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Liquids and Solids

Bonds: ionic, covalent, or due to intermolecular forces

All intermolecular forces are electrostatic:

ion-dipole forces (between ions and polar molecules), dipole-dipole forces (between polar molecules), London dispersion forces (between any kind of molecules), hydrogen-bonding forces (involving molecules in which hydrogen is bonded to N, O, or F atoms).

London dispersion forces the stronger the larger the molecular weight. London forces stronger for linear shapes than for other shapes.

Melting point and boiling point depend on intermolecular forces

The stronger the intermolecular forces, the higher the melting and boiling point

normal boiling point = the temperature at which a liquid boils at 1 atm pressure

viscosity and surface tension are dependent on intermolecular forces and the structures of liquids.

Vapor pressure

In P = constant - (ΔHvap / RT)

ΔHvap = enthalpy of vaporization

Crystalline solids (with well-defined surfaces; includes metallic elements) or amorphous solids.

The arrangement of atoms, ions, and molecules in a crystal can be determined by x-ray diffraction.

Crystalline solids divided into:

metallic (s- and d-block elements)

ionic - hard, rigid, brittle, high melting and boiling points

network (includes graphite and diamonds) - hard, rigid, brittle, very high melting points

molecular (includes ice) - relatively low melting and boiling points

phase = the physical form of a substance, which may be gas, liquid, or a variety of solid states

Unit cells in crystalline solids: primitive cubic, body-centered cubic, face-centered cubic

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Coordination number: the number of equidistant neighbors. The coordination number is 6 in primitive cubic units (occupied space: 52 %), 8 in body-centered cubic units (occupied space: 68 %), and 12 in face-centered cubic units (occupied space: 74 %)

2 types of close-packing for face-centered units: hexagonal close packing (layers ABABABAB); cubic close packing (layers ABCABCABC).

Types of solids:

Molecular solids - forces: London, dipole, hydrogen - example: sucrose, frozen CO2

Covalent-network solids - forces: covalent bonds - example: diamond, quartz (SiO2)

Ionic solids - forces: electrostatic attractions - examples: salts such as NaCl

Metallic solids - forces: metallic bonds - examples: all metallic elements


3 steps in forming a solution:

1. separating particles of a solute (requires energy: ΔH1>0)

2. separating particles of a solvent (requires energy: ΔH2>0)

3. combining particles of solvent with particles of solute (frees energy: ΔH3<0)

The sum of the three processes can be positive (higher enthalpy, endothermic) or negative (lower enthalpy, exothermic)

Processes in which the energy content of a system decreases (exothermic processes) and/or in which the disorder (entropy) increases tend to occur spontaneously.

mass percentage = parts per 100 (units: gram/100gram = g/100g)

ppm = parts per million (units: gram/1,000,000gram = milligram/kilogram = mg/kg)

ppb = parts per billion (units: gram/1,000,000,000gram = microgram/kilogram = µg/kg)

mole fraction (x) = (moles of one component)/(total moles of all components of a solution)

mole fraction (x) = the mole percentage of a substance in a mixture

molarity = (moles of solute)/(liters of solution)

molarity (small cap M) = mol / L

molality = (moles of solute)/(kilogram of solvent)

molality (m) = mol / kg

Enthalpy of solution = heat per mole produced when a substance dissolves

Lattice enthalpy = heat required to turn a solid ionic compound into an ionic gas

Enthalpy of solvation (for water, this is called enthalpy of hydration) = heat released when the solution is formed

Enthalpy of solution = lattice enthalpy + enthalpy of solvation (the sum may be positive or negative; if positive, there is a net inflow of energy and the process is endothermic; if negative, there is a net outflow, and the process is exothermic).

The solubility of a gas is proportional to its (partial) pressure.

For gases, an increase in temperature normally lowers solubility.

Enthalpy of solution: heat released or absorbed when a substance dissolves.

Enthalpy of solution is the sum of the lattice enthalpy and the enthalpy of solvation.

If the solvent is water, salvation is called hydration.

The lattice enthalpy corresponds to the energy needed (always endothermic) to break apart an ionic solid to form two widely separated ionic gases: NaCl (s) → Na+ (g) + Cl- (g)

The enthalpy of hydration corresponds to the energy released (always exothermic) to dissolve ionic gases in water: Na+ (g) + Cl- (g) → Na+ (aq) + Cl- (aq)

The sum of the two processes can either be negative or positive.

Henry law constant

Solubility of a gas (molarity; small cap M; mol per liter) = Henry’s law constant (mol/[liter x atm]) x partial pressure of a gas over a solution (atm)

Molar solubility = constant x partial pressure

S = kH x P

Units: mol/L = mol/(L x atm) x atm

kH depends on the gas, the temperature, and the pressure. At constant temperature, the solubility of a gas is proportional to its partial pressure.

Pressure of a gas mixture - partial pressure of gas A + partial pressure of gas B + ACA

The quantity of each gas in a mixture is determined as mole fractions.

Colligative properties:

(lowering of) vapor pressure

(raising of) boiling point

(lowering of) freezing point

tendency of a solvent to pass through a membrane

These properties depend on the relative number of solute molecules, and NOT on their chemical identity.

Mole fraction = (moles of solute) / (moles of solute + moles of solvent)

Molality = (moles of solute) / (mass of solvent)

Achtung: Molality is not per (mass of solution), only per (mass of solvent)

Raoult law Vapor pressure of a solution = mole fraction of solvent x vapor pressure of pure solvent

P = xsolvent x Ppure

The vapor pressure of a solvent in the presence of a nonvolatile solute is proportional to the mole fraction of the solvent.

Freezing-point depression = ΔT

ΔT = ikf x molality

kf = freezing-point constant

Unit of freezing-point constant: (K x kg) / mol

i = van Hoff factor: 2 for MX salts (the kind of NaCl), 3 for MX2 or M2X salts (the kind of CaCl2), and 1 for nonelectrolytes

Osmotic pressure = П (uppercase pi)

П = iRT x molarity

Units: [(L x atm) / (K x mol)] x K x (mol / L) = (L x atm x K x mol) / (K x mol x L) = atm

i = van Hoff factor: 2 for MX salts (the kind of NaCl), 3 for MX2 or M2X salts (the kind of CaCl2), and 1 for nonelectrolytes

R = gas constant

T = absolute temperature


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