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Questions and Answers
Which of the following best describes a buffer solution?
Which of the following best describes a buffer solution?
- A solution of a strong acid and a strong base.
- A solution of a weak base and its conjugate acid.
- A solution of a weak acid and its conjugate base (or a weak base and its conjugate acid), both at reasonable concentration. (correct)
- A solution of a strong acid and its conjugate base.
The pH of a buffer solution is significantly affected by reasonable dilution.
The pH of a buffer solution is significantly affected by reasonable dilution.
False (B)
What happens to the pH of blood during hyperventilation, and why?
What happens to the pH of blood during hyperventilation, and why?
The pH increases (alkalosis) due to a decrease in $H_2CO_3$ relative to $HCO_3^-$.
The Henderson-Hasselbalch equation states that $pH = pKa + log(\frac{[A^-]}{[HA]})$. When $[A^-] = [HA]$, the pH is equal to the ______.
The Henderson-Hasselbalch equation states that $pH = pKa + log(\frac{[A^-]}{[HA]})$. When $[A^-] = [HA]$, the pH is equal to the ______.
Match the following scenarios with their effect on a buffer solution containing acetic acid ($CH_3COOH$) and sodium acetate ($CH_3COONa$):
Match the following scenarios with their effect on a buffer solution containing acetic acid ($CH_3COOH$) and sodium acetate ($CH_3COONa$):
Which expression correctly calculates the pH of a buffer solution, according to the Henderson-Hasselbalch equation?
Which expression correctly calculates the pH of a buffer solution, according to the Henderson-Hasselbalch equation?
Adding a strong acid to a buffer solution containing acetic acid and sodium acetate will increase the concentration of acetate ions.
Adding a strong acid to a buffer solution containing acetic acid and sodium acetate will increase the concentration of acetate ions.
What determines the buffering capacity of a buffer solution?
What determines the buffering capacity of a buffer solution?
To prepare a buffer solution using a weak acid and its conjugate base, one can add less than 1 mole equivalent of ______ to a solution of the weak acid.
To prepare a buffer solution using a weak acid and its conjugate base, one can add less than 1 mole equivalent of ______ to a solution of the weak acid.
Match the following components of blood with their role in maintaining pH:
Match the following components of blood with their role in maintaining pH:
What is the primary mechanism by which rapid breathing (hyperventilation) affects the pH of blood?
What is the primary mechanism by which rapid breathing (hyperventilation) affects the pH of blood?
Addition of NaOH to a solution of $CH_3COOH$ will always result in a buffer solution.
Addition of NaOH to a solution of $CH_3COOH$ will always result in a buffer solution.
If a buffer solution contains equal concentrations of a weak acid and its conjugate base, what is the pH of the solution relative to the pKa of the acid?
If a buffer solution contains equal concentrations of a weak acid and its conjugate base, what is the pH of the solution relative to the pKa of the acid?
When a buffer's capacity is exhausted, further addition of acid or base will cause a ______ change in pH.
When a buffer's capacity is exhausted, further addition of acid or base will cause a ______ change in pH.
Which of the following actions will likely lead to a decrease in the pH of a buffer solution containing $NH_3$ and $NH_4^+$?
Which of the following actions will likely lead to a decrease in the pH of a buffer solution containing $NH_3$ and $NH_4^+$?
Flashcards
Buffer solution
Buffer solution
A solution of a weak acid and its conjugate base (or a weak base and its conjugate acid), both at reasonable concentrations, which will maintain a constant pH on addition of significant amounts of H3O+ or OH- ions.
Simple buffer
Simple buffer
Buffers contain acetic acid and sodium acetate in roughly equal concentrations.
pH control
pH control
The pH of a buffer solution is controlled by the ratio of the concentrations of the conjugate base to the weak acid.
Henderson-Hasselbalch equation
Henderson-Hasselbalch equation
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Equal buffer concentrations
Equal buffer concentrations
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Making a buffer
Making a buffer
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Buffer capacity
Buffer capacity
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PH of the blood increases
PH of the blood increases
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Study Notes
Buffer Solutions
- Blood pH is approximately 7.4; a shift to 7 or 8 could be fatal, making maintenance by a buffer system essential.
- A buffer solution comprises a weak acid and its conjugate base (or vice versa), both at reasonable concentrations.
- Buffer solutions maintain a stable pH despite the addition of acids or bases or dilution
Simple Buffers
- Simple buffer solutions can be created with acetic acid and sodium acetate in roughly equal concentrations
- The solution will contain CH3COOH + H3O+ (from the acid) and CH3COO- and Na+ (from the salt)
- A strong acid is consumed by conjugate base: CH3COO-(aq) + H3O+(aq) → CH3COOH(aq) + H2O(l)
- A strong base is consumed by acid: CH3COOH(aq) + OH-(aq) → CH3COO-(aq) + H2O(I)
- Reactions go to completion
Henderson-Hasselbalch Equation
- The Henderson-Hasselbalch equation is a method to calculate pH in buffer solutions: pH = pKa + log([A-]/[HA]).
- pH of a buffer solution is controlled by the ratio of [A-] / [HA].
- When [A-] / [HA] = 1, then pH = pKa because log(1) = 0.
- At pH = pKa, the concentrations of A- and HA are equal, making the buffer equally effective against added acid or base.
Expressing the Henderson-Hasselbalch Equation
- The Henderson-Hasselbalch equation can be expressed in terms of numbers of moles, rather than concentrations: pH = pKa + log(nA/nHA).
Preparation of Buffers
- Buffers can be formed by adding less than 1 mole equivalent of NaOH(aq) to a solution of CH3COOH(aq) where CH3COOH(aq) + NaOH(aq) → CH3COO-(aq) +Na+(aq) + H2O
- A weak base can also be used by adding HCl(aq) to a solution of NH3(aq) where NH3(aq) + HCl(aq) → NH4+(aq) + Cl-(aq)
Buffer Capacity
- Buffers can only absorb so much acid or base before they are exhausted
- Once capacity is said to be exhausted, the pH begins to change significantly
- After adding 50.0 mL of NaOH solution to a solution, will only have CH3COOÖ¾ and be an ineffective buffer solution
Examples of Natural Buffers
- The pH of blood is kept constant at 7.4 via the following equilibria
- Carbonate: H2CO3 + H2O = H3O+ + HCO3-
- Hemoglobin and oxyhemoglobin: HHb + H2O = H3O+ + Hb- or HHbO2 + H2O = H3O+ + HbO2-
- Phosphate: H2PO4 + H2O = HPO42- + H3O+
Hyperventilation
- Rapid breathing (hyperventilation) expels too much CO2, causing lightheadedness
- pH of the blood increases by increasing the ratio [HCO3-] /[H2CO3]
- Increased pH can be cured by breathing into a paper bag (rebreathing CO2.)
- Rebreathing decreases the [HCO3] / [H2CO3] ratio, lowering the pH to a normal value over time
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