AP Chemistry Lunch Review

Nucleus: protons and neutrons ◦ Isotopes – different number of neutrons

Electrons: ◦ Quantum numbers: four numbers- describe electron ◦ n – principle energy level (distance from nucleus)

◦ l – sublevel shape (values are 0 to n-1)

s (sphere) = 0

p (peanut) = 1

d (double peanut) = 2

f (flower) = 3

◦ m – orbital position (magnetic)

Orbital position in space (values are – l to + l ) ◦ s – spin (+½ or -½)

clockwise or counterclockwise

Electron Configuration –

Rules:

◦ Pauli Exclusion Principle – no two electrons in same atom have same four quantum numbers

◦ Hunds Rule – electrons will not pair until there is one in each degenerate orbital (one in each before doubling)

◦ Aufbau Principle – sublevels fill in order of increasing energy - look at arrangement of periodic table

+1 -3

Many different charges

+2

N O B E L G A S E S

-1 -2 +2 +4 -4

+3

Many Different Charges

Atomic Radius - size

Ionization Energy – energy to remove electron

Electronegativity – pull on pair of bonded electrons

Electron Affinity – pull electron in to neutral atom

Reasons down – increase in energy levels and shielding Reason across – increase in effective nuclear charge (pulls electrons in tighter)

Arrows show increase in indicated direction

True Bonds hold atoms together ◦ Ionic – transfer electrons ◦ Covalent – share electrons ◦ Metallic – delocalized electrons (free to move around)

Intermolecular forces (van der Waals) – weak forces between molecules ◦ Hydrogen bond - H directly bonded to N, O or F

Strongest type of polar force

Responsible for all strange properties of water

◦ Polar – unequal pull on bonded electrons In bond – different electrongativities

In molecule – electrons must be asymmetrical - leads to ion-dipole, dipole-dipole force between polar molecules

◦ London dispersion – weakest but always present Attraction between temporary dipoles

Lewis structures - draw based on valence electrons and formal charge

VSEPR shapes – electrons repel as far away as possible (review handout)

Hybridization of orbitals – merge of s, p and sometimes d orbitals

Sigma bonds – on axis, first bond formed – all single bonds are this type

Pi bonds – sideways overlap of unhybridized p orbitals (all double or triple bonds are this after sigma is formed)

Gases: ◦ Random, straight lines until collision

◦ Diffusion/effusion – Graham’s Law

◦ Ideal Gas Law ignores

volume of individual molecules

attractive force between molecules

Liquids and Solids: ◦ Condensed states

◦ Less energy – movement is mostly vibration around moving or fixed point (liquids slip past)

◦ Phase Diagrams: show P and T where different phases are in equilibrium with each other

Molecular ◦ Held together by dispersion forces

◦ Molecules at lattice positions

Ionic ◦ Held together by electrostatic force of ionic bond

◦ Ions at lattice positions

Metallic ◦ Held together by mobile delocalized electrons

◦ Atoms at lattice positions

Covalent Network ◦ Held together by covalent bonds

◦ Atoms at lattice positions

Look up properties of each type

consider bp/mp, solub., conduct.

Factors affecting solubility: ◦ Identity

interactions between solute and solvent

(nonpolar, polar, ionic forces of attraction)

◦ Pressure

↑ solubility of gases in a liquid

Henry’s Law – gas solubility is proportional to partial pressure of the gas above the liquid

◦ Temperature

Increasing temperature favors the endothermic process to use up the heat