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11 Buffers

The document discusses buffers and how they maintain pH levels in the blood. It specifically focuses on the carbonic acid-bicarbonate buffer system, which is the most important buffer. This buffer system involves the equilibrium between carbonic acid, which donates hydrogen ions, and bicarbonate ions, which accept hydrogen ions. This helps regulate pH levels and protects against changes that could be harmful or fatal.

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Buffers
Buffers are solutions that resist changes in pH on the addition of small amounts of acids or bases A buffer consists of either A weak acid and the salt of the acid or A weak base and the salt of the base Examples Ethanoic acid and sodium ethanoate Ammonia and ammonium chloride
How does a buffer work? HA  A - + H + Weak acid MA  M + + A - In the mixture there is a relatively high concentration of undissociated acid (HA) and fully dissociated salt (MA)‏
When acid is added: H + combines with A - to give undissociated acid (HA)‏ H + + A -  HA It does this because HA is a weak acid so nearly all the H + will be removed.  the pH only changes slightly
When OH - is added OH - will combine with H + to give water OH - + H +  H 2 O The equilibrium is disturbed and more HA dissociates to form H + and A - There is a continual supply of H + to neutralise the OH - the pH only changes slightly
Essentially a buffer works because A high [A - ] traps added H + A high [HA] supplies H + to trap OH -
Finding the pH of buffer solutions pH = pK a + log [salt] [acid] A solution was made containing propionic acid at a conc. of 0.10mol/L and sodium propanoate at a conc. of 0.10 mol/L. Find the pH of the solution K a (propionic acid) = 1.34 x 10 -5 mol/L pKa = -log K a = -log 1.34 x 10 -5 mol/L = 4.87 pH = 4.87 = log 0.10 = 4.87 + log 1.0 = 4.87 0.10
A solution was made containing propionic acid at a conc. of 0.10mol/L and sodium propanoate at a conc. of 0.20 mol/L. Find the pH of the solution K a (propionic acid) = 1.34 x 10 -5 mol/L pKa = -log K a = -log 1.34 x 10 -5 mol/L = 4.87 pH = 4.87 = log 0.20 = 4.87 + log 2.0 = 5.17 0.10
Adding H + or OH - to a buffer solution Calculate the effect of adding 10ml of hydrochloric acid of concentration 1.0 mol/L to 1.0L of a buffer containing 0.10mol/L ethanoic acid and 0.10mol/L sodium ethanoate. pK a of ethanoic acid = 4.75 Step 1 Find the pH of the buffer solution Using the equation pH = pK a + log [salt] [acid] pH = 4.75 + log 0.10 = 4.75 0.10
Step 2 Work out the moles of hydrochloric acid added Amount of hydrochloric acid added = 10ml of 1.0mol/L = 0.010mol Step 3 Find the resulting pH The added H + combines with ethanoate ions to form undissociated ethanoic acid. The concentration of the undissociated acid increases by the amount of H + added and the concentration of the salt decreases by the same amount. Using the equation for the pH of a buffer find the pH
PH = 4.75 + log 0.10 – 0.010 = 4.66 0.10 + 0.010 The pH had decreased by 0.09 For the addition of OH - ions the OH combines with H + ions to form water so the undissociated acid concentration will decrease and the salt concentration will increase. To the same buffer as above calculate the resulting pH when 10ml of sodium hydroxide of concentration 1.0mol/L is added. PH = 4.84 the pH increases by 0.09
http://www.chemistry.wustl.edu/~edudev/LabTutorials/Buffer/Buffer.html During exercise, the muscles use up oxygen as they convert chemical energy in glucose to mechanical energy. This O 2 comes from hemoglobin in the blood. CO 2 and H + are produced during the breakdown of glucose, and are removed from the muscle via the blood. The production and removal of CO 2 and H + , together with the use and transport of O 2 , cause chemical changes in the blood.
These chemical changes, unless offset by other physiological functions, cause the pH of the blood to drop. If the pH of the body gets too low (below 7.4), a condition known as acidosis results. This can be very serious, because many of the chemical reactions that occur in the body, especially those involving proteins, are pH-dependent. Ideally, the pH of the blood should be maintained at 7.4. If the pH drops below 6.8 or rises above 7.8, death may occur. Fortunately, we have buffers in the blood to protect against large changes in pH.
The Carbonic-Acid-Bicarbonate Buffer in the Blood By far the most important buffer for maintaining acid-base balance in the blood is the carbonic-acid-bicarbonate buffer. Carbonic acid (H 2 CO 3 ) is a weak acid. CO 2( g )  CO 2( aq ) + H 2 O ( l )  H 2 CO 3( aq )  H + ( aq ) + HCO 3 - ( aq )
In blood plasma, the carbonic acid and hydrogen carbonate ion equilibrium buffers the pH. In this buffer, carbonic acid (H 2 CO 3 ) is the hydrogen-ion donor (acid) and hydrogen carbonate ion (HCO 3 - ) is the hydrogen-ion acceptor (base). H 2 CO 3(aq)  H + (aq) + HCO 3 - (aq)
Additional H + is consumed by HCO 3 - and additional OH - is consumed by H 2 CO 3 Carbonic acid concentration is controlled by respiration, that is through the lungs. Carbonic acid is in equilibrium with dissolved carbon dioxide gas. H 2 CO 3 (aq)                      CO 2 (aq) + H 2 O(l)