CH301 Unit 2 Exam Review

Duet Rule

Hydrogen achieves the greatest stability when it’s valence shell contains 2 electrons. Different from the octet rule where other elements become stable upon achieving 8 valence electrons.

Bond Polarity

Is proportional to the length of the dipole between 2 atoms. Greater ΔEN, the greater polarity of the bond.

Bond Polarity v. Molecular Polarity

Molecules can have polar bonds and be non-polar molecules. Molecular polarity is the overall distribution of electron density in the molecule, while bond polarity is how uneven the sharing of electrons are between 2 atoms in the molecule. Note that if perimeter atoms are not the same, the molecule is l

Coulomb’s Law

Attraction gets stronger when charges are larger and ionics can compact closer together, which is possible when ionic radius is smaller.

Lattice Energy

Energy in the bonds of a crystalline solid salt. Compare charge first, THEN compare ionic radius ONLY IF charges are the same. Higher charges and smaller ions means stronger attraction and more LE.

Drawing 2D Lewis Structures

1. Add up the valence electrons in a molecule

2. Spread out perimeter atoms around the central atom

3. Put either 2 (for H) or 8 (for all others) electrons around perimeter atoms

4. Dump the rest of the unused valence electrons in the middle (central atom)

5. If the central atom does not have 8 or more electrons (Octet rule), share e- from perimeter atoms to make multiple bonds (double or triple)

DO NOT DO S = N -- A. It does not work all the time.

Octet Rule

Atoms prefer to obtain noble gas configuration (filled valence shells).

Electronegativity Trend

Increases and highest at the UPPER RIGHT of the table. Shielding increases down a group, meaning electrons aren’t pulled as closely to the nucleus, which directly affects EN.

In what cases is ΔEN = 0?

Br2, I2, N2, Cl2, H2, O2, F2 = All exist as diatomics.

Ionic Radius Trend

Decreases from left to right, and increases down a group. Largest in the LOWER LEFT corner of the table.

Metallic Character Trend

Most metallic in the LOWER LEFT corner of the table. Cs is the most metallic and F is the least metallic (excluding noble gases).

ENC Trend

Increases left to right, and remains constant down a group. ENC equation: # of protons - # of shielding electrons.

Ionization Energy Trend

Increases left to right, and decreases down a group. Largest in the UPPER RIGHT of the periodic table.

Naming Ionic Compounds

Leave cation as is + polyatomic ion + anion base name ending in -ide.

Transition Metal charges to know.

Ag —> 1+

Zn —> 2+

Al —> 3+

Hypovalency

Electron deficiency around the central atom. Will commonly be Be and B. Boron forms up to 3 bonds (3ve-), and Beryllium forms up to 2 bonds (2ve-).

Radicals

Odd number of valence electrons. Add up the number of valence electrons to determine if the molecule is radical. They are VERY reactive, and therefore unstable, so it’s correct for the central atom to not achieve the octet rule. Note that the radical electron is on the LESS EN atom.

Hypervalency

Some elements can have expanded octet rules. Elements with an n value of 3 or higher have access to the d-orbitals. Think 3d^10 for n=3 shells.

Regarding charges when drawing Lewis Structures

With anions, add the charge to the total e- count, and subtract the charge from the total e- count in the case of cations.

Deciding Central Atom of Lewis Structures

- Typically first element in the formula.

- The element that is the LEAST electronegative.

- Halogens are usually perimeter atoms

- H always outer/perimeter atom.

Formal Charges

# of valence electrons - # of dots - # lines = FC

- FC of 0 is ideal.

- FC of +2/-2 is not ideal.

- FC of -1 is preferable on the relatively more electronegative atom in the molecule if unavoidable.

Resonance

When 1 molecule has multiple Lewis Structures. Each is equivalent in energy and stability to one another. Formal charge for each is effectively minimized, and totals the charge on the polyatomic.

Defined Resonance Structures in CH301

O3, NO2(-), NO3(-), CO3(2-), RCOO(-), NO2F, IO2(-), HCO3(-)

Tips to know for Functional Groups

For carbon atoms, they have 4 bonds and no lone pairs. Nitrogen atoms have 3 bonds and 1 lone pair, and oxygen atoms have 2 bonds and 2 lone pairs.

VESPR Theory

Electron rich regions (regions of high electron density: RHED’s) repel one another and spread out as much as possible. Lone pairs, single, double and triple bonds all count as RHED’s.

Linear

Has 2 RHED’s and is made of <180 degree angles. sp hybridization.

• Linear: 2 bond regions, 0 lone pairs, AB₂. NONPOLAR

Trigonal Planar

Has 3 RHED’s and is made of <120 degree angles. sp² hybridization.

• Trigonal Planar: 3 bond regions, 0 lone pairs, AB₃. NONPOLAR

• Bent: 2 bond regions, 1 lone pair, AB₂U.. POLAR

Tetrahedral

Has 4 RHED’s and is made up of <109.5 degree angles. sp³ hybridization.

• Tetrahedral: 4 bond regions, 0 lone pairs, AB₄. NONPOLAR

• Trigonal Pyramidal: 3 bond regions, 1 lone pair, AB₃U. POLAR

• Bent: 2 bond regions, 2 lone pairs, AB₂U₂. POLAR

Trigonal Bipyramidal

Has 5 RHED’s and <120 and <90 degree angles. sp³d hybridization.

• Trigonal Bipyramidal: 5 bond regions, 0 lone pairs, AB₅. NONPOLAR

• Seesaw: 4 bond regions, 1 lone pair, AB₄U. POLAR.

• T-Shaped: 3 bond regions, 2 lone pairs, AB₃U₂. POLAR

• Linear: 2 bond regions, 3 lone pairs, AB₂U₃. NONPOLAR

Octahedral

Has 6 RHED’s and <90 and 1 <180 degree angles. sp³d² hybridization.

• Octahedral: 6 bond regions, 0 lone pairs, AB₆. NONPOLAR

• Square Pyramidal: 5 bond regions, 1 lone pair, AB₅U. POLAR

• Square Planar: 4 bond regions, 2 lone pairs, AB₄U₂. NONPOLAR

Valence Bond Theory

Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals that yield a pair of electrons shared between the two bonded atoms. Hybridization, the creation of new orbitals, are designed to make electrons available for bonding.

Formula for orbitals in hybridization

sp —> 1 s orbital overlaps with 1 p orbital. (Linear)
sp² —> 1 s orbital overlaps with 2 p orbitals. (Trigonal Planar)
sp³ —> 1 s orbital overlaps with 3 p orbitals. (Tetrahedral)

sp³d —> 1 s orbital overlaps with 3 p orbitals and 1 d orbital. (Trigonal Bipyramidal)

sp³d² —> 1 s orbital overlaps with 3 p orbitals and 2 d orbitals. (Octahedrals)

Sigma bonds

End-on-end overlap of any 2 orbitals. They are the first connection by any 2 atoms, and can have free rotation.

Pi bonds

Side-on-side overlap of any 2 orbitals; they do NOT have free rotation, and are only p-orbitals. Double bonds have 1 pi bond, and triple bonds have 2 pi bonds.

Molecular Orbital Theory

Based on quantum mechanics, orbital wave functions combine when orbitals (waves) overlap. This theory can explain O₂’s magnetism.

Deconstructive Interference v. Constructive Interference

Deconstructive interference is when 2 different orbitals (waves) of opposing directions overlap, which is anti-bonding. Constructive interference is when 2 similar orbitals (waves) of similar directions overlap, which is bonding.

Bond Order

The number of chemical bonds present in a molecule. (# of bonding electrons - # of anti-bonding electrons)/2 = Bond Order. As BO increases, both bond strength and energy increase, while bond length decreases.

Bond Orders for Diatomic Species

H₂: 1, 2 electrons
He₂: 0, 4 electrons

Li₂: 1, 6 electrons

Be₂: 0, 8 electrons

B₂: 1, 10 electrons

C₂: 2, 12 electrons

N₂: 3, 14 electrons

O₂: 2, 16 electrons

F₂: 1, 18 electrons

Ne₂: 0, 20 electrons

Li₂ to N₂ MO order

sigma 1s, 1s, sigma 1s*, sigma 2s, 2s, sigma 2s*, pi 2p, sigma 2p, 2p, pi 2p*, sigma 2p*

O₂ to Ne₂ MO order

sigma 1s, 1s, sigma 1s*, sigma 2s, 2s, sigma 2s*, sigma 2p, pi 2p, 2p, pi 2p*, sigma 2p*

Aldehyde

Ketone

Amide