HS Chemistry - Reversible Reactions

Unit Summary

Reversible Reactions & Dynamic Equilibrium

• In a reversible reaction, both the forward and reverse reaction occur at the same time, although not necessarily at the same rate.

• Dynamic equilibrium is when the rate of the forward reaction equals the rate of the reverse reaction.

  • At that point, the concentration of reactant and product molecules remains constant, since any change by the forward reaction is negated by the reverse reaction, and vice versa.
  • Any system that isn't in equilibrium will correct itself until it is, with the side that is in surplus reacting until enough of the other side's molecules are formed to reach equilibrium.
  • Dynamic equilibrium requires a closed system so that matter doesn't enter or escape.

• The point at which equilibrium is achieved in a reversible reaction is called the position of equilibrium. It indicates whether there are more products or reactants formed, and by extension, which side the reaction favors.

• Le Chatelier's principle is that if one or more of the factors that affect equilibrium changes, the position of equilibrium shifts in the direction that opposes $reduces$ the change.

  • Concentration:

    • When the concentration of molecule$s$ on one side is increased, that side will react more quickly than the other side. This will cause more molecules of the other side to be formed before equilibrium is re-established, shifting the position of equilibrium away from the side whose concentration was increased.
    • Conversely, when the concentration of molecule(s) on one side is decreased, that side will react more slowly than the other side. This will cause more molecules of that side (the one whose concentration was reduced) to be formed before equilibrium is re-established, shifting the position of equilibrium towards from the side whose concentration was decreased.

  • Pressure $for gases only$:

    • When pressure is increased, whichever side has a greater number of molecules in the equation ends up with more reacting molecules closer together. That side then reacts more, causing more molecules of the other side to be formed before equilibrium is re-established, and shifting the position of equilibrium to the side with less molecules.
    • Conversely, when pressure is decreased, whichever side has less molecules in the written equation will react more, shifting the position of equilibrium to the side with more molecules in the written equation.

  • Temperature:

    • Every reaction has an associated enthalpy change. In a reversible reaction, either the forward or reverse reaction must be exothermic, and therefore the other one has to be endothermic, as it simply reverses the other one.
    • As the temperature of a system increases, there is more heat available in the surroundings, so the endothermic reaction's rate increases. That side reacts more, causing more molecules of the other side to be formed before equilibrium is re-established and shifting the position of equilibrium away from the side that reacts endothermically.
    • Conversely, as the temperature of the system decreases, the exothermic reaction starts reacting more, causing more molecules of the other side to be formed before equilibrium is re-established and shifting the position of equilibrium towards the side that reacts endothermically.

More About Dynamic Equilibrium

• The equilibrium constant, K_c, of a reversible reaction shows how far the reaction goes towards the products.

• K_c = $\(concentration of product1$ × $concentration of product2$ × $concentration of product3$ × …) ÷ $\(concentration of reactant1$ × $concentration of reactant2$ × $concentration of reactant3$ × …)

  • If K_c > 1, there will be more product molecules than reactant molecules when the system is at equilibrium.
  • If K_c < 1, there will be more reactant molecules than product molecules when the system is at equilibrium.
  • If K_c = 1, the number of product molecules will be equal to the number of reactant molecules when the system is at equilibrium.

• When dealing with gases in a reversible reaction, there is also a pressure constant, K_p:

  • K_p = K_c × RT^∆n

  Where R is the Universal Gas Constant, T is the temperature of the system at equilibrium, and ∆n = $coefficients of all product molecules multiplied together$
  ÷ $coefficients of all reactant molecules multiplied together$

• An ICE table can be used to find the molar concentrations and/or pressure of a substance if we don't know it.

Collision Theory

• The collision theory is the idea that chemical reactions occur when molecules collide with one another.

  • An effective collision is when molecules collide with each other with an energy equal to or greater than the activation energy, causing a reaction between those molecules.

  • An ineffective collision is when molecules collide with each other with an energy less than the activation energy, and a reaction does not occur between those molecules.

• The temperature of a system is a measure of the average kinetic energy of the particles in the system. At higher temperatures, more particles have a higher kinetic energy, which means more particles have an energy greater than the activation energy. More effective collisions will occur when the temperature is higher, and therefore higher temperatures increase the rate of a reaction.

  • The amount of particles in a system with an energy equal to or greater than the activation energy can be visualized using a Boltzmann diagram.

• A higher concentration of reactants will increase the rate of reaction. In a reversible reaction, this can change the position of equilibrium.

• Catalysts decrease the activation energy, and thus increase the rate of reaction. However, this also means that catalysts won't change the position of equilibrium.

  • Catalysts can be classified as homogenous or heterogenous. If the reactants, products, and the catalyst(s) are all in the same state of matter, then the catalysts is homogenous. Otherwise, the catalyst is heterogenous.

• The rate of a reaction can be determined by measuring the change in concentration of its reactants or its products over time.

  • As a non-reversible reaction progresses, there are fewer and fewer reactants available as more products are formed. This decreases the rate of reaction, as there are fewer reactants available. Thus, the formula gives an average reaction rate.

• Another way of measure the rate of a reaction is through the reaction's half-life, which is the amount of time it takes to react 50% of the available reactants present in the system. The half-life remains constant regardless of the concentrations.

  • Using the half-life, the equation's rate law can be found.

Brønsted–Lowry Theory

• A Brønsted–Lowry acid is a proton $H^\+^$ donor.

• A Brønsted–Lowry base is a proton $H^\+^$ acceptor.

• In a neutralization reaction, an acid neutralizes a base, creating a salt $an ionic compound$ and another substance.

• After losing an H⁺ atom, the acid turns into its conjugate base; after gaining an H⁺ ion, the base turns into its conjugate acid.

• The extent to which an acid or base ionizes into its conjugate acid or base in solution is measured by its pH.

  • pH is measured on a scale from 1 - 14, where 7 is neutral. A substance with a pH lower than 7 is acidic, and a substance with a pH higher than 7 is basic.
  • pH = −log[H⁺] where [H⁺] is the concentration of H⁺ ions that dissociate, or ionize, from the substance.
  • The further the pH is from 7, the stronger the acid/base is, and the more it will ionize in solution.
  • Since there will be more free ions in solution, a substance with a pH farther from 7 will conduct electricity better.
  • Stronger acids are more reactive with reactive metals because there are more free ions in their solution.

• Salts $ionic compounds$ also have a pH:

  • If the salt forms an acid when it dissolves in water, the salt is acidic $its pH is less than 7$.
  • If the salt forms a base when it dissolves in water, the salt is basic $its pH is greater than 7$.

• Substances that can function both as acids and as bases are called amphoteric substances.

• The molar concentration of a substance is called its molarity:

  • M = $number of moles$ ÷ $Volume in Liters$

• A titration experiment can be used to find out the pH of an unknown acid or base, using an acid or base with a known pH, a known quantity of that acid/base, and a known quantity of the unknown acid/base.

• The equivalence point of a neutralization reaction is when enough base has been added to an acid to completely neutralize the acid, or when enough acid has been added to base to completely neutralize the base.

• An indicator is a weak acid or base whose conjugate base has a different color in solution than the original acid. It's used in titration experiments to show when the equivalence point has been reached. At the equivalence point, the conjugate base of the indicator will be the only substance affecting the pH $as everything else will cancel each other out$, and so its color will become clearly visible, indicating that the equivalence point has been reached.

Haber Process

• The Haber process is the industrial process used to produce ammonia $NH~3~$.

• It uses the following reaction to produce ammonia:

  • N₂ $g$ + 3N₂ $g$ ⇌ 2NH₃ $g$

• Since the forward reaction is endothermic, in order to maximize the product yield, the temperature at which the reaction is typically carried out is 450 °C.

• The greater the pressure of the system, the more NH₃ the reaction produces. Due to money and safety considerations, the pressure at which the reaction is typically carried out is 200 atm.

• Iron is typically used to catalyze the reaction.