Chapter 10: Solids, Liquids, and Gases

Particles and Phase Change

Solid

• Particles in solids are close together and are held in a fixed place.

• Definite shape and volume.

Liquid

• Particles in liquids are close together and move freely past each other.

• Definite volume, adopts to shape of the container its in.

Gas

• Particles in gases are far apart and have little interaction.

• Adopts the shape and volume of the container its in.

Phase Change

• Phase change - a transition from one state of matter to another.

Influencing Physical Properties

• Stronger forces between particles = higher melting and boiling points.

  • More energy (heat) is needed to break stronger forces and allow a substance to change phases.

Ionic Substances

• Lattices - rigid frameworks of atoms. molecules or ions.

• Most ionic substances have high melting points.

Metallic Substances

• Form lattices of tightly packed atoms.

• Electrons move easily between atoms.

• Conducts electricity

• Shapes of metals are easily altered.

  • Malleable - metals can be pounded into different shapes.

  • Ductile - metals can be stretched into wires.

Molecular Substances

• Forces within molecules: Covalent bonds/molecules

• Forces between molecules: Intermolecular forces

  • Composed of discrete molecules.

  • Forces between individual molecules are much weaker lower melting and boiling points.

Covalent Networks and Polymers

• Covalent networks - lattices of covalent bonds that form giant molecules.

• Have long sequences of covalent bonds resulting in huge molecules.

• Polymers - contain long chains of covalently-bonded atoms.

Intermolecular Forces

• Intermolecular forces - forces between molecules.

• Three key groups:

  • Dipole-dipole interactions (DD)

  • Hydrogen bonds

  • Dispersion forces

Dipole-Dipole Interactions

• Dipole-dipole: Attractions between polar covalent molecules.

  • Molecules of a compound stick together because of the attraction of the positive and negative poles (like magnets).

• Polar

• All polar molecules will have dipole-dipole forces.

• The sides have either a slight positive or slight negative charge.

• Net Dipole: Polar molecule

  • High melting point and boiling point due to stronger DD forces.

  • Takes more heat energy to pull them apart.

• No Dipole: Non-polar molecule

  • Lower melting point and boiling point.

  • Due to weaker forces.

  • No DD.

Hydrogen Bonding

• Hydrogen bonding - strongest of intermolecular forces.

• Polar

• Between molecules containing H-F, H-O, or H-N bonds.

• For hydrogen bonding to occur, the hydrogen must interact with two different electronegative atoms.

London Dispersion

• Dispersion - the weakest of the intermolecular forces.

• Non-polar

• All molecules and atoms have dispersion forces.

• Temporary dipole form when atoms or molecules interact.

• Very short lived, very weak.

Polar and Non-polar Intermolecular Forces

• Polar molecules

  • Always have dipole-dipole and dispersion.

  • Sometimes have H bond.

• Non-polar molecules

  • Dispersion only.

Describing Gases

• Molecules that are spread apart as far as possible.

• Very little interaction between particles, if they do interact they bounce off of each other.

Ideal Gas

• The volume of particles is much less than the volume of the container.

• Particles have NO attraction for each other.

• Bounce off of each other.

Pressure

• Pressure - the force that gases exert on their surrounding.

• Equation:

  • Pressure = Force/Area

• Barometer - Device that measures atmospheric pressure.

• Millimeters of mercury (mm Hg)

  • 1 mm Hg = 1 torr

• Gauge pressure - the different between the pressure of compressed has vs atmospheric pressure.

• 1 atmosphere = 760 mm Hg (torr).

The Gas Laws

Boyle’s Law

• The pressure and volume of a gas are inversely related.

• If pressure (P) goes up, then volume (V) goes down.

• If pressure (P) goes down, then volume (V) goes up.

• Equation:

  • $P_{1}V_{1}=P_{2}V_{2}$

  • P1 = initial pressure

  • V1 = initial volume

  • P2 = final pressure

  • V2 = final volume

Charle’s Law

• At constant pressure, the volume of a gas is directly proportional to its temperature.

• If temperature (T) goes up, volume (V) goes up.

• If temperature (T) goes down, volume (V) goes down.

• V is proportional to T

• Equation:

  • $\dfrac{V_{1}}{T_{1}}=\dfrac{V_{2}}{T_{2}}$

  • V1 = initial volume

  • T1 = initial temperature

  • V2 = final volume

  • T2 = final temperature

The Combined Gas Law

• If multiple variables are changing, use the combined gas law.

• Equation:

  • $\dfrac{P_{1}V_{1}}{T_{1}}=\dfrac{P_{2}V_{2}}{T_{2}}$

  • 1 = initial pressure, volume, temperature

  • 2 = final pressure, volume, temperature

The Kelvin Scale

• Absolute zero = -273.15 degrees Celsius or 0 K

• kelvin = C + 273.15

• Gas laws must always be in kelvin.

Avogadro’s Law

• If temperature and pressure are constant, the volume of a gas is proportional to the number of moles of gas present.

• The more gas is present the larger the volume it occupies.

The Ideal Gas Law

• Relates the amount of gas to its pressure, volume, and temperature.

• Equation:

  • $PV=nRT$

  • P = pressure

  • V = volume

  • n = # of moles of gas

  • R = the gas constant

    • R = 0.0821 L.atm/mol.K

  • T = temperature

• T must be in kelvin

Mixtures of Gases: Partial Pressure

• Partial pressure - the pressure caused by one gas in a mixture.

• Adding up all partial pressures = total pressure

• Adding up moles of each gas = total moles of gas

Diffusion

• Diffusion - the spread of particles through random motion.

• Lighter particles diffuse more quickly.

• Example: a candle

Effusion

• Effusion - the process of a gas escaping from a container.

• Depends on the velocity of the gas particles.

• Lighter particles are faster.

• Heavier particles are slower.

• Example: Helium escaping from a balloon.

  • Light gas/atoms so it deflates fast.