HEAT ENERGY & CHEMICAL REACTIONS – AT A GLANCE

 

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 → MgO
  
  

• Neutralization
  
  HCl + NaOH → NaCl + H₂O

• Dissolving 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