The Henderson Hasselbalch Equation is the equation that can be used to determine the pH of a buffer solution by relating pKa and the molar concentrations of a weak acid and its conjugate base.

**What is a Buffer? **

A buffer is a solution comprised of a weak acid (HA) and its conjugate base (A^{–}). An example of a buffer is hypochlorous acid, HClO, and its conjugate base of sodium hypochlorite, NaClO.

A buffer can resist significant changes in pH by keeping H^{+} and OH^{–} ions constant. How is this possible?

When a strong acid is added the conjugate base of the buffer will neutralize it.

When a strong base is added the weak acid of the buffer will neutralize it.

As long as the weak acid and conjugate base components of the buffer are greater in amount than the strong acid or strong base then the buffer will be maintained.

**Deriving the Henderson Hasselbalch Equation**

When a weak acid dissociates in water an equilibrium is established:

From this equation we can set up our equilibrium expression with the use of our acid dissociation constant, K_{a}. Remember that we ignore solids and liquids.

By rearranging our equilibrium expression we obtain the following format:

By taking the –log of both sides we obtain:

By inverting the ratio HA and A^{–} we finally obtain the Henderson Hasselbalch Equation:

The values of HA and A^{–} within the Henderson Hasselbalch Equation can either be in moles or molarity.

**Using the Henderson Hasselbalch Equation**

The Henderson Hasselbalch Equation is a convenient way to calculate the pH of a buffer solution.

**PRACTICE 1: **Calculate the pH of a solution that is 0.27 M in HF and 0.11 M in NaF. The acid dissociation constant of HF is 3.5 x 10^{-4}.

**STEP 1:** Identify the weak acid and its conjugate base.

Since the K_{a} value of HF is less than 1 it represents the weak acid and because NaF possesses 1 less hydrogen than HF it represents the conjugate base.

**STEP 2: **Take the –log of K_{a} to determine the pKa value.

**STEP 3:** Plug in your given and calculated values into the Henderson Hasselbalch Equation.

Now what happens when we are given the volume and molarity of the buffer components?

**PRACTICE 2: **Calculate the pH of a buffer solution that is 25.0 mL of 0.60 M in NH_{4}^{+} and 30.0 mL of 0.75 M in NH_{3}. (*K*_{b} = 1.8 × 10^{−5 }for NH_{3})

**STEP 1:** Identify the weak acid and its conjugate base.

Another way of looking at buffer is it being composed of a conjugate acid and its weak base. Since the K_{b} value of NH_{3} is less than 1 it represents the weak base and because NH_{4}^{+} possesses 1 more hydrogen than NH_{3} it represents the conjugate acid.

**STEP 2: **Change K_{b} into K_{a}.

**STEP 3:** Take the –log of K_{a} to determine the pKa value.

**STEP 4: **When given both the volume and molarity of the buffer components we must calculate their moles.

**STEP 5:** Plug in your given and calculated values into the Henderson Hasselbalch Equation.

The Henderson Hasselbalch Equation is a useful tool in determining the pH of a buffer solution without the use of an ICE Chart. This equation will also play a role in acid-base identification, acid-base titrations and titration curves.

Calculate the pOH of a solution that results from mixing 16 mL of 0.13 M methylamine (CH3NH2) with 32 mL of 0.1M CH3NH3CI. The Kb value for CH3NH2 is 3.6 x 10-4.

Enter your answer in the provided box. What is the component concentration ratio, [CH3COO-]/[CH3COOH], of a buffer that has a pH of 4.09? (Ka of CH3COOH = 1.8 x 10-5)

Enter your answer in the provided box. Find the pH of a buffer that consists of 0.39 M boric acid (H3BO3) and 0.41 M sodium borate (NaH2BO3). (pKa of boric acid = 9.24)

You may want to reference (pages 754-765) section 16.2 while completing this problem. Calculate the pH of the solution that results from each of the following mixtures. Part A160.0 mL of 0.27 M HF with 225.0 mL of 0.32 M NaF.Express your answer using two decimal places.

How many grams of CaCl2 are needed to make 495.3 g of a solution that is 30.5% (m/m) calcium chloride in water? Note that mass in not technically the same thing as weight, but % (m/m) has the same meaning as % (w|w). How many grams of water are needed to make this solution?

Phosphoric acid is a triprotic acid (Ka1 = 6.9 x 10-3, Ka2 = 6.2 x 10-8, and Ka3 = 4.8 x 10-13). To find the pH of a buffer composed of H2PO4- (aq) and HPO42- (ag), which pKa value would you use in the Henderson-Hasselbalch equation? (i) pKa1 = 2.16 (ii) pKa2 = 7.21 (iii) pKa3= 12.32 Calculate the pH of a buffer solution obtained by dissolving 11.0 g of KH2PO4 (s) and 28.0 g of Na2HPO4 (s) in water and then diluting to 1.00 L.

Be sure to answer all parts.What is the [H3O+] and the pH of a buffer that consists of 0.14 M HF and 0.38 M KF? (Ka of HF = 6.8 x 10-4) [H3O+] = _______ x 10 _______ pH = _________

What is the pH of a solution that consists of 0.50 M H2C6H6O6 (ascorbic acid) and 0.75 M NaHC6H6O6 (sodium ascorbate)? a. 3.76 b. 4.34 c. 3.99 d. 4.57

100.0 mL of 0.0500 M CH3COOH(aq) is mixed with 100.0 mL of 0.500 M NaCH3COO (aq). Ka = 1.75 x 10-5 at 25°C The pH of the resulting mixture is (i) 4.92 (ii) 4.66 (iii) 10.24 (iv) 5.76 (v) 3.79

Write the Henderson-Hasselbalch equation for a propanoic acid solution (CH3CH2CO2H, pKa = 4.874) using the symbols HA and A-, and the given pKa value for propanoic acid in the expression. Using the equation above, calculate the quotient [A-]/[HA] at pH (a) 4.556 (b) 5.454 (c) pH 4.874

Be sure to answer all parts.What is the [H3O+] and the pH of a buffer that consists of 0.38 M HNO2 and 0.65 M KNO2? (Ka of HNO2 = 7.1 x 10-4)

Determine the pH and pOH of 0.250 L of a buffer that is 0.0215 M boric acid and 0.0400 M sodium borate: pKa for B(OH)3 = 9.00 at 25°C. pH = pOH =

What is the pH of a solution which is 0.021 M in weak base and 0.035 M in the conjugate weak acid (Ka = 7.1 x 10-6)?

Enter your answer in the provided box. What is the component concentration ratio, [NO2-]/[HNO2], of a buffer that has a pH of 3.99? (Ka of HNO2 = 7.1 x 10-4)

A certain indicator, HA, has a Ka value of 0.000016. The protonated form of the indicator is blue and the ionized form is red. What is the pKa of the indicator? pKa = What is the color of this indicator in a solution with pH = 7? (i) blue (ii) red (iii) purple

A buffer solution is made using a weak acid, HA, with a pKa of 7.10. If the ratio of A- to HA is 1.0 x 102, what is the pH of the buffer?

How many moles of solid NaF would have to be added to 1.0 L of 1.65 M HF solution to achieve a buffer of pH 3.35? Assume there is no volume change. (Ka for HF = 7.2 x 10-4) (a) 2.7 (b) 0.34 (c) 0.98 (d) 1.0 (e) 1.6

Enter your answer in the provided box.Find the pH of a buffer that consists of 0.36 M HBrO and 0.58 M KBrO (pKa of HBrO = 8.64).

Enter your answer in the provided box. What is the component concentration ratio, [Pr-]/[HPr], of a buffer that has a pH of 4.60? (Ka of HPr = 1.3 x 10-5) Report your answer to 3 significant figures.

Part AWhat mass of ammonium chloride should be added to 2.50 L of a 0.150 M NH3 in order to obtain a buffer with a pH of 9.50? Express your answer using two significant figures.

If a buffer solution is 0.430 M in a weak acid (Ka = 6.0 x 10-6) and 0.220 M in its conjugate base, what is the pH?

Determine the pH of the following solutions. (a) a 0.54 M CH3COOH solution (b) a solution that is 0.54 M CH3COOH and 0.24 M CH3COOHNa.

How many grams of glyoxylic acid and sodium glyoxylate are needed to prepare 2.25 L of a 2.40 M buffer at pH 3.50? The pKa of glyoxylic acid is 3.34.