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Equilibrium 223
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the two ions The larger K value predominates If the larger value is Ka, the solution is acidic If the larger value is Kb, the solution is basic In the rare case where the two values are equal, the solution would be neutral The following table summarizes this information:
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CATION FROM ANION FROM SOLUTION
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Strong Base Strong Base Weak Base Weak Base
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Strong Acid Weak Acid Strong Acid Weak Acid
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Neutral Basic Acidic Must be determined by comparing K values
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For example, suppose you are asked to determine if a solution of sodium carbonate, Na2CO3, is acidic, basic, or neutral Sodium carbonate is the salt of a strong base (NaOH) and a weak acid (HCO3 ) Salts of strong bases and weak acids are basic salts As a basic salt, we know the final answer must be basic (pH above 7)
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Buffers
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Buffers are solutions that resist a change in pH when an acid or base is added to them The most common type of buffer is a mixture of a weak acid and its conjugate base The weak acid will neutralize any base added, and the weak base of the buffer will neutralize any acid added to the solution The hydronium ion concentration of a buffer can be calculated using an equation derived from the Ka expression: [H3O+ ] = K a [HA] [A ]
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Taking the negative log of both sides yields the Henderson Hasselbalch equation, which can be used to calculate the pH of a buffer: pH = p K a + log [A ] [HA]
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The weak base Kb expression can also be used giving: [OH ] = K b [B] [HB+ ] and pOH = p K b + log [B] [HB+ ]
These equations allow us to calculate the pH or pOH of the buffer solution knowing K of the weak acid or base and the concentrations of the conjugate weak acid and its conjugate base Also, if the desired pH is known, along with K, the ratio of base to acid can be calculated The more concentrated these species are, the more acid or base can be neutralized and the less the change in buffer pH This is a measure of the buffer capacity, the ability to resist a change in pH Let s calculate the pH of a buffer What is the pH of a solution containing 200 mol of ammonia and 300 mol of ammonium chloride in a volume of 100 L K b = 1 81 10 5 NH3 + H 2O There are two ways to solve this problem NH+ + OH 4
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[ NH+ ][OH ] (3 00 + x )(x ) 4 = = 1 81 10 5 Kb= [ NH3 ] ( 2 00 x ) Assume x small 1 81 10 5 = 3 00x 200 x = 1 21 10 5 pOH = 4918 pH = 14 00 0 - 4 918 = 9 082 [ NH+ ] 4 [ NH3 ]
Alternate solution: pOH = log 181 10 5 + log 300 200 = 4918 pH = 9082 = 4742 + log
Titration Equilibria
An acid base titration is a laboratory procedure commonly used to determine the concentration of an unknown solution A base solution of known concentration is added to an acid solution of unknown concentration (or vice versa) until an acid base indicator visually signals that the end point of the titration has been reached The equivalence point is the point at which a stoichiometric amount of the base has been added to the acid Both chemists and chemistry students hope that the equivalence point and the end point are close together If the acid being titrated is a weak acid, then there are equilibria which will be established and accounted for in the calculations Typically, a plot of pH of the weak acid solution being titrated versus the volume of the strong base added (the titrant) starts at a low pH and gradually rises until close to the equivalence point, where the curve rises dramatically After the equivalence point region, the curve returns to a gradual increase This is shown in Figure 153 In many cases, one may know the initial concentration of the weak acid, but may be interested in the pH changes during the titration To study the changes one can divide the titration curve into four distinctive areas in which the pH is calculated 1 Calculating the initial pH of the weak acid solution is accomplished by treating it as a simple weak acid solution of known concentration and Ka 2 As base is added, a mixture of weak acid and conjugate base is formed This is a buffer solution and can be treated as one in the calculations Determine the moles of acid consumed from the moles of titrant added that will be the moles of conjugate base formed Then calculate the molar concentration of weak acid and conjugate base, taking into consideration the volume of titrant added Finally, apply your buffer equations 3 At the equivalence point, all the weak acid has been converted to its conjugate base The conjugate base will react with water, so treat it as a weak base solution and calculate the [OH ] using Kb Finally, calculate the pH of the solution 4 After the equivalence point, you have primarily the excess strong base that will determine the pH
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