Energy
Energy is the ability to do work.
Forms of energy:
Kinetic, potential, heat, light, nuclear, solar, etc.
Law of Conservation of Energy:
Energy cannot be created or destroyed, only changed from one form to another.
Types of Energy in Matter
Potential Energy: Energy due to position or stored chemical bonds
Kinetic Energy: Energy of motion of particles
Internal Energy (U): Total kinetic + potential energy of a system
Heat and Temperature
Q = mc△T
Where:
Q = heat absorbed
m = mass
c = specific heat capacity
ΔT = temperature change
Enthalpy (H)
Total heat content of a substance.
H{reaction} = H{products} - H{reactant}
Exothermic Reactions
Give out heat (ΔH is negative)
Examples:
Combustion
Mg + O2 → MgONeutralization
HCl + NaOH → NaCl + H2ODissolving NaOH in water
Endothermic Reactions
Absorb heat (ΔH is positive)
Example:
CaCO3→CaO + CO2
Chemical Bonds & Heat
Bond breaking → absorbs energy (endothermic)
Bond forming → releases energy (exothermic)
Activation energy: minimum energy needed to start a chemical reaction
Heat Changes
| Type | Meaning |
|---|---|
Heat of formation | Heat evolved or absorbed when 1 mole is formed |
| Heat of neutralization | Heat liberated when 1 mole of H+ ion from acid reacts with 1 mole of OH- from a base to form 1 mole of H2O |
Heat of combustion | Heat liberated when 1 mole burns |
Heat of solution | Heat when substance dissolves |
Thermodynamics
Study of heat and energy is known as Thermodynamics
First Law:
△U = q - w
Second Law:
A reaction is spontaneous if there is an increase in entropy of a system
Entropy (S)
Measure of randomness
Solid → Liquid → Gas = Entropy increases
S = S{products} - S{reactants}
Gibbs Free Energy
△G = △H - T△S
| Value of ΔG | Meaning |
|---|---|
| Negative | Reaction is spontaneous |
| Zero | System at equilibrium |
| Positive | Not spontaneous |
Important Tip
A reaction is spontaneous when ΔG is negative