1 Acids and Bases  Acids & Bases are one of the most important classes of chemicals  Acids and bases have been know to human for a long time  Acids taste sour (in fruit), change colour of certain dye  Bases taste bitter and feel slippery (like in soap, lime water)  Acids and bases are widely present in nature,  especially in plants, electrolyte balance in life system cycle etc  Acids and bases are widely used in industry for various purpose  Dissolving chemicals, e.g. HF, aqua regia (HCl:HNO 3 =3:1)  Reagents for producing various chemicals  Catalysing various types of reactions  Titration in volumetric analysis  etc Chemical Reactions

2 Acids and Bases - Definition  Classical definition Acids - Substances that, when dissolved in water, increase the concentration of H + ions e.g. HCl(g) H + (aq) + Cl - (aq) Note: H +, which is a proton only (no e - ), is actually bond with water molecule forming H 3 O +, the rxn is HCl(g) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) For simplicity, we often use H + instead of H 3 O +. Bases - Substance that, when dissolved in water, increase the concentration of OH - ions e.g.NaOH OH - (aq) + Na + (aq) NH 3 + H 2 O  NH OH -  Brønsted-Lowry definition Acid is proton donor and Base is proton acceptor (because H + is a proton and OH - of a base reacts with H + giving water) Chemical Reactions H2OH2O H2OH2O

3 Conjugate Acid and Base Pairs  An acid & a base always work together to transfer proton (donate-accept). A substance can function as an acid only if another substance behaves simultaneously as a base.  When an acid or a base is dissolved in water, ions are released - this process involves proton transfer. To mark the process and link the ions with its original acid or base, conjugate acid-base pairs are defined.  Acid and conjugate base always appear in pair; likewise base and conjugate acid appear in pair  When an acid losses proton (H+) it becomes the conjugate base of that acid (e.g. HX to X - ) when a base receives a proton (H+) it becomes the conjugate acid of that base (H 2 O to H 3 O + )  If an acid dissolves in water, H 2 O is a base; if a base dissolves in water, H 2 O becomes an acid. Chemical Reactions remove H + HX(aq) + H 2 O (l)  X - (aq) + H 3 O + (aq) acid base conjugate base conjugate acid add H + remove H + HCl(aq) + H 2 O(l)  Cl - (aq) + H 3 O + (aq) acid base conjugate conjugate base acid add H + NH 3 (aq) + H 2 O(l)  NH 4 + (aq) + OH - (aq) base acid conjugate conjugate acid base remove H +

4 Strengths of Acids and Bases  The strength of acids and bases  The strength of an acid is the ability to donate proton, or increase [H + ] when acid is dissolved in water.  likewise, the ability to accept proton, or [OH - ], determine the strength of a base  Common acids and their relative strengths  Strong acids, paired with bases with negligible basicity - Able to completely transfer their proton to water - Their conjugate bases are the weakest, with negligible tendency to accept proton  Weak acids, paired with week bases - These acids are partially dissociated to ions - Their conjugate bases are also weak, with limited ability of accepting proton  Acids with negligible acidity, paired with strong bases - These class of acids, though carrying H, give out no [H + ] - Their conjugate bases, however, are strong bases  Water can act as acid as well as base Chemical Reactions acidbase HClCl - H 2 SO 4 HSO 4 - HNO 3 NO 3 - H 3 OH 2 O HSO 4 SO 4 2- H 3 PO 4 H 2 PO 4 HFF - HC 2 H 3 O 2 C 2 H 3 O 2 - H 2 CO 3 HCO 3 - H 2 SHS - H 2 PO 4 HPO 4 2- NH 4 NH 3 HCO 3 CO 3 2- HPO 4 PO 4 3- H 2 OOH - OHO 2- H 2 H - CH 4 CH 3 - acid strength increase base strength increase negligibleweak strong

5 Acid and Base Equilibrium  The extent of ionisation of an acid or a base in water  Some acids (or bases) ionise in water completely, leaving no molecules behind  Other acids (or bases) ionise partially in water, forming an equilibrium between molecules and ions e.g. HF(aq) + H 2 O (l)  F - (aq) + H 3 O + (aq)(1) NH 3 (g) + H 2 O (l)  NH 4 + (aq) + OH - (aq)(2)  The tendency of ionisation of an acid (or a base) varies with the type of acids, we can use the concept of reaction equilibrium to indicate the degree of ionisation. The ‘equilibrium constant’ used to describe the degree of ionisation of an acid is called acid-dissociation constant, K a, which is defined as for equili. (1) for equili. (2) Chemical Reactions ions molecule The higher the K a value, the higher ion conc., the higher acidity/basicity

6 Quantifying the Strength of Acids and Bases  [H + ] and [OH - ] are the measure of the strengths of acids and bases  We know that an acid when dissolved in water releases [H + ] and a base gives [OH - ]  We also know that the strengths of an acid or a base depend on the [H + ] and [OH - ]  It comes naturally that [H + ] & [OH - ] are used to indicate the strengths of acids/bases  The range of [H + ] and [OH - ]  Dilute aqueous solutions at 25°C always give, K w =[H + ][OH - ]=1.0x  For an acid [H + ]>[OH - ], K w =[H + ][OH - ]=1.0x  For a base [OH - ]>[H + ], K w =[H + ][OH - ]=1.0x  For pure water, which is neutral [H + ]=[OH - ]=1.0x10 -7, K w =[H + ][OH - ]=1.0x  pH scale  For convenience the low value of [H + ] and [OH - ], we use the scale of log 10 [H + ] define pH= -log 10 [H + ] Scale: Acid pH=0-7 [H + ]>[OH - ]; strong acids have low pH Base pH=7-14 [OH - ]>[H + ]; strong bases have high pH Note: When using [OH - ] (which is less used), we have pOH= -log 10 [OH - ] (=14-pH) Chemical Reactions water can act as an acid as well as a base at equilibrium H 2 O  H + + OH - Equili. constant at 25 °C is found to be Further examine other aqueous solution the same relation holds In pure water [H 2 O] is constant Known [H + ], [OH - ] can be calculated by this eqn.

7 Calculation of pH Example 1: Calculate pH of 0.05M HNO 3 solution HNO 3 + H 2 O  H 3 O + + NO 3 - HNO 3 is a strong acid, HNO 3 ionizes completely in water, i.e. [H 3 O + ]= 0.05M  pH = - log 10 [0.05] = 1.3 Example 2 : Calculate pH and pOH of 0.05M NaOH solution NaOH + H 2 O  Na + + OH - NaOH is a strong base, NaOH ionizes completely in water, i.e. [OH - ]=0.05M, K w = [H 3 O + ][OH - ] = 1 x M 2  [H 3 O + ] = 1 x M 2 / [OH - ] = 1 x M 2 / 0.05 M = 2 x M  pH = - log 10 [2 x ] = 12.7  pOH = 14 - pH = = 1.3 (why is this result the same as that of example1?) Chemical Reactions

8 Calculation of pH Example 3: What is the [OH - ], in mol/L, in a solution whose pH is 9.72? Known: pH = - log 10 [H 3 O + ] = 9.72  [H 3 O + ] = 1.9 x (mol/L) for any aqueous solution K w = [H 3 O + ][OH - ] = 1.0 x (mol/L) 2  [OH - ] = K w / [H 3 O + ] = 11.0 x (mol/L) 2 / 1.9 x (mol/L) = 5.3 x (mol / L) Example 4: The acid-dissociation constant, K a, of hydrofluoric acid is 6.8x What is the [H 3 O + ] in a 2M HF solution? What is the pH of the solution? HF(aq) + H 2 O (l)  F - (aq) + H 3 O + (aq) initial 200 at equili. 2 - xxx By definition Solve the eqn for x ( = [H 3 O + ]) pH = - log 10 [H 3 O + ] = - log 10 (0.0365) = 1.44 Chemical Reactions

9 Aqueous Equilibria and Some Applications  In chemistry many aqueous systems involve equilibria  Human body fluids are in electrolyte equilibria in order to function properly  Electrolyte: aqueous solutions that contain ions  Plants contain weak acids, which maintain right balance for plants to grow  Many properties of a solution that has ions are affected by its equilibrium state.  etc. (In a broad sense, harmony=balance=equilibria)  Many phenomena in chemistry can be studied by means of equilibria. We will look at:  The behaviour of an equilibrated electrolyte solution when other ions are added Applications  Buffer effect  Acid-base titration  Solubility of ionic substances and the factors affecting it Chemical Equilibria

10 The Common-Ion Effect from Equilibrium  Considering the following two cases Case 1. What is the pH of 0.3M acetic acid HC 2 H 3 O 2 solution, (K a =1.8x10 -5 )? HC 2 H 3 O 2 (aq)  H + (aq) + C 2 H 3 O 2 - (aq) initial at equilibrium 0.3-xxx By definition Solve eqn for x Chemical Equilibria Note: HC 2 H 3 O 2 is a weak acid (K a <<1) and the water solut n of HC 2 H 3 O 2 is an electrolyte solution.

11 The Common-Ion Effect from Equilibrium  (cont’d) Case 2. What is the pH of solut n contains 0.3M acetic acid HC 2 H 3 O 2 & 0.3M NaC 2 H 3 O 2 ? HC 2 H 3 O 2 (aq)  H + (aq) + C 2 H 3 O 2 - (aq) initial at equilibrium 0.3-xx0.3+x By definition Solve equ for x  Compare cases 1 & 2 : The extent of ionisation of HC 2 H 3 O 2 is reduced by the presence of NaC 2 H 3 O 2 (which has C 2 H 3 O 2 - ion in common with HC 2 H 3 O 2 )  This is called the Common-ion Effect. It works in many equilibrated electrolyte solutions such as buffer solutions, solubility of ionic compounds etc. Chemical Equilibria Note: NaC 2 H 3 O 2 ionises in water completely NaC 2 H 3 O 2 (aq)  Na + (aq) + C 2 H 3 O 2 - aq) Note: The presence of NaC 2 H 3 O 2 & Na + does not change K a value Note: C 2 H 3 O 2 - is the conjugate base of HC 2 H 3 O 2

12 Buffered Solutions  Behaviour of a solution containing a weak conjugate acid-base pair equilibrium of weak acidHX(aq)  H + (aq) + X - (aq) acid-dissociation constant If a base, OH -, is added, OH - (aq) + HX(aq)  H 2 O(aq) + X - (aq)  [HX]  & [X - ]  If an acid, H +, is added, H + (aq) + X - (aq)  HX(aq)  [X - ]  & [HX]  When the addition of OH - or H + is small compared to [HX] & [X - ], the change to [HX] & [X - ] is very small, so does the ratio [HX] / [X - ]  the [H + ] thus pH will remain almost constant.  A Buffered Solution (also called Buffer) contains a weak conjugate acid-base pair. It can resist drastic change of pH upon the adding strong acid or base.  Buffers solutions are widely used in biology and biochemistry because of the need of maintaining certain pH for some reactions/process to occur properly. Note: Buffer solutions can be made for all pH ranges. The amount of acid or base it can neutralise before pH begins to change (called buffer capacity) depends on the [HX] & [X - ]. Chemical Equilibria As the HC 2 H 3 O 2 and C 2 H 3 O 2 - pair in case 2