Chemistry lecture 13

• Reaction Quotient (Q)


  An algebraic expression relating the mole ratio concentrations of reactants [R] and products [P] at any time during a chemical reaction
• Q varies with concentrations
• Dynamic equilibrium


  Achieved when rate of forward equals rate of reverse reaction
• Equilibrium constant (Kc)


  The special name for the Reaction Quotient (Q) at dynamic equilibrium
• Kc is constant
• A reaction graph shows the position of equilibrium
• Solids and solvents are not included in the equilibrium expressions
• Metal solid/aqueous ion equilibria

  • If placed in a solution of their own ions, metals will establish an equilibrium
  • Zn(s) ⇌ Zn2+(aq) + 2e-
  • K = [Zn2+]
  • The equilibrium position depends upon the reactivity of the metals
• Phase Change Equilibria


  • Water in a sealed container will soon reach an equilibrium with the water vapour saturating the spaces above the liquid surface
  • H2O(l) ⇌ H2O(g)
  • K = [H2O(g)]
• Effect of solids, liquids and gases on the equilibrium constant k

  • N2(g) + 3H2(g) ⇌ 2NH3(g), Kc = [NH3]2 / [N2][H2]3
  • H2(g) + I2(g) ⇌ 2HI(g), Kc = [HI]2 / [H2][I2]
  • 2SO2(g) + O2(g) ⇌ 2SO3(g), Kc = [SO3]2 / [SO2]2[O2]
  • 2CO(g) + 2NO(g) ⇌ 2CO2(g) + N2(g), Kc = [CO2]2[N2] / [CO]2[NO]2
  • CaCO3(s) ⇌ CaO(s) + CO2(g), Kc = [CO2]
  • CH3COOH(aq) + H2O(l) ⇌ CH3COO-(aq) + H3O+(aq), k = [CH3COO-][H3O+] / [CH3COOH]
• Catalysts


  • Substances which alter the rate of a chemical reaction without themselves being used up in the overall reactions
  • Because catalysts alter the rate of forward and reverse reactions equally, there is NO overall effect on the position of equilibrium
  • The equilibrium constant k is not altered by a catalyst
• Factors affecting equilibrium


  • Concentration of reactants or products
  • Pressure if any gases are involved
  • Temperature
• Concentration


  The ratio of the quantity of solute to either the quantity of solvent or to the quantity of solution
• Temperature-independent concentration units


  Weight Fraction, Weight Percentage (w/w), parts per thousand (ppt), parts per million (ppm), parts per billion (ppb), Molality
• Temperature-dependent concentration units


  Molarity (M) - moles of solute per litre of solvent
• Normality


  A concentration unit that expresses the number of equivalents of solute per litre of solution
• Increasing pressure results in reduction of volume so some gas particles must go into solution (equilibrium shifts right)
• Lowering pressure on solution results in equilibrium shift to the left and some dissolved gas leaves the solution
• Some solutes dissolve more at high temperatures (endothermic reactions)
• Other solutes are soluble at low temperature (exothermic reactions)
• Le Chartelier's Principle


  If a change is made to a reaction at equilibrium, the position of equilibrium alters to oppose the effect of the change
• Important cations


  • sodium (Na+)
  • potassium (K+)
  • ammonium (NH4+)
  • calcium (Ca2+)
  • aluminium (Al3+)
  • Iron(ii) (Fe2+)
  • Iron(iii) (Fe3+)
  • Chromium(iii) (Cr3+)
  • Zinc (Zn2+)
  • Copper(i) (Cu+)
  • Copper(ii) (Cu2+)
  • silver (Ag+)
• The colours of ionic compounds depend on cation(s) present
• Le Chartelier's Principle


  Governs the specific manner in which concentration, pressure, and temperature affect the equilibrium position
• Le Chartelier's Principle


  States that if a change is made to a reaction at equilibrium, the position of equilibrium alters to oppose the effect of the change
• Ionic solids are made up of

  • Cations
  • Anions
• Important cations


  • Sodium (Na+)
  • Potassium (K+)
  • Ammonium (NH4+)
  • Calcium (Ca2+)
  • Aluminium (Al3+)
  • Iron(II) (Fe2+)
  • Iron(III) (Fe3+)
  • Chromium(III) (Cr3+)
  • Zinc (Zn2+)
  • Copper(I) (Cu+)
  • Copper(II) (Cu2+)
  • Silver (Ag+)
• The colours of ionic compounds depend on the cation(s) present
• When copper sulfate dissolves in water it produces Cu2+(aq) - blue and SO42-(aq) - colourless
• Hydrated cations


  Cations become hydrated through ion-dipole attraction, the number of H2O involved depends on the size of the cation, the shape of the hydrated ions depends on the number of H2O involved
• Important anions


  • Chloride (Cl-)
  • Bromide (Br-)
  • Sulfide (S2-)
  • Hydroxide (OH-)
  • Carbonate (CO32-)
  • Nitrate (NO3-)
  • Sulfate (SO42-)
  • Chromate (CrO42-)
  • Dichromate (Cr2O72-)
  • Permanganate (MnO4-)
  • Thiosulfate (S2O32-)
  • Bicarbonate (HCO3-)
• Solution


  A mixture of a solvent and a solute, a homogeneous mixture where all the particles are very small (0.05-0.25 nm), solute particles are larger and regarded as colloid and suspension
• Driving forces for solution formation

  • Randomness - the tendency of a system to become increasingly disordered
  • Solute/solvent attraction - attraction forces between particles are a major factor in solution formation
• Polar solvents dissolve


  • Polar and ionic substances
• NaCl dissolving in H2O

  • Yields Na+(aq) and Cl-(aq) - disordered state
• Sugar dissolving in H2O

  • Yields sugar+(aq) and sugar-(aq) - disordered state
• Iodine dissolving in cyclohexane

  • Both are non-polar covalent, so (solute/solvent) attraction is stronger than (solute/solute) attraction
• Iodine not dissolving in H2O

  • I2/I2 covalent (solute/solute) attraction is stronger than (solute/solvent) attraction between I2 and H2O
• NaCl not dissolving in cyclohexane

  • NaCl is polar and (solute/solute) attraction is stronger than (solute/solvent) attraction between NaCl and cyclohexane
• Colligative properties


  Solution properties which depend only on total concentration of solute particles