Rates And Equilibrium

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• Exothermic: energy is transferred from system to surroundings
• Endothermic: energy is transferred from surroundings to system

Collision Theory:

1. For a reaction, particles must collide with sufficient energy that is greater than activation energy, and in a favourable orientation
2. More successful collisions per second leads to a greater reaction rate
3. Factors affecting rate: 1. Concentration of reactants 1. Higher concentration means more particles in the same/smaller space. Increase frequency of collisions leads to increased frequency of successful collisions which leads to a higher rate of reaction 2. State of subdivision 1. Increased surface area leads to more particles able to collide, which increases frequency of collisions and this leads to increased successful collisions which leads to a higher rate of reaction 3. Temperature 1. Higher average kinetic energy leads to higher frequency of collisions. There are also more particles with sufficient energy to over come activation energy, thus proportion of successful collisions increases. Theses conditions lead to an increase in frequency of successful collisions, and thus rate of reaction increases 4. Catalyst 1. A catalyst provides an alternative reaction pathway with a lower activation energy. Thus, the proportion of successful collisions increases. Thus, frequency of successful collisions increases, and thus rate of reaction increases

Partial Pressure

• Partial pressure is the pressure a gas in a mixture would exert if it was the only gas occupying the space
• Directly proportional to concentration, as can be seen below $PV=nRT$ $\therefore P=\frac{n}{V}RT$

Systems

• A chemical reaction is regarded as a system, with everything around it (i.e. rest of universe) regarded as the surroundings
• In an endothermic reaction, the system absorbs energy from the surroundings
  • i.e. energy is transferred from surroundings to system
• In an exothermic reaction, the system releases energy to the surroundings
  • i.e. energy isIf transferred from system to surroundings
• Open systems: matter and energy can be exchanged with the surroundings
• Closed systems: energy but not matter can be exchanged with the surroudings

Reactions

• When some chemical reactions occur, the products cannot be converted back into the reactants
• Such reactions, which occur in only one direction, are called irreversable reactions
• However, other reactions are called reversible reactions, where the reactants can be converted back into the reactants, once they are formed
• Examples of reversible systems
  • Evaporation and condensation of water
  • Saturated sugar solution
  • Reaction of haemoglobin and oxygen gas
• Reversible reaction: reaction in which forward and reverse reaction can occur
• Non-reversible reaction: reaction in which only forward reaction can occur

Reversible Reactions

• When particles collide, the energy associated with collisions can break bonds in the reactants, causing them to rearrange to form products
• The energy required to break the bonds is known as activation energy
  • Activation energy: the amount of energy required for a chemical reaction to occur
• If newly formed products collide with enough energy to break their bonds (equal to activation energy of reverse reaction), then it is possible to form the original reactants
• If a forward reaction is endothermic, the reverse reaction is exothermic, and vice versa

Equilibrium

• Must have a closed system (no transfer of matter)
• Occurs when the rate of forward reaction is equal to the rate of reverse reaction

Dynamic Equilibrium

• The rate of forward reaction is equal to the rate of reverse reaction
• However, reactions are still occurring, hence we say that equilibrium is dynamic
• Forward reaction is simultaneously occurring with reverse reaction

Equilibrium Constant

For an equation $aA+bB\rightleftharpoons cC+dD$

The Equilibrium constant

$K=\frac{[C{]}^c[D{]}^d}{[A{]}^a[B{]}^b}=\frac{[\text{Products}{]}^{\text{concentration}}}{[\text{Reactants}{]}^{\text{concentration}}}$ where the numerator is the products, and the denominator is the reactants. $[X]$ is the concentration of element $X$

$K$ changes with temperature (only) Only gases and solutions have concentrations which change. Solids and liquids are given a value of 1.

Heterogenous - different states, only include gases and solutions (as above)

Homogenous - all states the same, all reactants and products are included $[X]$ denotes relative proportion in the liquid mix, rather than concentration

High K: products formed, reaction close to completion

Low K: reactants favoured, reaction barely progressed

Reaction Constant

It is the value for $K$ if not at equilibrium

If Q < K Forward reaction is going faster to approach equilibrium (need more products)

If Q > K Reverse process is happening faster to get to equilibrium (need more reactants)

tldr for a system at equilibrium

• Rate of forward reaction is equal to the rate of reverse reaction
• Q = K
• Dynamic equilibrium