Chem 1332 Exam 1 Review

University of Houston Fundamentals of Chemistry

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Chem 1332 Exam 1 Review by Mind Map: Chem 1332   Exam 1 Review

1. 1331 Review

1.1. Chapter 2

1.1.1. Naming ionic & covalent compounds

1.1.2. Polyatomic ions

1.1.3. Acids

1.1.4. Organic compounds

1.2. Chapter 4

1.2.1. Solubility rules

1.3. Chapter 9 & 10

1.3.1. Lewis Structures

1.3.2. Molecular shape

1.3.3. Polarity of molecules

2. Chapter 12: Solids & Liquids

2.1. Interparticle forces

2.1.1. Infinite arrays

2.1.1.1. Ionic

2.1.1.2. Metallic "sea of electrons"

2.1.1.3. Covalent

2.1.2. Molecular interactions

2.1.2.1. Hydrogen bonding H bonded to F, O, or N

2.1.2.2. Dipole-dipole

2.1.2.3. London dispersion (induced dipoles, incr. w/ mass)

2.2. Stronger IPFs

2.2.1. Incr. boiling point

2.2.1.1. (Decr. vapor pressure)

2.2.2. Decr. melting/freezing point

2.2.3. Incr. ΔH

2.3. Phase Changes

2.3.1. Phase diagram

2.3.1.1. Critical point, triple point

2.3.1.2. Line between (s) & (l) almost vertical: m.p. indep. of p; negative slope for water

2.3.1.3. At boiling point: vapor pressure = atmospheric pressure

2.3.2. ΔH(fus)

2.3.2.1. Endothermic (+)

2.3.2.1.1. Solid --> Liquid (melting)

2.3.2.2. Exothermic (-)

2.3.2.2.1. Liquid --> Solid (freezing)

2.3.3. ΔH(vap)

2.3.3.1. Endothermic (+)

2.3.3.1.1. Liquid --> Gas (vaporization)

2.3.3.2. Exothermic (-)

2.3.3.2.1. Gas --> Liquid (condensation)

2.3.3.3. ΔH°(vap) > ΔH°(fus)

2.3.4. ΔH(subl)

2.3.4.1. Endothermic (+)

2.3.4.1.1. Solid --> Gas (sublimation)

2.3.4.2. Exothermic (-)

2.3.4.2.1. Gas --> Solid (deposition)

2.4. Heating curves

2.4.1. Graph T vs q

2.4.1.1. Heating -->

2.4.1.2. Cooling <--

2.4.2. Heat material: q = mcΔT

2.4.3. Phase change: q = ΔH° x n

3. Chapter 13: Solutions

3.1. Concentration

3.1.1. mass % solute = (g solute / g solution) x 100%

3.1.2. volume %

3.1.3. mole fraction, X solute = mol solute / total mol

3.1.4. Molarity, M = mol solute / L solution

3.1.5. molality, m = mol solute / kg solvent

3.2. Colligative Properties

3.2.1. Vapor pressure

3.2.1.1. solute - non-volatile, non-electrolyte

3.2.1.1.1. v.p. solution = (v.p solvent) (X solvent)

3.2.1.2. Solute particles get in the way

3.2.1.3. volatile solvent & volatile solute

3.2.1.3.1. v.p. sol'n = (X solv.) (v.p. solv.) + (X solute) (v.p. solute)

3.2.2. Boiling point

3.2.2.1. b.p. solution > b.p. solvent

3.2.2.2. volatile solvent, non-volatile solute

3.2.2.2.1. Δb.p. = k(bp) x m

3.2.2.2.2. k(b.p.): boiling point constant

3.2.3. Freezing point

3.2.3.1. f.p. solution < f.p. solvent

3.2.3.2. Δf.p. = k(f.p.) x m

3.2.4. Osmotic pressure

3.2.4.1. π = MRT

3.2.4.2. R = 0.0821 L*atm/mol*K

3.2.5. If ionic compound is soluble in water, it dissociates and incr. the # of moles of solute particles in the sol'n. (Ex: 0.1 m NaCl (aq) --> 0.1 m Na+ & 0.1 m Cl-)

3.3. Solubility

3.3.1. "Like dissolves like": IPFs

3.3.2. ΔH(sol'n) = (IMF solv. & solute) - IMF sol'n endo (break) exo (make)

3.3.3. V. positive (endo) ΔH(sol'n) --> insoluble

3.3.4. Small or negative (exo) ΔH(sol'n) --> soluble

3.3.5. Ion-dipole depends on strength of ionic bond

3.3.6. ΔS: change in entropy (disorder)

3.3.6.1. (+) ΔS(sol'n): incr. in disorder, "nice"

3.3.6.2. Gas in a solid/liquid: ΔS(sol'n) is (-)

3.3.7. Effect of T on solubility

3.3.7.1. ΔH(sol'n) endothermic (+): solubility incr. as T incr.

3.3.7.2. ΔH(sol'n) exothermic (-): solubility decr. as T incr.

3.3.7.3. For (g) only: Henry's Law

3.3.7.3.1. solubility = K(H) x P

3.3.7.3.2. K(H), Henry's constant

4. Chapter 16: Kinetics

4.1. Rates of Reaction

4.1.1. Rate of reaction = Δ[x] / Δt (In mol/L*s)

4.1.2. Rate Law

4.1.2.1. rate = k[reactant 1]^x [reactant 2]^y

4.1.2.1.1. x,y: orders, depend on rxn (calculate experimentally)

4.1.2.1.2. Units of k:

4.1.2.2. First order: ln[A(0)] - ln[A] = kt

4.1.2.3. Second order: (1/[A]) - [1/[A(0)] = kt

4.1.2.4. Half-Life, t(1/2)

4.1.2.5. Determine rate law & k from experimental data