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# Henderson-Hasselbalch Equation

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Sections
Intro to Buffers
Henderson-Hasselbalch Equation
Intro to Acid-Base Titration Curves
Strong Titrate-Strong Titrant Curves
Weak Titrate-Strong Titrant Curves
Acid-Base Indicators
Titrations: Weak Acid-Strong Base
Titrations: Weak Base-Strong Acid
Titrations: Strong Acid-Strong Base
Titrations: Diprotic & Polyprotic Buffers
Solubility Product Constant: Ksp
Ksp: Common Ion Effect
Precipitation: Ksp vs Q
Selective Precipitation
Complex Ions: Formation Constant

Henderson-Hasselbalch Equation deals with conjugate acid-base pairs and allows us to calculate pH of a buffer without the use of ICE Chart.

Concept #1: Henderson-Hasselbalch Equation Example #1: Calculate the pH of a solution containing 2.0 M nitrous acid (HNO2) and 1.48 M lithium nitrite (LiNO2).

Ka = 4.6 x 10-4.

Practice: The Kb of C6H5NH2 (aniline) is 3.9 × 1010. Determine pH of a buffer solution made up of 500 mL of 1.4 M C6H5NH2 and 230 mL of 2.3 M C6H5NH3+.

Practice: Determine the buffer component concentration ratio (CB/WA) for a buffer with a pH of 4.7. Ka of boric acid (H3BO3) is 5.4 × 1010.

Practice: Calculate mass of NaN3 that needs be added to 1.8 L of 0.35 M HN3 in order to make a buffer with a pH of 6.5. Ka of hydrazoic acid is 1.9 × 105.

Concept #2: Calculating Buffer Range

Example #2: Determine the buffering range of a solution containing lactic acid (Ka = 1.4 x 10-4) and sodium lactate.

Practice: Which of the following weak acid-conjugate base combinations would result in an ideal buffer solution with a pH of 9.4?

a) formic acid (HCHO2) and sodium formate (Ka = 1.8 x 10-4)

b) benzoic acid (HC7H5O2) and potassium benzoate (Ka = 6.5 x 10-5)

c) hydrocyanic acid (HCN) and lithium cyanide (Ka = 4.9 x 10-10)

d) iodic acid (HIO3) and sodium iodate (Ka = 1.7 x 10-1)