Chem 1B Midterm 1

Energy

The capacity to do work.

System

The part of the universe being studied; the body undergoing change.

Surroundings

Everything outside the system.

Potential energy

Energy due to position or chemical composition.

Kinetic energy

Energy of motion.

Mechanical energy

The sum of kinetic and potential energies.

Internal energy (U)

Total energy stored in a system; a state function.

State function

A property that depends only on the current state, not on the path taken.

Thermal equilibrium

When no net heat is exchanged between systems in contact.

Heat (q)

Thermal energy transferred between system and surroundings.

Enthalpy (H)

Heat content of a system at constant pressure; q ≈ ΔH.

Temperature

Measure of the average kinetic energy of particles; drives heat flow.

q = mcΔT

Equation relating heat transferred to change in temperature (m is mass, c is specific heat).

Work (W)

A force acting over a distance; energy transfer through work.

W = PΔV

Work done during expansion against external pressure.

Constant pressure

Pressure remains the same during the process.

Constant volume

Volume remains the same during the process.

q and H are equivalent at constant pressure

At constant pressure, heat exchanged equals enthalpy change (q ≈ ΔH).

q and ΔU are equivalent at constant pressure with little work

When PV-work is small, q ≈ ΔU.

q and ΔU are equivalent at constant volume

At constant volume, heat added equals the change in internal energy (q = ΔU).

Bomb calorimeter

Calorimeter used to measure heat transfer at constant volume.

Coffee cup calorimeter

Calorimeter used to measure heat transfer at constant pressure.

Bond energy

Energy required to break one mole of bonds in a compound.

Enthalpy of formation

ΔHf: Enthalpy change when one mole of a substance forms from its elements in standard state.

Standard state

Standard 1 atm pressure and 25°C; reference state for enthalpies.

Enthalpy of formation for elements

Defined as zero in their standard states.

Hess’s Law

The overall enthalpy change for a reaction equals the sum of enthalpy changes of steps.

ΔHtotal (sum of steps)

The total enthalpy change is the sum of stepwise enthalpy changes.

Multiplier -1

If you reverse a reaction, multiply its ΔH by -1.

Multiply by n

If you multiply a reaction by n, multiply its ΔH by n.

Lattice energy

Energy required to convert a solid ionic lattice into gaseous ions; depends on charge and radius.

Higher charge → larger lattice energy

Ions with higher charges create stronger lattice energy.

Larger radius → smaller lattice energy

Lattice energy decreases as ionic radii increase.

Born-Haber cycle

A graphical representation of the steps forming an ionic compound.

Gas

A state of matter with no fixed shape or volume.

Velocity

Speed and direction of gas molecules; at higher temperatures, speed increases.

Pressure

Force per unit area; increases with more collisions with container walls.

Units of gas pressure

atm, mmHg (torr), Pa, kPa, psi.

Volume of gas molecules

In an ideal gas, molecules occupy negligible volume.

Elastic collisions

Collisions where kinetic energy is conserved.

Boyle’s Law

Relates pressure and volume at constant temperature: P ∝ 1/V.

Charles’s Law

Relates volume and temperature at constant pressure: V ∝ T.

Avogadro’s Law

Relates volume and number of moles at constant T and P: V ∝ n.

Ideal gas law

PV = nRT; relates pressure, volume, temperature, and amount.

Mole, volume, temperature interconversion

Unit conversions using the ideal gas law connect moles, volume, and temperature.

Partial pressure

P_i: pressure contribution of a component in a gas mixture.

Mole fraction

X_i: moles of component i divided by total moles.

Pi = Xi P_total

Partial pressure equals mole fraction times total pressure.

Kinetic energy

Energy of motion; (1/2)mv^2 for a molecule.

Kinetic Molecular Theory (KMT)

Model assuming point masses, random motion, and elastic collisions.

Point masses

KMT assumption that particles have negligible volume.

Elastic collisions

Collisions where kinetic energy is conserved.

Temperature and kinetic energy

Rises with increasing temperature; higher temperature means higher average kinetic energy.

Distribution of velocities

Gases have a distribution of molecular speeds, not a single value.

Diffusion

Spontaneous mixing of gases due to random motion.

Effusion

Gas moves through a small opening into a vacuum.

Rate ∝ 1/√M

Rate of diffusion is inversely proportional to the square root of molar mass.

High temperature, low pressure

Gases behave more ideally under these conditions.

Ideal gas at atmospheric pressure

True: under typical conditions many gases approximate ideal behavior.

Van der Waals equation

Equation of state that corrects PV for non-ideal gases by including pressure and volume terms.

Van der Waals constants (a and b)

Constants in the equation of state accounting for intermolecular forces (a) and finite volume (b).