CHMY 121N - Chapter 8 and 9

change of state

The change of a substance from one state of matter (gas, liquid, or solid) to another

melting point

The temperature at which the liquid phase is in equilibrium with the solid phase

boiling point

The temperature at which the gas phase is in equilibrium with the liquid phase

intermolecular forces (IMFs)

• forces that act between molecules or discrete atoms and hold them close to one another

• aka van der Waals forces

intramolecular forces

Forces within atoms of the same molecule

london dispersion force

• temporary attractive forces between molecules created by electron movement

• weak strength

• get stronger as the size of the molecule increases

• aka induced dipole force

do london dispersion forces have a short-lived or long-lived polarity?

• short lived

• at any given instant there may be more electrons at one end of a molecule than at the other

do all molecules experience london dispersion force, or only some?

all molecules

The larger the molecular weight and surface area, the greater the _______ of a molecule

temporary polarization

dipole-dipole forces

• occurs between the positive end of one polar molecule and the negative end of another polar molecule

• gets stronger for more polar molecules

• weak strength

do molecules that contain polar covalent bonds have a net polarity?

yes, they could.

Dipole-Dipole forces are stronger the more _____ the molecule.

polar

ion dipole force

Interaction between a fully charged ion and a polar molecule

hydrogen bond

• occurs between molecules with O-H, N-H, and/or F-H bonds

• gets stronger for more polar molecules

• moderate strength

hydrogen bond donor

H atom with a covalent bond with N, O, or F (most electronegative elements) can provide a hydrogen atom

hydrogen bond acceptor

Another molecule with a lone pair of electrons on N, O, or F (most electronegative elements) accepts hydrogen atoms

can non-polar molecules form hydrogen bonds?

no

In liquids and solids, the stronger the intermolecular forces, the ______ the melting and boiling points.

higher

what are the 3 intermolecular forces from strongest to weakest?

Hydrogen bonds > dipole-dipole > london dispersion

ideal gas

A gas that obeys all the assumptions of the kinetic-molecular theory

non-ideal gas

don’t completely follow these assumptions because of intermolecular forces

Kinetic-Molecular Theory of Gases

1. Gas particles move randomly and don’t attract or stick to each other.

2. The particles are tiny compared to the space between them.

3. Hotter gases (higher temperature) have faster-moving particles.

4. When gas particles collide with each other or the container, they bounce off without losing energy.

Pressure (P)

the force per unit area pushing against a surface

what is the equation for pressure?

pressure = force / area

what is the SI unit for pressure?

pascal (Pa)

how do you get kPa?

multiply pascal by 1000

what are common units of pressure?

torr or millimeters of mercury (mmHg)

atmospheric pressure

measures the amount of pressure that gases in our atmosphere puts on one square inch

what makes up 1 atm?

14.7 lbs/in² (psi)

what is the conversion factor for atm?

1 atm = 760 mmHg = 760 torr

gas laws

A series of laws that predict the influence of pressure (P), volume (V), and temperature (T) on any gas or mixture of gases

directly proportional

if when one variable increases, so does the other variable.

inversely proportional

if when one variable increases, the other decreases.

are area and force directly or inversely proportional?

directly

are volume and pressure directly or inversely proportional?

inversely

Boyle’s Law

P₁V₁ = P₂V₂

What is the proportionality between temperature and volume?

direct

Charles’ Law

V₁ / T₁ = V₂ / T₂

is temperature in kelvin or celcius?

kelvin (°C + 273)

What is the proportionality between temperature and volume?

direct

Gay-Lussac’s Law

P₁ / T₁ = P₂ / T₂

combined gas law

P₁V₁ / T₁ = P₂V₂ / T₂

Avogadro’s Law

V₁ / n₁ = V₂ / n₂

what does n represent?

moles

Standard Temperature and Pressure (STP)

0 °C (273 K) and 1 atm (760 mmHg)

Standard Molar Volume

22.4 L per 1 mol of any gas

Ideal Gas Law

PV = nRT

when should you use the combined gas law?

when the properties of a gas are changing, but the amount of a gas stays the same.

when should you use the idea gas law?

when we know 3 of 4 gas variables (pressure, volume, temperature, amount) and the gas isn’t changing

The higher the molar mass, the _____ the boiling points (due to London dispersion forces).

higher

partial pressure

The contribution of a given gas in a mixture to the total pressure mixtures of gases

does each particle in a gas act independently?

yes, so the chemical identity of its neighbors is irrelevant.

do mixture of gases behave the same as pure gases?

yes, and they obey the same laws

what does the pressure exerted by each gas depend on?

the frequency of collisions of its molecules with the walls of the container.

Dalton’s Law

P_total = P_{gas 1} + P_{gas 2} + P_{gas 3}        

atmospheric pressure

the sum of the partial pressures of all the gases in it

Viscosity

The measure of a liquid’s resistance to flowing (moving).

when does viscosity increase?

when intermolecular forces increase

surface tension

the energy, or work, required to increase the surface area of a liquid due to intermolecular forces.

what is surface tension caused by?

by the stronger inward pull on surface molecules compared to those inside the liquid.

crystalline solids

solids in which the atoms, molecules, or ions are rigidly held in an ordered arrangement.

ionic solids

• particles are ions

• composed of alternating positive and negative ions in 3D arrangement

• held together by ionic bonds

• ex: Na(s) + ½ ​Cl₂(g) → NaCl(s)

properties of ionic solids

• brittle and hard

• high melting point

molecular solids

• made of molecules

• held together by intermolecular forces

• ex: ice, wax, dry ice

properties of molecular solids

• soft

• low to moderate melting points

covalent networks

• individual atoms held together by covalent bonds in giant 3D arrays

• one huge molecule

• ex: a diamond

properties of covalent networks

• very hard

• very high melting point

metallic solids

• individual metal atoms held together by metallic bonds

• valence electrons are delocalized

  • forming a “sea” of electrons that can move freely through the structure

  • this is what makes metals conductive

properties of metallic solids

• lustrous

• can be soft (Na) or hard (Ti)

• high melting point

amorphous solids

• particles are randomly arranged

• no long-range structure

• ex: glasses, tar, some plastics

properties of amorphous solids

• noncrystalline

• no sharp melting point

• able to flow, though may be very slow

• curved edges when shattered

vapor pressure

the partial pressure of vapor molecules in equilibrium with a liquid

What happens to liquid molecules in a closed container?

Some evaporate, but random motion brings some back into the liquid.

What is dynamic equilibrium in a closed container?

Evaporation = condensation → vapor concentration stays constant.

What affects vapor pressure?

Temperature and the liquid’s intermolecular forces.

How do intermolecular forces affect vapor pressure?

Stronger forces → fewer molecules escape → lower vapor pressure.

How does temperature affect vapor pressure?

Higher temperature → more energy → higher vapor pressure.

normal boiling point

the temperature where boiling occurs, at a pressure of exactly 1 atm

Heat of Fusion (H_fus)

The quantity of heat required to completely melt one gram of a substance once it has reached its melting point

Heat of Vaporization (H_vap)

The quantity of heat needed to completely vaporize a liquid at its boiling point

equation for melting point

heat (cal or J) = mass (g) x heat of fusion (cal or J/g)

• grams cancel out

equation for boiling point

heat (cal or J) = mass (g) x heat of vaporization (cal or J/g)

• grams cancel out

equation for heat (with no phase change)

Q = m × ΔT × C

• q = heat (cal)

• m = mass (g)

• ΔT = change in temp (in Celsius)

• c = specific heat (cal / (g*celcius))