2.1 chemical equilibrium

• Some reactions never go to completion but are in a state of dynamic equilibrium. The reaction needs to be in a closed system.

  At equilibrium the concentrations of reactants and products are constant meaning both forward and backward reactions do not stop. Both rates of reaction (forward and backward) are equal. There is no change in composition of the equilibrium mixture at dynamic equilibriu.
• Increasing concentration of reactants or removing the product shifts equilibrium to the right.
• Increasing the concentration of products or removing the reactant shifts equilibrium to the left.
• Increasing temperature shifts equilibrium in the direction of the endothermic reaction.
• Decreasing temperature shifts the equilibrium in the direction of the exothermic reaction.
• Catalysts have no effect on equilibrium position, but they speed up the rate of reaction.
• The equilibrium constant is given by the symbol K and can be described by the general expression: $K=$$[products]/[reactants]$

  The square brackets symbolise the concentration is the species present at equilibrium.

  The reversible reaction here is: $reactants ⇌ products$
• Equilibrium constants are independent of changes of concentration or partial pressure of a species in any given reaction.

  When these change, the reaction also changes to equilibrium position until the ratio is re established.
• The equilibrium position is used to describe how an equilibrium solution behaves when reaction concentration changes. So K can stay constant while the position of equilibrium can shift and adjus.

  K states where the thermodynamically most favourable ratio of reactants and products is.
• Knowing the values of the equilibrium constant, K, allows us to determine:

  • The direction in which a reaction will proceed to achieve equilibrium.
  • The ratio of the concentration of reactants and products when equilibrium is reached.
• • If K is less than 1, the equilibrium lies more to the reactants therefore the equilibrium concentration of the reactants is large.
  • If K is greater than 1, the equilibrium lies more to the products therefore the equilibrium concentration of the products is large.
  • If K is approximately 1, then neither reactants or products are favoured.
• • For endothermic reactions, a rise in temperature causes an increase in K and the yield of the product is increased.
  • For exothermic reaction, a rise in temperature causes a decrease in K and the yield of the product is decreased.
• Catalysts speed up the rate of the forward reaction and backward reaction equally and so have no effect on the equilibrium position. There is also no effect on the equilibrium constant and only allows the equilibrium position to be achieved at a faster rate.
• Acids have a H+ ion which is what makes the substance acidic. Substances with a lower pH value have a high concentration of H+ ions.

  Considering that H+ is essentially a proton, we can describe acidic substances as a proton donor.
• In an aqueous solution, hydrogen ion only exists when surrounded by water molecules. These are known as hydronium ions and can be written as H3O+(aq).

  When we talk about H+ ions we are really talking about the hydronium ion.
• If an acid is seen as a proton donor then a base is a proton acceptor.

  When an acid donates a proton, the species left is known as the conjugate base.

  When a base accepts a proton, the species formed is the conjugate acid.
• Water dissociates very slightly. The ionization of water can be written as: $H2O(l) +$$H2O(l) ⇌ H3O+$$(aq) +$$OH-(aq)$
• In pure water, the concentration of hydroxide and hydronium ions are equal and the pH is neutral. Water can be called amphoteric, as it can behave as an acid or a base.
• The dissociation constant for the ionisation of water is known as the ionic product and is represented by Kw : 𝐾𝑤 = [𝐻3𝑂 + (𝑎𝑞)] [𝑂𝐻− (𝑎𝑞)] or simply 𝐾𝑤 = [𝐻 + (𝑎𝑞)] [𝑂𝐻−𝑎𝑞 ]

  The reaction is temperature dependant and has a value of 1 x 10-14 at 25°C.
• If [H+] > [OH-], the solution will be acidic.

  If [H+] < [OH-], the solution will be alkaline.

  If [H+] = [OH-], the solution will be neutral.
• The pH of any aqueous solution can be calculated using the expression: pH = -log10[H3O+] and [H3O+] = 10-pH

  The pH scale is a logarithmic scale; so any change in pH requires a factor of 10 change in the hydrogen ion concentration. We can therefore calculate the concentration of H+, OH- or the pH for acidic and alkaline solutions.