Presentation on theme: "Chemical Equilibrium Equilibrium, strong and weak acids and bases, solutions, concentration and pH."— Presentation transcript:
1Chemical EquilibriumEquilibrium, strong and weak acids and bases, solutions, concentration and pH
2Index Equilibrium Le Chatelier’s principle Equilibrium and water Weak and strong acids and alkalis, pHTitration and stoichiometryBasesAqueous solutions and pH
3Equilibrium Reactants Products Graph 2 Graph 1 A reversible reaction can go both ways, both products and reactants existtogether. The final equilibrium position is the same whether you start with theproducts or reactants.Eventually the rate for the forward reaction equals the backward rate. (graph 1)A state of Dynamic Equilibrium is reached. (in a closed system).Depending on the percentage of products to reactants, the equilibrium is said tolie to the left or right. (graph 2)ConcentrationTimeForwardBackwardEquilibrium lies to the rightGraph 2equilibriumReactionrateTimeForward reactionBackward reactionGraph 1
4Reversible Reactions. CoCl2.6H2O CoCl2 + 6H2O pink blue CoCl2.6H2O ForwardCoCl H2OblueCoCl2.6H2OpinkheatReverseCoCl H2OblueCoCl2.6H2OpinkheatEthanoic acid + ethanolEthyl ethanoate waterNo matter which side of the equation you start from the concentrationsin the equilibrium mixture will be the same and the forward and backwardrates will be the same.
5Le Chatelier’s principle Many reactions in the chemical industry are in equilibrium.E.g. Haber Process, Synthesis gas manufacture, Contact Processand Nitric acid manufacture.Concentration, pressure and temperature effect the position of theequilibrium. Their effect can be predicted usingLe Chatelier’s Principle. This states:‘If a system at equilibrium is subjected to a change, the system will adjust to oppose the effect of the change.’1. Catalysts do not change the equilibrium position. Catalysts speed upboth the forward and backward reactions. So the state of equilibriumwill be reached quicker but the amounts of the products and reactantswill stay the same.
6Le Chatelier’s principle 2. Changing concentration Fe3+ (aq) CNS- (aq)FeCNS2+ (aq)Pale yellowcolourlessDeep redTo increase the proportion of the products you can:Increase the amount of reactantsBy increasing the amount of Fe3+ (aq) the forward reaction will increaseproducing more products, however, the reverse reaction will also increaseuntil the equilibrium is reached again.The new equilibrium position will have moved to the right in order toreduce the increased amount of Fe3+ (aq). So the deep red darkens(ii) Decrease the amount of products.By removing Fe3+ (aq), the equilibrium will shift to the left in order totry and replace the Fe3+ (aq) ions that have been removed. So thepale yellow darkens.
7Le Chatelier’s principle 3. Changing Pressure Pressure only effects the equilibrium in reactions which involve gases.Increasing the pressure will cause a shift to an equilibrium mixturewith a smaller number of gas molecules. i.e. less molecules would meanless overall pressure, so, in effect reducing the overall pressure.In the examples below increasing the pressure in example 1 will cause theequilibrium to move to the right in an attempt to reduce the overall numberof molecules, the opposite is true of example 2.Increased pressureExample 1CO (g) + 2H20 (g)CH3OH (g)Decreased pressureDecreased pressureExample 2CH4 (g) + H20 (g)CO (g) + 3H2 (g)Increased pressure
8Le Chatelier’s principle 4. Changing Temperature Chemical reactions are either exothermic or endothermic.Reversible reactions are exothermic in one direction and endothermicin the other.An increase in temperature shifts the equilibrium position to a mixturewhich is formed by absorbing thermal energy. i.e. the equilibrium movesin the direction which will remove the added thermal energy.In examples 1&2, if the temperature is increased the equilibrium will moveto the left, as this is endothermic, so removing the increase thermalenergy.CO (g) + 2H20 (g)CH3OH (g)ExothermicEndothermicExample 1C2H4 (g) + H20 (g)C2H5OH (g) H = -44kJ mol-1 H = + 44kJ mol-1Example 2
9Equilibrium and water H2O (l) H+ (aq) OH - (aq) + Covalent water exists in equilibrium with tiny traces of ionic water,with the equilibrium lying to the left.In fact only 1 in every 555 million water molecules dissociates, ionises.Water can conduct electricity, even in its pure form, it will conductelectricity very slightly because some ions are free to move.pHThis value depends on the concentration of hydrogen ions in solutionThe concentration of H+ ions in pure water is 1 x 10-7 mol l-l a pH of 7.
10Acids and Alkali’s pH Strong acid Weak acid 1234567891011121314pHStrong acidWeak acidFully dissociated in aqueous solutionNot fully dissociated in aqueous solutionWeak alkaliStrong alkaliNot fully dissociated in aqueous solutionFully dissociated in aqueous solutionDiluteConcentratedA little solute in a lot of water.e.g. 0.2 mol l-1A lot of solute in a little of water. e.g. 2.0 mol l-1
11Ionic Product for water = [H+] (aq) x [OH-] (aq) = 10–7 x 10–7 Concentration and pHThe pH range, (p comes from the German word Potenz meaning power and H from [H+]), can vary from –1 to 15.H2O (l)H+ (aq)OH - (aq)+When pure water dissociates one H+ ion is produced for every OH- ion.So [H+] = [OH-]Since the pH is 7 so [H+] = [OH-] = 10–7 mol l-1i.e. [H+] (aq) = 10–7 mol l-1and [OH-] (aq) = 10–7 mol l-1The Ionic Product for water is the product of the concentration of thehydrogen and hydroxide ions.Ionic Product for water = [H+] (aq) x [OH-] (aq) = 10–7 x 10–7= mol2 l-2
12[OH-, [H+] and pHThe table below shows the relationship between [OH-, [H+] and pH .[H+] aqmol l-1[H+]pH[OH-][OH-] aq0.11 x 10-111 x 10-130.0011 x 10-331 x 10-111 x 10-551 x 10-91 x 10-7791113Note the relationship between [H+] and the pH value.
13Strong and weak acidsAcids are substances which dissociate in water to form H+ ions.Strong Acids, e.g. HCl, H2SO4 and HNO3. These are molecular compoundsThese acids dissociate completely, the acid is strongly ionised.H3O+ ion is usuallywritten as H+(aq)HCl(g) H2O (l)Cl- (aq) + H3O+ (aq)The relationship between pH and hydrogen ion concentration:[H+](aq) mol l-1[H+]pH101 x 10 1-10.0011 x 10 -331 x 10 -660.1 mol l-1 of HCl would contain 0.1 molof hydrogen ions (H+ (aq) ).This would mean the concentration ofhydrogen ions would be 1 x 10 –1.Its pH would be 1
14Strong and weak acids Weak Acids, e.g. CH3COOH, HCOOH, H2C03, H2SO3 These acids only partially dissociate, the acid is only partially ionised.CH3COOH H2O (l)CH3COO- (aq) + H3O+ (aq)0.1 mol l-1 of CH3COOH has a pH of 3.This suggests a lower [H+] than the equivalent strength of HCl.Comparison of strong and weak acids, consider 0.1 mol l–1 of:HCl (Strong)CH3COOH (Weak)pH13ConductivityHighLowReaction with MgFastSlowReaction with CaCO3
15TitrationThe neutralisation of an acid by an alkali can be represented by:H+ (aq) OH- (aq)H20 (l)If you titrate 25 ml of 1.0 mol l-l HCl against 1.0 mol l-l NaOH, whatvolume of NaOH would you need?Vacid x Cacid = VAlk x CalkSo 25 x 1 = 25 x Vol of NaOHAns: 25 ml NaOHWhat happens if you use 1.0 mol l-l of weak acid instead?Ans: 25 ml NaOHThe reasonAs OH- ions are been added to the weak acid, H+ ions are been used up.However, as the H+ ions are used up, more of the weak acid will dissociateproducing more H+ ions, which in turn will use up more OH- ions.This will continue until all the molecules of the weak acid have dissociated.
16Stoichiometry of reactions Stoichiometry of a reaction means the numerical proportions of the substances involved.3 Na+OH- (aq)Fe3+ (OH-)3 (aq) Na+NO3- (aq)Fe 3+ (NO-)3 (aq) +3 NH4+OH- (aq)Fe3+ (OH-)3 (aq) NH4+NO3- (aq)Even though a 1mol l-1 NaOH solution has about 250x the concentration of hydroxide ions as 1mol l-1 NH4OH, the weight of Iron(III) Hydroxide precipitate in both reactions is the same.The reasonAs Fe3+ ions are been added to the weak acid, OH- ions are been used up.However, as the OH- ions are used up, more of the weak base will dissociateproducing more OH- ions, which in turn will use up more Fe3+ ions.This will continue until all the molecules of the weak base have dissociated.
17Strong and weak basesBases are substances which can react with H+ ions (aq) to form water.Unlike strong acids, strong bases are ionic and also dissociate completely.Na+ OH- (s)Na+(aq) OH-(aq)waterWeak bases do not completely dissociateNH3 (g) waterNH4OH (aq)NH4+(aq) + OH-(aq)Only around 0.4% of the molecule, NH4OH, dissociatesComparison of strong and weak bases, consider 0.1 mol l–1 of:NaOHAmmoniapH1311-12ConductivityHighlow
18Aqueous solutions (Covalent and Ionic Oxides) Covalent oxidesSulphur dioxide and carbon dioxide both dissolve in water to formweak acids. (Partial dissociation) the equilibrium lies to the left)CO2 (aq) + H2O (l)2H+ (aq) CO32-(aq)Covalent solutionIonic solution (carbonic acid)SO2 (aq) + H2O (l)2H+ (aq) SO32-(aq)Ionic solution (sulphurous acid)Ionic oxidesCalcium oxide and magnesium oxide are soluble basesCa 2+O2- (s) H2O (l)Ca 2+(aq) OH-(aq)All oxide ions are converted into hydroxide ions, because of thiscomplete dissociation these oxides for strong alkaline solutions
19Aqueous solutions OH- Solutions of salts When salts dissolve in water they become fully ionised. Sometimesthese ions can disturb the water equilibrium giving an acidicor alkaline solution. This interaction is called hydrolysis.Salts of a weak acid and a strong alkali gives an alkaline solution.E.g. soapsSodium ethanoate CH3COONa, not all the molecules ionise in water, someCH3COO- and Na+ ions associate with each other.H2O (l)H+ (aq)OH (aq)+CH3COONa(aq)CH3COO- (aq) + Na+ (aq)H2O (l)H+ (aq)OH - (aq)+OH-+CH3COONa(aq)CH3COO- (aq) + Na+ (aq)Some of the CH3COO- ions associate with H+ ions, so removing the H+.This causes more H2O to dissociate and so an excess of OH- ions results.
20Aqueous solutions2. Salts of a strong acid and a weak alkali gives an acid solution.Ammonium Chloride NH4Cl, in water not all the molecules ionise, someCl- and NH4+ ions associate with each other.H2O (l)H+ (aq)OH - (aq)+NH4Cl (aq)Cl- (aq) NH4+ (aq)H2O (l)H+ (aq)OH -+NH4Cl (aq)Cl- (aq) NH4+ (aq)H++NH4OH(aq)Some of the NH4+ ions associate with OH- ions, so removing them.(remember the position of the equilibrium for ammonium hydroxidesolution, it’s to the left.) This causes more H2O to dissociate and soan excess of H+ ions results.The H+ and Cl- ions have no tendency to join because HCl is a strong acid.
21Salts made from a strong acid and strong base Sodium ChlorideH2O (l)H+ (aq)OH - (aq)+NaCl (s)Na + (aq) Cl- (aq)Both the sodium ions and the chloride ions will not associatewith either the hydrogen or hydroxide ions because both HCl andNaOH, being a strong acid and a strong alkali are completely dissociateinto their respective ions.So the concentration of [H+] and [OH-] ions are unaffected.For this reason, sodium chloride solution has a pH of 7.