Chem 101 Unit 2

Ionic Bonding

Occurs when one atom steals the electron of the other atom, resulting in an ionic bond. No sharing, only taking

Cations

Positively charged ions have fewer electrons than their neutral counterparts. Typically, result from elements that want to give up electrons since it is easier to remove them

Anions

Negatively charged ions that have more electrons than their neutral counterparts. Typically come about from elements that want to gain electrons to complete their octet.

Deducing charges on atoms in ionic formulas

Use the criss-cross method for charges for example, NaOH would be Na^+1 and (OH)^-1 but you have to watch out for atoms with ± 2,3 charges that simplify to 1

Working with ionic compounds with polyatomic ions

Polyatomic ions are charged, and the charge applies to the combination of atoms in the polyatomic ion. Use the same criss-cross method to deduce charges

Covalent bonding as shared electrons

Covalent bonding includes sharing of electrons between atoms. No one atom takes an electron from the other. More eN atoms have a stronger pull on those electrons, pulling them closer to them and giving them a partial charge.

Assigning bond dipoles to polar covalents

Bond dipoles point in the direction of the more eN atom and only apply to N, O, F, Cl, Br, and S

Octet and duet rule

Atoms want to have a closed shell. For most atoms, this is an octet, and for hydrogen, this is a duet. Completing an octet or duet is the most stable for an atom.

Hypervalent atoms

Hypervalent atoms share more than 8 valence electrons. They typically have the same number of valence electrons in their bonds as they do in the periodic table, which is why they have so many electrons.

lewis dot structures for bonding patterns on neutral atoms

Must ensure all octets and duets are complete, lone pairs are assigned, bond order is correct for all atoms

Formal charge

When an atom has a charge due to either owning more of less electrons than its neutral counterpart. Can be positive or negative.

Lewis structures for hypervalent atoms

Hypervalent atoms in lewis structure do not abide by the octet rule as they have more than 8 valence electrons and tend to have 4+ bonds.

Factors for assessing relative stability of lewis structures

Stability can be assessed by looking for the Lewis structures with the least amount or no charge and by assessing whether or not all octets and duets are satisfied. Assessing charge placement (placing negative charges on the most eN atoms and positives on the least eN is most stable).

Resonance structures and their relationships

Resonance structures show different ways of structuring the same molecule. Some resonance structures are more stable than others. They can also show partial bonds in charged molecules.

ranking relative energies of resonance structures

Resonance structures can be ranked in energy by spotting charge placement and the number of charges. The resonance structures with the least amount of charges and placement on atoms that prefer those charges lowers the total energy of the resonance structure making it more stable

VSEPR (2.3.4), electronic geometry, bond angles

Can be found by counting the number of electron elements surrounding an atom (bonds and lone pairs)

• VSEPR 2, 180 degrees, linear geometry

• VSEPR 3, 120, trigonal planar

• VSEPR 4, 109.5, tetrahedral (can be smaller if there are more lone pairs)

Bond dipoles and molecular polarity

If bond dipoles cancel out and the molecule is symmetrical (typically only seen when the central atom is surrounded by the same atom on each side), then the molecule can have polar bonds but be nonpolar overall.

Impact of resonance structures on molecular polarity

With charged structures, the resonance structures show that the charge is distributed between all atoms so as long as the central atom is surrounded by the same atom on all sides, the molecule is nonpolar

Hybridization theory

Demonstrates the mixing of orbitals based on the VSEPR #

S and p orbitals mix to make hybrid orbitals that form sigma bonds, while the unhybridized p orbitals form pi bonds. The s and p orbitals combine if they are in the same phase

Hybrid orbitals are somewhere in between the energy of s and p orbitals

hyrbid orbitals vs unhybridized p’s in bonding

VSEPR 2 —> 2 sp obritals and 2 unhybrid p orbs.

VSEPR 3 —> 3 sp² orbitals and one unhybrid p orb.

VSEPR 4 —> 4 sp³ orbitals and no unhybrid p orbs.

Molecular orbital theory for homonuclears

Two s orbitals in the same phase combine to form a bonding orbital, while two in opposite phases do not overlap and form an antibonding orbital. Show magnetism and provide bond order for the homonuclears

Bonding and antibonding orbitals

Bonding orbitals are more stable and therefore at lower energies

Sigma and pi overlap and impact on MO diagram for period 2

Sigma overlap tends to be at a lower energy since it is more stable in the MO diagram; however, sigma p bonding is higher energy than pi p bonding for B, C, and N.

One set of p orbitals can have sigma bonding whereas the other two sets have pi bonding

Identifying correct MO diagram based on diatomic

B, C, and N use the diagram with sigma 2p and pi 2p switched, where pi is lower

O and F use the diagram where sigma 2p is lower than 2p

Bond order based on electrons in bonding & antibonding

By subtracting the total number of electrons in the antibonding MO’s from the number of electrons in the bonding MO’s and dividing by 2, we get the predicted bond order of the homonuclears

paramagnetism and diamagnetism

Paramagnetism is present when there is at least one unpaired electron

• Show magnetic properties

Diamagnetism is present when there are no unpaired electrons

• show no magnetic properties

Frontier MO’s: highest occupied and lowest unoccupied

The highest occupied MO: the most reactive electrons that would move around

Lowest unoccupied MO: the most stable place to receive new electrons

States of matter

Solid: atoms packed tightly together, a lot of IMFs

Liquid: atoms slightly looser, slightly less IMFs than a solid

Gas: atoms bouncing everywhere and very free, very little IMFs

Phase transitions

Solid to liquid = melting: ΔH of fusion (endothermic)

liquid to gas = vaporization: ΔH of vaporization (endothermic)

gas to liquid = condensation: -ΔH of vaporization (exothermic)

liquid to solid = freezing: -ΔH of fusion (exothermic)

Dispersion forces

Present in every single atom and molecule

Very weak and caused by a temporary weak charge when electrons move to one side of the electron cloud

Increases when molecular mass increases, as there are more electrons

Increases with surface area as there is more space for interaction

Dipole-dipole forces

When polar molecules are attracted to partial charges of other molecules. O in one water molecule is attracted to the H’s of another water molecule.

• Increases with the magnitude of the dipole: the more electronegative the atom, the greater the dipole-dipole forces. Doesn’t always mean a higher boiling point if masses are different

hydrogen bonding

Occurs when a hydrogen is directly bonded to either F, O, or N. It forms a hydrogen bond with other molecules by bonding to a lone pair on an F, O, or N atom of another molecule

Ion-dipole forces

Occurs when an ionic salt is dissolved in a polar liquid (typically water), the polar molecules are attracted to the charged atoms of the ionic compound

Predicting IMFs of a molecule

All molecules have dispersion forces, only polar molecules can have the other 3.

IMF strength increases in this order

Dispersion, Dip-Dip, H bonds, Ion-Dip

Impact of IMFs on BP and VP

The more IMFs, the greater the BP because it takes more heat to break up those IMFs; it also lowers VP, which is why BP is higher, since it also takes more heat to make the VP equal atm pressure.

• Boiling point is equal to the temperature at which vapor pressure equals atmospheric pressure

Relative BP and VP predictions

A molecule with few IMFs has a higher VP and lower BP

A molecule with many IMFs has a lower VP and greater BP

melting vs vaporization

Vaporization requires more energy because you have to break almost all the IMFs to turn a liquid to gas whereas melting only required loosening up IMFs so the atoms move more freely

difference in DeltaH vap and DeltaH fus

DeltaH Vap is the amount of energy it takes to go from liquid to gas or gas to liquid

DeltaH fus is the amount of energy it takes to go from solid to liquid or liquid to solid

Simple and ideal gas laws in calculations

Ideal gas law PV=nRT

P increases when V decreases

V increases when T increases

n increases when V increases

Use gas laws to solve problems

Use PV=nRT if all but one variable is given

use PV/nRT = PV/nRT if trying to calculate the change in a variable before and after

Mole fraction

proportion of the moles of one atom over the total moles present

If there are 2 moles of N2 and 5 moles present, the mole fraction is 2/5

Partial pressure

Partial pressure is calculated by using the moles of one atom to calculate the partial pressure or by multiplying the mole fraction by the total pressure (if given)

Total pressure of gas mixtures

Total pressure can be found by using the total moles present to calculate or by obtaining the sum of all the partial pressures.

Rotation of sigma and pi bonds

Sigma bonds can freely rotate on their axis, whereas pi bonds cannot because rotating would break their side-on bonding

• Because of this, only single bonds can freely rotate