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ch17 MSJ jlm

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1 Additional Aqueous Additional Aqueous Equilibria Equilibria Chapter 17 Chapter 17 Lawrence J. Henderson 1878-1942. Discovered how acid-base equilibria are maintained in nature by carbonic acid/ bicarbonate system in the blood. Developed buffer equation. Karl A. Hasselbalch 1874-1962 Developed logarithmic form of buffer equation.

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2 Use this “decision tree” to calculate pH Use this “decision tree” to calculate pH values of solutions of specific solutions. values of solutions of specific solutions. 1. Is it pure water? If yes, pH = 7.00. 2. Is it a strong acid? If yes, pH = -log[HZ] 3. Is it a strong base? If yes, pOH = -log[MOH] or pOH = -log (2 x [M(OH) 2 ]) 1. Is it a weak acid? If yes, use the relationship K a = x 2 /(HZ – x), where x = [H +] 1. Is it a weak base? If yes, use the relationship K b = x 2 /(base – x), where x = [OH - ] 1. Is it a salt (MZ)? If yes, then decide if it is neutral, acid, or base; calculate its K value by the relationship K a K b = K w , where K a and K b are for a conjugate system; then treat it as a weak acid or base. 1. Is it a mixture of a weak acid and its weak conjugate base? It is a buffer; use the Buffer Equation.
3 Common Ion Effect So far, we’ve looked at solutions of weak acids and solutions of weak bases . Weak acid equilibrium: HX H + + X - (described by K a ) Weak base equilibrium: X - + H 2 O HX+ OH - (described by K b ) What if you had both HX and X - in the same solution? This could be obtained by adding some Na + X - salt to a solution of HX. The result is called a buffered solution. A buffered solution resists changes to pH when an acid or base is added.

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4 Buffered Solutions Buffered Solutions A buffer consists of a mixture of a weak acid (HX) and its conjugate base (X - ). An example of preparing a buffered solution: Add 0.10 mole of lactic acid (H-Z) and 0.14 mole of sodium lactate (Na-Z) to a liter of solution.
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