Energetics Things to Know

Thermo Chemistry

Study of energy changes due to chemical reactions

Temperature

Average kinetic energy of particles

Heat

Thermal energy transfer between two bodies at different temperatures

System

Area of interest

Surroundings

Everything else in the universe

Universe

System plus surroundings

Open System

Mass and energy can be exchanged with surroundings (usually in the form of heat)

Closed System

Can exchange energy with surroundings but not mass (example: can of pop)

Isolated System

Mass and energy cannot be exchanged with surroundings (example: cooler chest)

Enthalpy (H)

Total energy of a system; some is stored as potential energy in chemical bonds

Change in Enthalpy (ΔH)

The difference between the enthalpy of products and reactants (Hproducts − Hreactants)

Exothermic Reaction

Reaction where heat from the system is released to the surroundings; ΔH is negative; bonds are made

Endothermic Reaction

Reaction where heat from surroundings is absorbed by the system; ΔH is positive; bonds are broken

Specific Heat Capacity (c)

The amount of energy required to raise the temperature of one gram of a substance by one degree Celsius (or kelvin); for water c = 4.18 J/g°C

Heat Capacity

Amount of energy needed to raise the temperature of a particular substance by one degree Celsius (or kelvin); calculated as m × c

Heat (q) Formula

q = m × c × ΔT or q = C × ΔT

ΔT

Change in temperature, calculated as Tfinal − Tinitial

Heat of Fusion (ΔHfusion)

Phase change from solid to liquid; 334 J/g

Heat of Vaporization (ΔHvaporization)

Phase change from liquid to gas; 2260 J/g

Heat of Condensation (ΔHcondensation)

Phase change from gas to liquid; −2260 J/g

Heat of Solidification (ΔHsolidification)

Phase change from liquid to solid; −334 J/g

Heat of Sublimation (ΔHsublimation)

Phase change from solid directly to gas

Specific Heat of Ice (Cice)

2.09 J/g°C

Specific Heat of Water (Cwater)

4.18 J/g°C

Specific Heat of Steam (Csteam)

2.00 J/g°C

Standard Enthalpy Change of Formation (ΔHf°)

Enthalpy change when one mole of a compound is formed from its elements in their standard states at 298 K and 1 atm pressure

ΔHf° for an Element

Enthalpy of formation for an element in its standard (most stable) state equals 0 kJ/mol

Hess’ Law

Total ΔH for converting reactants to products is always constant, no matter how the change occurs

Bond Enthalpies

Amount of energy required to break a particular bond in one mole of gaseous molecules

Entropy (S)

Measure of randomness in a system

Spontaneous Process

A reaction that occurs under a given set of conditions without external assistance

Gibbs Free Energy (G)

Relates energy obtained from a chemical reaction to ΔH, ΔS, and temperature

ΔG < 0

Reaction is spontaneous (products favored)

ΔG > 0

Reaction is nonspontaneous (reactants favored)

ΔG = 0

Reaction is at equilibrium

Gibbs Chart: +ΔH and +ΔS

ΔG depends on temperature; spontaneity calculation needed

Gibbs Chart: +ΔH and −ΔS

ΔG is positive; reaction is not spontaneous

Gibbs Chart: −ΔH and +ΔS

ΔG is negative; reaction is spontaneous

Gibbs Chart: −ΔH and −ΔS

ΔG depends on temperature; spontaneity calculation needed

Lattice Enthalpy (U or ΔHlattice)

Energy required to completely separate one mole of a solid ionic compound into gaseous ions

Ionization Energy

Amount of energy required to remove an electron from an atom; smaller radius means higher ionization energy

Electron Affinity

Amount of energy released when an atom gains an electron; smaller radius means higher electron affinity