4.19 Buffer Solutions

Transcription

1 4.19 Buffer Solutions Buffer solution: p.319. (i.e. it minimizes the change in ph when A or B added) Buffers are made by: high conc of a weak acid + equal conc of its conj base Add base in salt form eg) CH 3 COOH + H 2 O CH 3 COO - + H 3 O + (1M) (1M) 1 M CH 3 COO - = 1 M Na CH 3 COO But Na is a Ka = [CH 3 COO - ] [H 3 O + ] = [H3O + ] = 1.8 x 10-5 [CH 3 COOH] therefore: ph = pka Recall: ph = - log [H 3 O + ] and pka = - log[ka] A solution of CH 3 COOH by itself is not a buffer A. Making a Buffer e.g.: HF as a buffer 1 L solution with 1.00 mole of HF is a weak monoprotic acid the ionization of HF occurs as follows: 1.00 M Very Very by adding 1.00 mole of NaF a is created. NaF dissociates completely and contributes more F - than HF does (Na + is spectator) Now F- is = This causes a shift to the left in the above And ph (less ) A WEAK Acid and Base don t each other like SA and SB They co-exist in equilibrium unless disturbed! 1

2 B. How a Buffer Works 1. What happens if we add a small amount of HCl to the above buffering system? By adding a small amount of HCl a little H 3 O + is produced [H 3 O + ] and ph But recall, H 3 O + is a on the above equilibrium And we have lots of F - equilib will shift and [H 3 O + ] will again The graph of H 3 O + versus time in a buffering system But not quite to its original value [H 3 O + ] Time HCl added 2. What happens if we add a small amount of NaOH to the buffer solution from above? a little OH - is produced [H 3 O + ] and ph BUT, have lots of HF equilib will shift and [H 3 O + ] will again But not quite to its original value 2

3 Draw ph vs. Time and and H 3 O + vs Time graphs: H 3 O + ph Time Time NaOH added NaOH added C. Two Kinds of Buffer Solutions 1. Weak Acid and the Salt of It s Conjugate Base eg. 1.0 M CH 3 COOH & 1.0 M NaCH 3 COO Useful as buffers in the acidic solutions (ph < 7 ) Any WA and a salt of it s conj.b will work as Acidic Buffer. 2. Weak Base and the Salt of It s Conjugate Acid eg. 1.0 M NH 3 & 1.0 M NH 4 Cl Useful as buffers in basic solutions (ph > 7) Any WB and a salt of it s conj.a will work as Basic Buffer. Question: Why can t strong acids and strong bases be used to prepare a buffer solution? D. Limitations of Buffers eg) add 1.5 moles of HCl to 1L of this buffer solution 1.00 M 1.00 M Very small [H 3 O + ] immediately increases to 1.5 M. Read Hebden: p not enough to react with the increased H 3 O + will have an of H 3 O + Do Questions # ph will crease significantly. we have overcome the limitations of our buffer..it cannot hold ph relatively constant 3

4 Example of buffer being overcome: 14 WA-SB Titration Buffer Region 8 ph here buffer is overcome by the base and is unable to maintain ph Volume of Base (ml) Buffers only work when amounts of acid or base are added to them! Acid Rain ph of rain always < 7 because CO 2 in air dissolves in rainwater: CO 2 + H 2 O H 3 O HCO 3 Rain with ph < 5.6 is called Sources of acidity in Acid Rain 1. Sulphur Most fuels (including coal & oil ) contain S Recall: Non-Metal Oxides react with water (rain) to form When fossil fuels are burned: S + O 2 SO 2 Then, in air: 2SO 2 + O 2 2SO 3 Both SO 2 and SO 3 react with rain: SO 2 + H 2 O H 2 SO 3 SO 3 + H 2 O H 2 SO 4 (Sulphurous acid) (Sulphuric acid) Mixtures of SO 2 and SO 3 react called SO x 4

5 2. Nitrogen Combusion (eg. car engines) cause N 2 and O 2 in the air to react N 2 + O 2 2NO..2NO + O 2 2NO 2 N 2 + 2O 2 2NO 2 2NO 2 + H 2 O HNO 2 + HNO 3 Mixtures of NO and NO 2 called NO x HNO 2, HNO 3, H 2 SO 3, H 2 SO 4 constitute the mixture in acid rain Protection Against Acid Rain Most lakes have moderate CO 2 and HCO 3 - systems. But, too much acid rain will overcome this buffer and life in/around the lake Limestone (CaCO 3 ) rich areas can help acid rain H 2 SO 4(aq) + CaCO 3(s) CaSO 4(s) + CO 2(aq) + H 2 O (l) Limestone powder has been dumped into lakes to reverse effects of acid rain. Environmental Problems with Acid Rain 1. Fish and plant growth 2. minerals are leached out of rocks a. can be moving poison substances from into groundwater b. can be removing nutrients out of topsoil 3. Metal and stone structures (bldgs and statues) damaged The big picture Acid rain produced in industrial area and falls on your house, or a forest - enormous cleanup costs People suffer from water contamination from: - the acid itself - subsequent leached poisons Read Hebden: p Do Questions # Food crops are damaged or destroyed 5

6 4.19 Buffer Solutions Buffer solution: a solution which resists changes in ph when a small amount of A or B is added. (i.e. it minimizes the change in ph when A or B added) Buffers are made by: high conc of a weak acid + equal conc of its conj base Add base in salt form eg) CH 3 COOH + H 2 O CH 3 COO - + H 3 O + (1M) (1M) 1 M CH 3 COO - = 1 M Na CH 3 COO But Na is a spectator ion Ka = [CH 3 COO - ] [H 3 O + ] = [H3O + ] = 1.8 x 10-5 [CH 3 COOH] therefore: ph = pka Recall: ph = - log [H 3 O + ] and pka = - log[ka] A solution of CH 3 COOH by itself is not a buffer A. Making a Buffer e.g.: HF as a buffer 1 L solution with 1.00 mole of HF is a weak monoprotic acid the ionization of HF occurs as follows: 1.00 M Very small Very small by adding 1.00 mole of NaF a buffer is created. NaF dissociates completely and contributes more F - than HF does (Na + is spectator) Now F- is = 1.00 This causes a shift to the left in the above And ph increases (less acidic) A WEAK Acid and Base don t neutralize each other like SA and SB They co-exist in equilibrium unless disturbed! 6

7 B. How a Buffer Works 1. What happens if we add a small amount of HCl to the above buffering system? By adding a small amount of HCl a little H 3 O + is produced [H 3 O + ] increases and ph decreases But recall, H 3 O + is a disturbance on the above equilibrium And we have lots of F - equilib will shift left and [H 3 O + ] will decrease again The graph of H 3 O + versus time in a buffering system But not quite to its original value [H 3 O + ] As equilibrium shifts left, [H 3 O + ] decreases to partially compensate for the sudden increase A very small net increase in the [H 3 O + ] Time HCl added 2. What happens if we add a small amount of NaOH to the buffer solution from above? a little OH - is produced [H 3 O + ] decreases and ph increases BUT, have lots of HF equilib will shift RIGHT and [H 3 O + ] will increase again But not quite to its original value 7

8 Draw ph vs. Time and H 3 O + vs Time graphs: [H 3 O + ] A very small net decrease in the [H 3 O + ] As equilibrium shifts right, [H 3 O + ] increases to partially compensate for the sudden decrease Time NaOH added Volunteer readers C. Two Kinds of Buffer Solutions 1. Weak Acid and the Salt of It s Conjugate Base eg. 1.0 M CH 3 COOH & 1.0 M NaCH 3 COO (acid is stronger than base) Useful as buffers in the acidic solutions (ph < 7 ) Any WA and a salt of it s conj.b will work as Acidic Buffer. 2. Weak Base and the Salt of It s Conjugate Acid eg. 1.0 M NH 3 & 1.0 M NH 4 Cl (base is stronger than acid) Useful as buffers in basic solutions (ph > 7) Any WB and a salt of it s conj.a will work as Basic Buffer. Question: Why can t strong acids and strong bases be used to prepare a buffer solution? SA s or SB s will not remain in an equilib mixture! Read Hebden: p D. Limitations of Buffers Do Questions # eg) add 1.5 moles of HCl to 1L of this buffer solution 1.00 M 1.00 M Very small Volunteer readers [H 3 O + ] immediately increases to 1.5 M. not enough F - to react with the increased H 3 O + will have an excess of H 3 O + ph will decrease significantly. we have overcome the limitations of our buffer..it cannot hold ph relatively constant 8

9 Example of buffer being overcome: 14 WA-SB Titration Buffer Region 8 ph 6 4 here buffer is overcome by the base and is unable to maintain ph Volume of Base (ml) Buffers only work when SMALL amounts of acid or base are added to them! E. Common Uses of Buffers Control of ph in industrial reactions Used in maintaining water quality Pools and hot tubs ph balanced shampoos and deodorants Soil ph Minimizing effects of acid rain Acid Rain ph of rain always < 7 because CO 2 in air dissolves in rainwater: CO 2 + H 2 O H 3 O HCO 3 Rain with ph < 5.6 is called ACID RAIN Sources of acidity in Acid Rain 1. Sulphur Most fuels (including coal & oil ) contain S Recall: Non-Metal Oxides react with water (rain) to form ACIDS When fossil fuels are burned: S + O 2 SO 2 Then, in air: 2SO 2 + O 2 2SO 3 Both SO 2 and SO 3 react with rain: SO 2 + H 2 O H 2 SO 3 SO 3 + H 2 O H 2 SO 4 (Sulphurous acid) (Sulphuric acid) Mixtures of SO 2 and SO 3 react called SO x 9

10 2. Nitrogen Combusion (eg. car engines) cause N 2 and O 2 in the air to react N 2 + O 2 2NO..2NO + O 2 2NO 2 N 2 + 2O 2 2NO 2 2NO 2 + H 2 O HNO 2 + HNO 3 Mixtures of NO and NO 2 called NO x HNO 2, HNO 3, H 2 SO 3, H 2 SO 4 constitute the mixture in acid rain Protection Against Acid Rain Most lakes have moderate CO 2 and HCO - 3 buffering systems. But, too much acid rain will overcome this buffer and kill life in/around the lake Limestone (CaCO 3 ) rich areas can help neutralize acid rain H 2 SO 4(aq) + CaCO 3(s) CaSO 4(s) + CO 2(aq) + H 2 O (l) Limestone powder has been dumped into lakes to reverse effects of acid rain. Environmental Problems with Acid Rain 1. Fish and plant growth 2. minerals are leached out of rocks a. can be moving poison substances from into groundwater b. can be removing nutrients out of topsoil 3. Metal and stone structures (bldgs and statues) damaged Volunteer readers The big picture Acid rain produced in industrial area and falls on your house, or a forest - enormous cleanup costs People suffer from water contamination from: - the acid itself - subsequent leached poisons Food crops are damaged or destroyed Read Hebden: p Do Questions #