Chem 131 Exam 3 🤬

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When is it appropriate to form double or triple bonds in a Lewis structure?

When the central atom does not have an octet of electrons

How many covalent bonds are usually formed by the element C?

four covalent bonds because it has four valence electrons and needs four more to complete its octet. (common pattern for atoms in Group 14)

State the number of valence electrons in a chlorine atom.

7

State the total number of valence electrons in O₂.

12 (one oxygen has 6)

If the electronegativity difference is small the bond is

nonpolar covalent

Nonpolar Geometries (symmetric)

Linear, Trigonal Planar, Tetrahedral, Trigonal bipyramidal, Octahedral, Linear, Square planar

Polar Geometries (asymmetrical)

Bent, Trigonal pyramidal, Seesaw, T-shaped, Square pyramidal, Distorted octahedral

Bonding Order

Non-bonding - antibonding / 2

What elements can form an expanded octet?

silicon, phosphorus, sulfur, chlorine, bromine, and iodine

how many electron groups in AB2?

2

Steric # =

lone pairs + bonded atoms

2 bonds, 0 lone pairs, 180 degrees (AB2) (Ex: BeCl2, HgCl2, CO2)

linear

120 degreess, sp2, 3 bonds, 0 lone pairs (AB3) (Ex: BF3, CO32-, SO3)

trigonal planar

2 bonds, 1 lone pair, slightly less than 120 bond angle, AB2E type (Ex. SO2)

Bent/V-shaped

109.5 degrees 4 bonds, 0 lone pairs (AB4) (Ex: CH4, XeO4, PO43-) (OR 3 bonds, 1 lone pair for some electron geometries)

tetrahedral

107.5 degrees, 3 bonds, 1 lone pair (AB3E) (Ex: NH3)

Trigonal Pyramidal

104.5 degrees 2 bonds, 2 lone pairs (AB2E2) (Ex: H2O)

Tetrahedral bent

4 bonds, 1 lone pair (AX4E) (ex: SF4)

Seesaw

3 bonds, 2 lone pairs, about 90 bond angle, AB3E2 type Ex. ClF3

T-shaped

2 bonds, 3 lone pairs (AB2E3) (Ex: XeF2)

OR 5 bonds, 0 lone pairs

linear (trigonal bipyramidal)

6 bonds, 0 lone pairs (AB6) (Ex: SF6)

Octahedral

5 bonds, 1 lone pair, (AB5E) (Ex: BrF5)

Square Pyramidal

4 bonds, 2 lone pairs (AB4E2) (Ex: XeF4)

Square Planar

bent angle with one lone pair

~120 degrees

bent angles with two lone pairs

107.5 degrees

The orientation in space of an atomic orbital is associated with

the magnetic quantum number (m1)

When a bond is broken

energy is absorbed

When a bond is formed

energy is released

Nonpolar

symmetrical

polar

asymmetrical

s orbitals (sphere)

Helium and groups 1-2

p -orbital (dumbell)

Groups 13-18

d orbital (clover)

Groups 3-12

f orbital (complex shape)

Groups 5-18 (lower levels)

Atomic Radius

Moving down: increases (size of outermost occupied orbital increases)

Moving across (left to right): decreases (Effective nuclear charge increases)

First ionization energy

Moving down: Decreasing (outermost electrons further away from nucleus)

Moving across (left to right): Increasing (Effective nuclear charge increases)

Electron affinity

Moving Up: Decreasing/more negative

Moving across (right): Decreasing/more negative (effective nuclear charge increases)

Metallic character

Moving Down: Increasing (ionization energy decreases)

Moving Across (right): Decreasing (ionization energy increases)

Ionization energy

Going Up: increases

Going Across (right): increases

Copper and Chromium Columns Exception Rule

remove s orbital electron and put it in p orbital

Bonds broken are positive because

energy is being absorbed

Bonds formed are negative because 

energy is being released

To know which resonance structure is best,

you should follow a hierarchy of rules that prioritize stability: structures where all atoms have a full octet are most stable, followed by structures with the fewest formal charges. If a choice still needs to be made, the structure with negative formal charges on the most electronegative atoms and positive charges on the least electronegative atoms is preferable.

Hydrogen Bonding

Look for H directly bonded to N, O, F

C-C bond

low electronegativity

C-F bond

high electronegativity

London Dispersion Forces

weak, temporary attractions between molecules or atoms that result from the random motion of electrons. These forces occur in all molecules and their strength depends on the number of electrons, molecular size, and shape. Larger molecules with more electrons and greater surface areas have stronger London dispersion forces, leading to higher boiling points. 

Intermolecular forces strength

ionic (metal) > hydrogen (H with N,O,F)> dipole-dipole (polar) > London dispersion forces (all)

If l is s,

0

If l is p

1

If l = d

2

if l = f

3

Steric Number (2-6)

depends on how many “collections” of bonds AND lone pairs there are surrounding the central atom

SN of 2

sp hybridization

SN of 3

sp2 hybridization

SN 4

sp3 hybridization 

SN 5

sp3d

SN 6

sp3d2

Sigma bonds are ____ than pi

stronger

Only ___ bond is sigma. The other ones (if double or triple) are all pi bonds

one

What atom becomes the central atom

Least electronegative (further to the left)

Large boiling point

small size, strong forces, if molar mass is greater

Cr to Cr⁺

[Ar] 4s1 3d5 → [Ar] 3d5

(remove from s orbital and add to d)

Cu to Cu⁺

[Ar] 4s1 3d10 → [Ar] 3d10

(remove from s orbital and add to d)

magnetic quantum (ml) is related to

orientation of atomic orbitals in space

Fewer protons means 

larger radius and weak pull