3. Look at each OH bond—determine bond polarity O H

- +

-

++

Is Water a polar molecule?

4. Rotate to look end on / smash into page

-

++

-

+

5. If different (+/-), then POLAR

(if same +/+ or -/-, then NONPOLAR)

Are the following molecules polar?

• CO2

• CCl4• CHCl3

Intermolecular Forces(IMF)

• Attractive forces between 2 or more molecules

• Need to consider molecule shape and polarity

3 types of IMF

1. Dispersion forces: between NONPOLAR molecules

(weakest)

2. Dipole: between POLAR molecules(intermediate)

3. Hydrogen bonds: special case of dipole forces, between H in one molecule and O, N, or F in another

(strongest)

Dipole Forces

• +/- attraction between POLAR molecules

Partial + on one molecule attracted to the partial – on a neighboring molecule

Hydrogen Bonds

• Special case of dipole attractions• In each molecule, must have H

bonded to O, N, or FPartial + (H) on one molecule attracted to the partial – (O, N, or F) on a neighboring molecule

Hydrogen bonds

• Individually, each H-bond is weak (compared to a covalent bond)

• Collectively, H-bonds are VERY strong, especially in large molecules like proteins or DNA

Dispersion Forces• Between 2 non-polar molecules• “temporary dipole – induced

dipole”“Temporary Dipole”

Dispersion Forces

The “temporary dipole” now “induces” a neighboring molecule to become a dipole

(- pushes e- away from it in the neighbor, making that end of the neighbor +)

Dispersion force = +/- attraction between temporary dipole and induced dipole

IMFs and States of Matter

• Think of the ability of a material to change phases as a measure of the strength of IMFs

Solids

Melting (s l)

• As Temperature increases, molecules vibrate/move/bounce more and more

• Gain enough Kinetic energy (KE) to overcome some IMFs

• Molecules can now slide around one another

Liquids

Boiling (l g)

• As temperature increases, molecules gain more and more energy

• Soon overcome all IMF• Molecules no longer “attached” to

each other• Escape to the gas phase

Gas

Solutions

• Homogenous mixture of 2 or more substances

• Solute = material being dissolved, – Usually a solid– Present in least amount

• Solvent = material in which the solute is being dissolved– Usually a liquid– Present in greatest amounts

Aqueous Solutions

• Solvent = water• Solute = some solid

Dissolving Process

• Think of what happens on a molecular level

Before mixing…

Interact w/each other

Strong IMFs

Close packed materials (solid)

Interact w/each other

medium IMFs

Liquid- so molecules still slide around

During Mixing…

Replace IMFs from like molecules with IMFs from others

Now each solute has an IMF interaction with a solvent molecule

“Like Dissolves Like”

• If the IMFs between a solute and a solvent are similar, then the solute will dissolve in the solvent!

• Ex- NaCl in H2O

– NaCl is polar (ionic) H2O is polar (H-bonds)

– Similar IMFs, so NaCl will dissolve in water

Will CH4 dissolve in H2O?

• CH4 is nonpolar… IMF = dispersion forces

• H2O is polar…..IMF = H-bonds

• Different IMFs so CH4 will NOT dissolve in H2O