Chemical Equilibrium and Reaction Rates Flashcards

Equilibrium

• Equilibrium is reached when the concentrations of reactants and products no longer change. It's important to note that this doesn't mean the reaction has stopped; rather, the forward and reverse reactions occur at the same rate, maintaining constant concentrations of reactants and products.
• Reversible reactions: These reactions can proceed in both forward and backward directions, indicated by a double arrow ([\rightleftharpoons]).
• Dynamic equilibrium: Equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction. At this point, the net change in concentrations of reactants and products is zero.
• Reactions at equilibrium are dynamic, not static. The forward and reverse reactions continue to occur, but at equal rates.
• Analogy: Imagine a balanced tug-of-war where both sides are pulling with equal force, resulting in no net movement.

Equilibrium Constant

• For a reversible reaction $aA + bB \rightleftharpoons cC + dD$, the equilibrium constant ($K<em>{eq}$) is defined as: $K</em>{eq} = \frac{[C]^c [D]^d}{[A]^a [B]^b}$
  • This equation relates the concentrations of reactants and products at equilibrium.
  • Where $[]$ denotes concentration, typically measured in molarity (M), which is moles per liter (mol/L).
  • The value of $K<em>{eq}$ indicates the extent to which a reaction will proceed to completion. A large $K</em>{eq}$ favors product formation, while a small $K_{eq}$ favors reactants.

Le Chatelier's Principle

• Le Chatelier's Principle states that if a stress is applied to a system at equilibrium, the system will shift in a direction that relieves the stress.

Change in Concentration

• Increase in substance concentration: The system will shift away from the added substance to consume it.
  • Example: If you add more reactant, the equilibrium will shift towards the product side to use up the excess reactant.
• Decrease in substance concentration: The system will shift towards the removed substance to replenish it.
  • Example: If you remove a product, the equilibrium will shift towards the product side to produce more of it.

Change in Volume/Pressure

• Boyle's Law: At constant temperature, decreased volume leads to increased pressure (inverse relationship).
  • More moles of gas: Pressure decreases if volume increases, and vice versa.
  • Fewer moles of gas: Pressure increases if volume decreases, and vice versa.
  • Equal moles of gas on both sides of the reaction: Changes in volume or pressure have no effect on the equilibrium position.

Change in Temperature

• Increased temperature: The equilibrium shifts away from heat, favoring the endothermic reaction (heat is absorbed).
  • Decreased temperature: The equilibrium shifts towards heat, favoring the exothermic reaction (heat is released).

Adding a Catalyst

• Catalysts speed up reactions in both directions, allowing equilibrium to be reached faster. However, catalysts do not shift the position of equilibrium; they only affect the rate at which equilibrium is achieved.

Reaction Rates

• Reaction rate is influenced by several factors:
  • Higher concentrations of reactants lead to more frequent molecular collisions, thus increasing the reaction rate.
  • Increased temperature increases the average kinetic energy of molecules, resulting in more frequent and effective collisions and a faster reaction rate.
  • Gases can be concentrated by compression, which increases the frequency of molecular collisions and thus the reaction rate.

Factors Increasing Reaction Rate

• Heating substances: Increases kinetic energy and collision frequency.
• Using higher concentrations: More molecules available to react.
• Mixing in a more finely divided form: Increases surface area for reactions to occur.
• Adding a catalyst: Lowers the activation energy required for the reaction.

Reaction Curves

• Graph 1 (Reactants): Concentration decreases over time; the steepest slope indicates the fastest reaction rate at the beginning because there are more reactant molecules available.
• Graph 2 (Products): Concentration increases over time; the steepest slope indicates the fastest reaction rate at the beginning as products are formed.

Measuring Reaction Rates

• To monitor the progress of a reaction, measure a characteristic of a product or reactant, such as:
  • Change in concentration over time.
  • Change in pressure (for gaseous reactions).
  • Change in pH (for reactions involving acids or bases).

Energy and Reactions

• Activation Energy: The minimum energy required for a reaction to occur. It's the energy barrier that must be overcome for reactants to transform into products.

Catalysts

• Catalysts are substances that speed up chemical reactions without being consumed in the process. They provide an alternative reaction pathway with a lower activation energy.

Types of Catalysts

• Homogeneous catalysts: These are uniformly mixed with the reactants in the same phase (e.g., all in solution).
• Heterogeneous catalysts: These are in a different physical form from the reactants (usually solid catalysts with liquid or gas reactants). The reaction occurs on the surface of the catalyst.

How Catalysts Work

• Catalysts enable intermediate molecules (activated complexes) to form with lower energy, reducing the activation energy of the reaction.
• Heterogeneous catalysts provide a surface (active sites) for reactions to occur. Reactant molecules can adsorb onto the surface, facilitating bond breaking and formation.

Catalysts and Inhibitors

• Catalysts increase reaction rates by lowering activation energy, while inhibitors slow them down by interfering with the reaction mechanism.
• Inhibitors can:
  • Interfere with catalyst function (e.g., poisoning the catalyst).
  • Tie up reactants, making them unavailable for the reaction.

Worksheet - Reaction Rates

Reaction: $C<em>6H</em>{12}O<em>6 (s) + 6 O</em>2(g) \rightarrow 6 H<em>2O (g) + 6 CO</em>2 (g)$

1. Concentrations:

• $C<em>6H</em>{12}O<em>6$ and $O</em>2$ decrease as the reaction proceeds, indicating they are being consumed.
• $H<em>2O$ and $CO</em>2$ increase as the reaction proceeds, indicating they are being produced.

1. Collision Theory:

• For $C<em>6H</em>{12}O<em>6$ and $O</em>2$ to react, they must collide with the correct orientation and sufficient energy to overcome the activation energy barrier.

1. Activation Energy:

• Activation energy is the minimum energy required to initiate the reaction. It is needed to break the bonds in the reactants and form the activated complex.

1. Changes in Conditions:

   a. Increasing the temperature:

   • Increases the reaction rate because molecules move faster and collide more frequently and with greater energy.

   b. Increasing the concentration of $C<em>6H</em>{12}O_6$:

   • Increases the reaction rate because there are more molecules available to collide and react.

   c. Decreasing the concentration of $O_2$:

   • Decreases the reaction rate because there are fewer molecules available to collide and react.

   d. Increasing the surface area:

   • Increases the reaction rate because there is more surface area for the reactants to collide.

   e. Decreasing the temperature:

   • Decreases the reaction rate because particles move slower, resulting in fewer and less energetic collisions.

   f. Increasing the pressure:

   • Increases the reaction rate, especially for gaseous reactants. Higher pressure leads to a higher concentration of gas molecules, increased collision frequency, and often a higher temperature.

   g. Decreasing the concentration of $H_2O$:

   • Increases the reaction rate because the equilibrium shifts to the right to produce more products, including $H_2O$.

   h. Increasing the volume:

   • Decreases the reaction rate because particles are further apart, reducing the frequency of collisions.

   i. Increasing the concentration of $CO_2$:

   • Decreases the reaction rate because the equilibrium shifts to the left, favoring the reactants.

   j. Using a catalyst:

   • Increases the reaction rate by lowering the activation energy required