Chem Notes Ch 7 Flashcards
Properties of solids
Crystalline or amorphous. Molecules held by rigid intramolecular bonds.
Ionic solids
Held together by ionic bonds. Properties: hard, non-conductive, brittle, and high melting point
Metallic solids
Held together by a sea of free electrons flowing around a lattice of
metal cations. Properties: malleable, ductile, conductive, high luster (shiny!), and variable melting points and hardness.
Covalent network solids
Held together by network of covalent bonds. Properties: hard, non-conductive, and high melting point.
Molecular solids
Held by intermolecular forces (hydrogen bonding, dipole-dipole forces, dispersion forces). Properties: soft, non-conductive, and low melting point.
Crystalline solids
Form an orderly array of molecules that follow a repeating pattern. The structure of the crystal pattern is defined by the smallest individual repeating unit cell.
Simple cubic
1 atom per unit cell; one atom at every corner of cube
Body-centered cubic
2 atoms per unit cell; one atom at center of cube and one atom at every corner of cube.
Face-centered cubic
4 atoms per unit cell; one atom at the center of all the faces of the cube and at every corner of the cube.
Properties of liquids
Free (and constantly) moving but held close together by intermolecular forces like van der Waals, dipole-dipole, and hydrogen bonding.
Van der Waals
Combination of weak intermolecular forces between all types of molecules. Caused by random shifts in electron density of particles that create temporary weak poles. Low Bond Strength. Occurs on all molecules. Increases with particle size
Dipole-Dipole
Occurs between molecules with permanent uneven distribution of elections. Caused by electronegative atoms that draw in neighbouring electrons creating poles of high and low charge. Medium Bond Strength. Occurs only on polar molecules. Increasing number of polar groups increases boiling point. Symmetrical dipoles cancel out
Hydrogen Bonding
Special type of dipole-dipole force; occurs when highly electronegative atoms like F, N, or O strip neighbouring hydrogen’s electron. High Bond Strength. Occurs between an H atom and an F, N, or O atom. Electronegative group and hydrogen can have multiple bonds
Heterogenous mixture
2 phases are separated and easily distinguishable
Homogenous mixture
The phases appear as one continuous phase
Solubility
Degree to which a solid (solute) can be dissolved in a liquid (solvent)
Miscibility
Degree to which a liquid mixes with another liquid
Stronger intermolecular forces lead to
Higher boiling point, higher heat of vaporization, higher viscosity, higher surface tension, and lower vapor pressure
Phase diagrams
Has lines on the phase diagram to indicate the boundaries where phase changes occur.
At the triple point, gas, liquid, solid phases are all in equilibrium. The critical point exists at very high temperatures and pressures and is defined as the point in a phase diagram at which the liquid and gas phases are indistinguishable.
The normal melting point is the melting point at 1 atm, and the normal boiling point is the boiling point at 1 atm. The point at which the dotted line intersects the solid-liquid equilibrium line is the normal melting point, and the point at which the dotted line intersects the liquid-gas equilibrium line is the normal boiling point.
Phase Equilibrium:
Gain in energy drives the solid → liquid → gas forward and loss of energy drives it back. Reaction may reach equilibrium where forward and reverse reactions are balanced. Temperature of a substance is the average kinetic energy of it’s particles, and some particles will have higher energy and others lower. Some particles change phase while the bulk remains in their current state. Substances are in phase equilibrium where some fraction of their particles are changing phase back and forth.
Solid-Liquid Equilibrium
Melting or fusion: Particles in the solid ice absorb enough energy, they vibrate strongly enough that they break free of bonds and can flow into the liquid water phase
Solidification, freezing, or crystallization: Some particles in liquid water lose enough energy that they slow down and form bonds with the ice, becoming solid
Freezing point depression: A solid having a lower freezing point due to interaction of another particle, marking it harder to form crystals. Δ(Tfreezing) = (i)(Kfreezing (°C·kg/mol))(M (mol/kg))
ΔTfreezing (°C) is drop in freezing point temperature, Kfreezing (°C·kg/mol) is cryogenic constant, a property of the solvent, M (mol/kg) is the moles of solute/mass of solvent (mol/kg) and i (unitless) is the van’t Hoff factor.
Liquid-Gas Equilibrium
Vaporization: When particles of liquid have enough energy to escape the weak bonds and become gas
Condensation: When pressure on vapour turns the gas into liquid
Vapour pressure: Pressure exerted on the lid by the buildup of gas and liquid in a container, increases with temperature
Boiling point: When vapour pressure is equal to the pressure on the surface of liquid
Raoult’s Law for 2 substances with different vapour pressures: Ptotal = (mole fraction of A)(vapour pressure of pure A) + (mole fraction of B)(vapour pressure of pure B)
Boiling point elevation. ∆(𝑇boiling) = (𝑖)(𝐾)(𝑀)
Solid-Gas Equilibrium
Sublimation: Solid transitioning to vapour, without water
Deposition: Vapour transitioning to solid, without water
Energy of phase equilibrium
Driven by thermodynamics. Phase change can be shifted with a change in composition, energy or pressure.
Sensible heat
The input of energy needed to increase temperature of a substance. Q = (m)(c)(deltaT). Q(J) is energy gain/loss, m (kg) is mass, c (J/kg·°C) is the specific heat capacity.
Latent heat or heat of transformation
The breaking or forming of intermolecular bonds.
Qlatent = (m)*(delta H vaporization)
Qlatent = (m)*(delta H fusion)
ΔHcondensation = -ΔHvaporization
ΔHfusion = -ΔHsolidification
Sensible heat vs latent heat
The energy in sensible heat goes into increasing or decreasing the kinetic energy of the particles but the energy of latent heat goes into breaking or forming intermolecular bonds. In this way the latent heat does not change the temperature of substance but instead causes the phase change to occur.
The latent heat of a phase change can be calculated with:
- Qlatent = m∆Hvaporization
- Qlatent = m∆Hfusion
Where Qlatent (J) is the energy gain/loss, m (kg) is the substance mass, ΔHvaporization/fusion (J/kg) is the latent heat of transformation. Usually, reference materials list the ΔHvaporization and the ΔHfusion because ΔHcondensation = -ΔHvaporization and ΔHfusion = - ΔHsolidification.
