Question

Adding a Strong Acid to a Buffer Learning Goal: To understand how buffers use reserves of conjugate acid and conjugate base to counteract the effects of acid or base addition on pH. A buffer is a mixture of a conjugate acid-base pair. In other words, it is a solution that contains a weak acid and its conjugate base, or a weak base and its conjugate acid. For example, an acetic acid buffer consists of acetic acid, $CH_3COOH$, and its conjugate base, the acetate ion $CH_3COO^-$. Because ions cannot simply be added to a solution, the conjugate base is added in a salt form (e.g., sodium acetate $NaCH_3COO$). Buffers work because the conjugate acid-base pair work together to neutralize the addition of $H^+$ or $OH^-$ ions. Thus, for example, if $H^+$ ions are added to the acetate buffer described above, they will be largely removed from solution by the reaction of $H^+$ with the conjugate base: $H^+ + CH_3COO^- ightarrow CH_3COOH$ Similarly, any added $OH^-$ ions will be neutralized by a reaction with the conjugate acid: $OH^- + CH_3COOH ightarrow CH_3COO^- + H_2O$ This buffer system is described by the Henderson-Hasselbaich equation $pH = pK_a + logfrac{[conjugate base]}{[conjugate acid]}$ Part A A beaker with 175 mL of an acetic acid buffer with a pH of 5.000 is sitting on a benchtop. The total molarity of acid and conjugate base in this buffer is 0.100 M. A student adds 6.00 mL of a 0.300 M HCl solution to the beaker. How much will the pH change? The $pK_a$ of acetic acid is 4.740. Express your answer numerically to two decimal places. Use a minus (-) sign if the pH has decreased. $Delta pH = $

Adding a Strong Acid to a Buffer
Learning Goal:
To understand how buffers use reserves of conjugate acid and conjugate base to counteract the effects of acid or base addition on pH.
A buffer is a mixture of a conjugate acid-base pair. In other words, it is a solution that contains a weak acid and its conjugate base, or a weak base and its conjugate acid. For example, an acetic acid buffer consists of acetic acid, $CH_3COOH$, and its conjugate base, the acetate ion $CH_3COO^-$. Because ions cannot simply be added to a solution, the conjugate base is added in a salt form (e.g., sodium acetate $NaCH_3COO$).
Buffers work because the conjugate acid-base pair work together to neutralize the addition of $H^+$ or $OH^-$ ions. Thus, for example, if $H^+$ ions are added to the acetate buffer described above, they will be largely removed from solution by the reaction of $H^+$ with the conjugate base:
$H^+ + CH_3COO^-
ightarrow CH_3COOH$
Similarly, any added $OH^-$ ions will be neutralized by a reaction with the conjugate acid:
$OH^- + CH_3COOH
ightarrow CH_3COO^- + H_2O$
This buffer system is described by the Henderson-Hasselbaich equation
$pH = pK_a + logfrac{[conjugate base]}{[conjugate acid]}$
Part A
A beaker with 175 mL of an acetic acid buffer with a pH of 5.000 is sitting on a benchtop. The total molarity of acid and conjugate base in this buffer is 0.100 M. A student adds 6.00 mL of a 0.300 M HCl solution to the beaker. How much will the pH change? The $pK_a$ of acetic acid is 4.740.
Express your answer numerically to two decimal places. Use a minus (-) sign if the pH has decreased.
$Delta pH = $

$Adding a Strong Acid to a Buffer Learning Goal: To understand how buffers use reserves of conjugate acid and conjugate base to counteract the effects of acid or base addition on pH. A buffer is a mixture of a conjugate acid-base pair. In other words, it is a solution that contains a weak acid and its conjugate base, or a weak base and its conjugate acid. For example, an acetic acid buffer consists of acetic acid, CH3COOH, and its conjugate base, the acetate ion CH3COO^-. Because ions cannot simply be added to a solution, the conjugate base is added in a salt form (e.g., sodium acetate NaCH3COO). Buffers work because the conjugate acid-base pair work together to neutralize the addition of H^+ or OH^- ions. Thus, for example, if H^+ ions are added to the acetate buffer described above, they will be largely removed from solution by the reaction of H^+ with the conjugate base: H^+ + CH3COO^- ightarrow CH3COOH Similarly, any added OH^- ions will be neutralized by a reaction with the conjugate acid: OH^- + CH3COOH ightarrow CH3COO^- + H2O This buffer system is described by the Henderson-Hasselbaich equation pH = pKa + logfrac[conjugate base][conjugate acid] Part A A beaker with 175 mL of an acetic acid buffer with a pH of 5.000 is sitting on a benchtop. The total molarity of acid and conjugate base in this buffer is 0.100 M. A student adds 6.00 mL of a 0.300 M HCl solution to the beaker. How much will the pH change? The pKa of acetic acid is 4.740. Express your answer numerically to two decimal places. Use a minus (-) sign if the pH has decreased. Delta pH =$

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