Unit 8: Kinetics and Equilibrium

Potential Energy Diagrams

• Negative $\Delta$$$\Delta$$H = exothermic reaction = energy/heat is on the product side of the chemical equation.

• Positive $\Delta$$$\Delta$$H = endothermic reaction = energy/heat is on the reactant side of the chemical equation.

Rates of Reaction

• Collision Theory

  • Atoms and/or molecules must collide in order to react.

  • They must collide in a correct orientation; otherwise, no reaction will occur.

  • They also must have enough energy to form products.

Factors that Affect the Rate of Reaction

• Presence of a Catalyst – lowers the activation energy of the reaction so the reaction begins easier (less energy required for a successful collision). Also, a catalyst provides an alternate pathway for the reaction to proceed.

• The Nature of the Reactants – The speed of a reaction sometimes depends upon the types of bonds between the atoms. Aqueous ionic compounds generally react faster than covalently bonded compounds.

  • Aqueous ionic compounds are already broken down into ions (loose Lego pieces), making forming bonds much faster.

  • Covalently bonded compounds would need to break existing bonds first (pre-built Lego structure) and then form new ones (different completed Lego structure), making this process slower compared to aqueous ionic.

• Concentration - Higher concentration leads to more effective collisions between the molecules (higher probably of a successful collision).

  • Note* - As a reaction proceeds, the concentration of the reactants will gradually decrease and the reaction will slow down. This is because you are using up reactants to produce products.

• Temperature - At a higher temperature the molecules will collide more frequently and with more kinetic energy (sufficient energy is required for a successful collision).

• Surface Area - Smaller pieces of a larger portion react faster than the larger portion as a whole because more surface area has been exposed. A large clump doesn’t have the inside of the clump exposed so it cannot react.

• Pressure - An increase in pressure will move the gas molecules closer together. This will make more collisions and increase the rate of reaction. Pressure only affects gases!

Phase Equilibrium

• Phase changes are reversible in a closed system and at equilibrium (the rate of the forward and reverse reactions is equal, and the concentrations of reactants and products will be constant) may be attained.

  • Ex. H2O(l) $\rightleftharpoons$$$\rightleftharpoons$$ H2O(g)

  • Ex. 2

  • When phase equilibrium is reached, for every liquid molecule that enters the gas phase, a gas molecule will enter the liquid phase.

  • Provided that the solution is saturated, for every molecule of solute that enters solution, a molecule of solute in solution will precipitate out of solution (for every undissolved molecule that becomes dissolved, a dissolved molecule will become undissolved). While the rates are equal, the amounts (concentrations) of excess solute and solute dissolved do not have to be equal.

Chemical Equilibrium

• In a closed system, where no reactants, products, or energy can be added to or removed from the reaction, a reversible reaction will reach equilibrium.

• At equilibrium, the rate of the forward reaction becomes equal to the rate of the reverse reaction.

• At equilibrium, the concentrations of each species remains constant.

• Chemical Equilibrium: Forward and reverse reactions are occurring at the same time at the equal RATES. Concentration of reactants and products appear to be CONSTANT, not necessarily equal.

• Stress - any kind of change in a system that upsets (disturbs) the equilibrium. Ex. Pressure, Volume, Temperature, Concentration, Addition of a catalyst

Le Châtelier’s Principle

• Equilibrium position shifts as a result of changing conditions (added stress). When a system in chemical equilibrium is disturbed by a change of temperature, pressure, or a concentration, the equilibrium will shift in such a way that tends to relieve the effects of the stress and a new point of equilibrium is established (in a closed system).

• The equilibrium always responds in such a way to counteract the stress.

Le Chat’s Tube

• Imagine (generic reaction at equilibrium) the tube below that is closed off on all sides except four openings that are closed using caps.

• The dotted line visually separates the reactants half of the tube from the products half of the tube.

• The tube is filled with liquid. To add liquid, you need to open one of the top two caps. To remove liquid, you would need to open one of the bottom two caps.

  • Example 1 (add B, a reactant) : equilibrium shifts to the right (the liquid flows to the right)

  • Example 2 (remove A, a reactant) : equilibrium shifts to the left (the liquid flows to the left)

  • Example 3 (remove C, a product) : equilibrium shifts to the right (the liquid flows to the right)

  • Example 4 (add D, a product) : equilibrium shifts to the left (the liquid flows to the left)

  • Example 5 (increase heat) : equilibrium shifts to the right (the liquid flows to the right) This depends on which side of the equation ‘energy/heat’ is on aka is it endothermic or exothermic.

  • Example 6 (decrease heat) : equilibrium shifts to the right (the liquid flows to the right) This depends on which side of the equation ‘energy/heat’ is on aka is it endothermic or exothermic.

Effect of an Equilibrium Shift

• An equilibrium shift to the right means the following:

  • An increase in the concentration of ALL products (favors the forward reaction)

  • A decrease in the concentration of ALL reactants (disfavors the reverse reaction)

• An equilibrium shift to the left means the following:

  • An increase in the concentration of ALL reactants (favors the reverse reaction)

  • A decrease in the concentration of ALL products (disfavors the forward reaction)

• Changing concentration of reactant or product.

• Changing temperature of reactant or product.

Increasing Pressure (Decreasing Volume)

• Think Boyle’s Law! $P1V1 = P2V2$$$P1V1 = P2V2$$

• Pressure and volume have an inversely proportional relationship.

  • Therefore, if you increase pressure, then volume decreases.

  • Pressure affects only gases of an equation at equilibrium

  • An increase in pressure causes the equilibrium to shift in the direction that has the fewer number of moles (less volume). The coefficient of each species indicates how many moles you have of that species!

Examples:

• $N2 (g) + 3H2 (g) \rightleftharpoons 2NH3 (g) \Delta H = -92 kJ$$$N2 (g) + 3H2 (g) \rightleftharpoons 2NH3 (g) \Delta H = -92 kJ $$ is the same as:$N2 (g) + 3H2 (g) \rightleftharpoons 2NH3 (g) + 92 kJ$$$ N2 (g) + 3H2 (g) \rightleftharpoons 2NH3 (g) + 92 kJ$$

  • Recall: negative (-) $\Delta$$$\Delta$$H is exothermic : energy is a product Consider heat as a reactant or product.

• $N2 (g) + O2 (g) \rightleftharpoons 2NO (g) \Delta H = +183 kJ$$$N2 (g) + O2 (g) \rightleftharpoons 2NO (g) \Delta H = +183 kJ$$ is the same as: $N2 (g) + O2 (g) + 183 kJ \rightleftharpoons 2NO (g)$$$N2 (g) + O2 (g) + 183 kJ \rightleftharpoons 2NO (g)$$

  • Recall: positive (+) $\Delta$$$\Delta$$H is endothermic : energy is a reactant Consider heat as a reactant or product.

• $3O2(g) \leftrightharpoons 2O3(g)$$$3O2(g) \leftrightharpoons 2O3(g)$$

  • An increase in pressure results a shift to the right, therefore there in a decrease in O2 and an increase in O3. This is because on the reactant side, there are 3 moles of gas (coefficient of all gases). On the product side, there are 2 moles of gas (coefficient of all gases). That means an increase in pressure will shift the equilibrium to the side with less moles of gas, the products. Also, a decrease in pressure will shift the equilibrium to the side with more moles of gas, the reactants.

Effect of a Catalyst

• Catalysts affect both the forward and reverse directions equally.

• A catalyst does not change the point of equilibrium but reduces the time required for the system to come to equilibrium.

• It lowers the activation energy by providing an alternate pathway still also speeds up the rate of both the forward and reverse reactions.

• Homogeneous Equilibrium - All reactants and products are in the same state of matter

• Heterogeneous Equilibrium - Reactants and products are in different states of matter

Spontaneous Processes

• Processes that can proceed without little to no outside intervention.

  • Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.

  • Depends on Enthalpy (H) and Entropy (S)

  • Processes that are spontaneous at one temperature may be nonspontaneous at other temperatures. Such as ice melting and ice freezing both being spontaneous processes that occur at different temperatures.

Entropy

• A measure of the randomness, disorder, messiness of a system. (chaos)

  • The change in entropy is designated with a $\Delta$$$\Delta$$S. ($\Delta$$$\Delta$$ means change in)

  • $\Delta$$$\Delta$$S (+) means that there is an increase in disorder/chaos, $\Delta$$$\Delta$$S (-) means that there is a decrease in disorder/chaos.

  • $\Delta$$$\Delta$$S = Sfinal - Sinitial

  • The entropy tends to increase with increases in: Temperature, Volume, more molecules formed in product, phase of matter change to a more mobile state Such as S(g) > S(l) > S(s).

  • Generally, when a solid is dissolved in a solvent, entropy increases. But when gas is dissolved in solvent, entropy decreases.

Nature’s Preferences

• Reactions tend to proceed in the direction that decreases the energy of the system (H, enthalpy). $\Delta$$$\Delta$$H= “-”

• Reactions tend to proceed in the direction that increases the disorder of the system (S, entropy). $\Delta$$$\Delta$$S= “+”