• Define heat of reaction (∆H) as the net change of chemical potential energy of the system.
• Define exothermic reactions as reactions which release a net amount of energy, and transfer potential energy into thermal energy.
• Define endothermic reactions as reactions which absorb a net amount of energy, and transfer thermal energy into potential energy.
• Identify that bond breaking is endothermic and that bond formation is exothermic. 
• Identify that a stronger bond requires more energy to break and releases more energy when formed.
• State that ∆H > 0 for endothermic reactions
• State that ∆H < 0 for exothermic reactions

• Do calculations of bond energies

• Define activation energy as the minimum energy required to start a chemical reaction.
• Colliding molecules must have THESE PROPERTIES before a reaction can take place:
  • the correct orientation, 
  • a kinetic energy equal to or bigger than the activation energy of a reaction 
• Define the activated complex as a temporary transition state between the reactants and the products

• Draw or interpret fully labelled sketch graphs (potential energy versus course of reaction graphs) of catalysed and uncatalysed endothermic and exothermic reactions (labeled with reactants, products, activation energy, activated complex and ∆H). 

• A strong chemical bond would be a difficult bond to break. 
• Bond energy is therefore a measure of the strength of a bond.

= E_absorbed - E_released

= (435+243) - 2(431)

= 678 - 862

= -184 kJ/mol => EXOTHERMIC REACTION

Definition:

ΔH = The net change of chemical potential energy of a system during a chemical reaction, measured in kJ.mol^-1.

The overall (net) change of energy in a reaction shows: 

• how much energy must be supplied to reaction systems in order to keep the reaction going 

• or how much chemical potential energy will be released as thermal energy during a reaction

• This change is known as the heat of reaction or enthalpy

A substance that increases the rate of a reaction but remains itself unchanged at the end of the reaction.

A catalyst provides an alternative route of reaction with lower energy than the activation energy of the uncatalyzed reaction. So, more molecules have sufficient energy to undergo effective collisions.

Steps to catalysis:

• Adsorption - reactants attach onto catalyst surface. 

• Alignment - reactant atoms align for effective overlap.

• Desorption - products detach from catalyst.

  
  

• A substance that speeds up the rate of a reaction by providing an alternative energy pathway with lower energy than the original activation energy of the reaction; it is not chemically changed by the end of the reaction.

• they collide with sufficiently high kinetic energy (to overcome the activation energy threshold).

• the molecules have the correct orientation.

• No reaction happens if 2 criteria above not met.

However, there is one way to make a reaction easier:

• ADD A CATALYST!

• Catalysts do not lower the activation energy, but they provide an alternative pathway with lower energy than the original activation energy

• This will allow the reaction to take place quicker/easier.