Unit 7- Kinetics & Equillibrium

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*Any code blocks () will represent subscripts. Ex: Carbon Dioxide= CO2

Reaction Rates:

→ The speed at which a reaction occurs.

Collision Theory:

• Reactions occur when the particles collide, in the perfect way.
  • The more collisions the faster a reaction occurs.

Influencing Factors:

• Concentration
  • When there is more substance, then there are more collisions, causing the reaction to happen faster.
  • Larger concentration = faster reaction rate
• Temperature
  • At higher temperatures, particles move around more quickly causing more collisions, and a faster reaction rate. (and the opposite happens with lower temps).
• Nature of reactants
  • Some types of reactants are more reactive than others.
  • Ionic compounds are more reactive than covalent.
    • 2 Ionic compounds would react faster than 2 covalent.
• Surface Area
  • When there is more area, more collisions will occur faster.
  • Powdery substances will react the fastest.
• Pressure (Mostly for gases)
  • When the pressure is higher, there is less space between particles, causing more collisions.
  • Opposite for lower pressure
• Catalysts:
  • The area where a reaction occurs, often causes a reaction, and the presence of a catalyst will force collisions to happen, speeding up the reaction.
  • Think about Enzymes in Biology.

Potential Energy: (Diagrams)

→ Graphs that show the Potential Energy as reactions moves forward

Potential Energy → stored energy in a given parts of system.

The Diagrams:

• There are three types, one is for Exothermic reactions, the other is for Endothermic reactions. There are also diagrams which show reactions with catalysts.

The symbols:

• on the graph Delta (The triangle)H, is the Heat of Reaction, or it can also be called the change in enthalpy.
  • To find it you take the energy of the products minus the energy of the reactants.
• There is also Ea, which means activation energy. The Energy needed to start the reaction.
• There is also the activated complex, which is the total energy of the reaction.

Exothermic:

• In this diagram the Reactants have more energy than the products, so it releases energy.
  • The change in enthalpy would be negative.
• The activation energy is also very little for this reaction.

Endothermic:

• The energy of reactants is lower than the products, so it absorbs energy.
  • The change in enthalpy would be positive.
• The activation energy is a lot higher for this reaction

With a catalyst:

→ This is what an Endothermic reaction would look like with a catalyst

• The solid line is the reaction without a catalyst, and the dotted line is with it,
  • The activation energy and activated complex gets lowered.
  • (The change in enthalpy remains the same).
• It would look very similar for an Exothermic reaction. (just lowers the curve).

Enthalpy and Entropy:

• Enthalpy → The energy of a system.
  • Nature generally prefers lower Enthalpy.
• Entropy → The chaos/disorder in a system.
  • Nature generally prefers higher Entropy.
  • Gasses tend to be the most chaotic, of the states. Solids are the least.

Closed and Open Systems:

• Closed System → Only energy can be transferred through the boundary, not matter.
• Open System → Both Matter and energy can be transferred through the boundary.

Chemical Equilibrium:

→ When the forward and reverse reaction are happening at the same time at the same rates.

• Usually represented by “⇌”
• Does NOT mean that the quantities of both reactions are equal. Just the rates are constant.

There is also solution equilibrium, and phase equilibrium.

• Solution Equilibrium → In a supersaturated solution the solute is cycling through dissolving and crystallization. (The solid at the bottom is constantly changing)
  • The dissolving and crystallizing rate are the same, and remain constant.
• Phase Equilibrium →when two phases exist at the same time.
  • Like at 0 C a substance is both liquid and solid.

Le Chatelier’s Principle:

→ Describes how systems of equilibrium deal with added stress.

The three types of stress which can be added are:

1. Concentration
2. Temperature/ Heat or Energy.
3. Pressure (for gasses)

• When these stresses are added to or taken away from a system, it will either produce more products or more reactants until it balances out.
  • When more products are made, the forward reaction rate get higher. This is called “favoring the products”.
  • When more reactants are made, the reverse reaction rate gets higher. This is called, “favoring the reactants”

Concentration:

• When you add concentration the system will shift away from where something was added.
  • If you add more reactant, the products are favored.
  • If you add more products, the reactants are favored.
• The opposite happens when you decrease the concentration. The reaction will shift toward the area which was decreased.
  • If you take away reactants, the reactants are favored.
  • If you take away products, the products are also favored.
• The system will favor one side or the other until it can balance itself out.

Example:

• 2H2 (g) + O2 (g) ⇌ 2H2O (l)
  • If we add more Hydrogen, then more Water will be made.
  • If we add more water, more Hydrogen and Oxygen will be made.
  • If we decrease the amount of Hydrogen, then more Hydrogen and Oxygen will also be made.

Temperature:

• Which side is favored when heat is added or taken away, depends on if the reaction is exothermic or endothermic. It will fill in where heat is missing.
  • For Endothermic, if heat is on the reactant side, so if heat is added, the products would be favored.
  • If heat is taken away, the reactant will be favoured.
  • For Exothermic reactions its the opposite. The heat is on the product side, so adding heat will favor the reactants.
  • If heat is taken away, the products will be favoured
• A good way to think about it is that, the system will balance out the heat. So it will favour the side that has less heat.

Example:

• PCl5(g) + Heat ⇌ PCl3 (g) + Cl2 (g) [Endothermic]
  • If heat is added, more PCl3 and Cl2 will be formed
  • If heat is taken away, more PCl5 will be formed.
• 2HI (g) ⇌ H2 (g) + I2 (g) + Heat [Exothermic]
  • If heat is added, more HI will be formed
  • If heat is taken away, more H2 and I2 will be formed.

Pressure:

• The pressure will affect which side of a reaction has more or less moles of gas.
  • If Pressure increases, the side with less moles of gas will be favoured.
  • If Pressure decreases, the side with more moles of gas will be favoured.
• This also works with volume.
  • When Volume is decreased, the pressure increases, so the side with less moles is favoured.
  • When volume is increased the pressure decreases, so the side with more moles is favoured.

Example:

• H2 (g) + 3N2 (g) ⇌ 2NH3 (g)
  • The reactants has 4 moles of gas, and the products has 2 moles of gas.
  • So if pressure is increased, the products would be favoured, since they have less moles. (same if volume is decreased)
  • If the pressure is decreased, the reactants would be favoured, since they have more moles. (Same if volume is increased)
• 2HBr (g) ⇌ H2 (g) + Br2 (g)
  • Both the reactants and the products have 2 moles of gas.
  • Increasing or decreasing the pressure won’t change the rate because both sides have the same amount of moles of gas.

Next Unit: Unit 8- Acids and Bases