Buffers
57 flashcards covering Buffers for the MCAT Chem / Phys / Psych / Soc section.
Buffers are solutions that help maintain a stable pH level when small amounts of acid or base are added. They typically consist of a weak acid and its conjugate base, or a weak base and its conjugate acid, working together to resist changes in hydrogen ion concentration. This concept is crucial in biological and chemical systems, such as blood maintaining a pH around 7.4 to support enzyme function and overall homeostasis.
On the MCAT, buffers appear in the Chemical and Physical Foundations section, often through questions on calculating pH using the Henderson-Hasselbalch equation, determining buffer capacity, or analyzing how buffers function in real-world scenarios like cellular environments. Common traps include mistaking buffers for strong acids or bases, overlooking the importance of pKa values, or failing to account for dilution effects. Focus on mastering equilibrium principles, pH calculations, and the role of buffers in physiological contexts to tackle these problems effectively.
Practice sketching titration curves for buffers to visualize pH changes.
Terms (57)
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What is a buffer?
A buffer is a solution that resists changes in pH when small amounts of acid or base are added, typically consisting of a weak acid and its conjugate base or a weak base and its conjugate acid.
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Acidic buffer
An acidic buffer is a solution made from a weak acid and its conjugate base that maintains pH below 7 by neutralizing added base.
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Basic buffer
A basic buffer is a solution made from a weak base and its conjugate acid that maintains pH above 7 by neutralizing added acid.
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Henderson-Hasselbalch equation
The Henderson-Hasselbalch equation is pH = pKa + log([A-]/[HA]), used to calculate the pH of a buffer solution from the concentrations of the conjugate base and weak acid.
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pKa in buffers
pKa is the negative logarithm of the acid dissociation constant, indicating the pH at which a weak acid is half-dissociated, and it determines the effective pH range for a buffer.
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Buffer capacity
Buffer capacity is the amount of acid or base a buffer can neutralize before its pH changes significantly, depending on the concentrations of the buffer components.
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Factors affecting buffer capacity
Buffer capacity is influenced by the concentrations of the weak acid and its conjugate base, as higher concentrations allow for greater neutralization of added acids or bases.
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Optimal pH for buffers
A buffer is most effective within one pH unit of its pKa, where the concentrations of the acid and base forms are roughly equal, maximizing resistance to pH changes.
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Preparing a buffer
Preparing a buffer involves mixing a weak acid with its salt or a weak base with its salt in appropriate ratios to achieve the desired pH, often using the Henderson-Hasselbalch equation.
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Acetate buffer
An acetate buffer consists of acetic acid and sodium acetate, maintaining pH around 4.76, the pKa of acetic acid, and is commonly used in biochemical experiments.
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Phosphate buffer
A phosphate buffer uses the H2PO4- and HPO4 2- pair, effective around pH 7.2, and helps maintain cellular pH in biological systems.
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Bicarbonate buffer in blood
The bicarbonate buffer in blood involves carbonic acid and bicarbonate ions, regulating pH around 7.4 by responding to CO2 levels and maintaining acid-base balance.
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Effect of adding acid to a buffer
Adding a strong acid to a buffer converts some of the conjugate base to its weak acid form, minimizing pH change as long as the buffer components are not depleted.
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Effect of adding base to a buffer
Adding a strong base to a buffer converts some of the weak acid to its conjugate base form, resisting pH increase until the buffer is exhausted.
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Buffer range
The buffer range is the pH interval over which a buffer effectively resists changes, typically from pKa - 1 to pKa + 1, based on the acid's dissociation.
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Conjugate acid-base pair
In a buffer, a conjugate acid-base pair consists of a weak acid and its corresponding base, which together maintain equilibrium to stabilize pH.
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Calculating pH of a buffer
To calculate the pH of a buffer, use the Henderson-Hasselbalch equation with the known pKa and the ratio of conjugate base to weak acid concentrations.
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Weak acid in buffers
A weak acid in a buffer partially dissociates in water, providing the acid component that reacts with added base to prevent large pH shifts.
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Weak base in buffers
A weak base in a buffer partially accepts protons, providing the base component that reacts with added acid to stabilize pH.
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Le Chatelier's principle in buffers
Le Chatelier's principle explains how buffers maintain pH by shifting equilibrium; for example, added acid converts base to acid, restoring balance.
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Buffer exhaustion
Buffer exhaustion occurs when all the weak acid or conjugate base is consumed, after which the solution can no longer resist pH changes effectively.
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Dilution of a buffer
Diluting a buffer with water slightly changes its pH because the ratio of acid to base remains the same, but the capacity decreases due to lower concentrations.
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pH at equivalence point in buffers
In a titration involving a buffer, the pH at the equivalence point is not 7 if the acid and base are weak, depending instead on the hydrolysis of the salt formed.
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Common trap: Strong acid buffers
A common trap is assuming strong acids can act as buffers, but they cannot because they fully dissociate and cannot resist pH changes effectively.
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Strategy for buffer problems
For buffer problems, first identify the weak acid and conjugate base, then use the Henderson-Hasselbalch equation to calculate or predict pH changes.
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Protein buffers
Protein buffers, like hemoglobin, use ionizable side chains to maintain pH in biological fluids by acting as weak acids or bases.
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Amino acid buffers
Amino acids can act as buffers due to their carboxylic acid and amine groups, with pKa values around 2-10, helping regulate pH in cells.
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Quantitative buffer capacity
Quantitative buffer capacity is measured as the moles of acid or base needed to change the pH of one liter of buffer by one unit, indicating its effectiveness.
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Polyprotic acid buffers
Polyprotic acids, like phosphoric acid, can form buffers at multiple pKa values, allowing them to maintain pH over a broader range in systems like blood.
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Carbonic acid buffer system
The carbonic acid buffer system in blood involves CO2 hydration to form carbonic acid, which dissociates to bicarbonate, crucial for respiratory acid-base control.
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Example: pH of acetate buffer
For a 0.1 M acetic acid and 0.1 M sodium acetate buffer, the pH is 4.76, the pKa of acetic acid, since the acid and base concentrations are equal.
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Temperature effects on buffers
Temperature can affect buffers by altering pKa values, as higher temperatures may increase dissociation, slightly shifting the pH of the solution.
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Ionic strength in buffers
Ionic strength influences buffer pH by affecting the activity of ions, so buffers in high-salt environments may require pKa adjustments for accuracy.
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Half-equivalence point in buffers
At the half-equivalence point in a titration, the pH equals the pKa, as half the weak acid has been converted to its conjugate base, forming a buffer.
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Buffer in enzyme activity
Buffers maintain optimal pH for enzyme activity, as enzymes have specific pH ranges where they function best, preventing denaturation.
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TRIS buffer
TRIS buffer, with a pKa around 8.1, is used in biochemical assays to maintain neutral to slightly basic pH, common in protein electrophoresis.
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Citric acid buffer
A citric acid buffer uses the citric acid-phosphate system, effective around pH 3-6, and is found in the Krebs cycle for maintaining cellular pH.
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Logarithmic pH scale in buffers
The logarithmic pH scale means that in buffers, a tenfold change in the acid-to-base ratio results in a one-unit pH change, as per the Henderson-Hasselbalch equation.
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Derivation of Henderson-Hasselbalch
The Henderson-Hasselbalch equation derives from the acid dissociation equilibrium, rearranging Ka = [H+][A-]/[HA] to pH = pKa + log([A-]/[HA]).
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Equilibrium constant in buffers
In buffers, the equilibrium constant Ka determines the ratio of conjugate base to acid at a given pH, guiding how the buffer responds to additions.
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Common misconception: Buffer pH is fixed
A common misconception is that buffer pH never changes, but it does shift with large additions of acid or base beyond the buffer's capacity.
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Worked example: Adding HCl to buffer
If 0.01 moles of HCl is added to 1 liter of a buffer with 0.1 M acetic acid and 0.1 M acetate, the pH decreases slightly as acetate reacts with HCl, calculated via Henderson-Hasselbalch.
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Selecting a buffer for pH
To select a buffer, choose one with a pKa close to the desired pH, ensuring maximum capacity within one unit of that pKa.
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Ratio of acid to base in buffers
The ratio of acid to base in a buffer determines its pH, with equal ratios yielding pH equal to pKa, as per the Henderson-Hasselbalch equation.
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Buffer in physiological systems
In physiological systems, buffers like bicarbonate prevent drastic pH changes from metabolic acids, maintaining homeostasis in blood and tissues.
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Limitations of buffers
Buffers have limitations, such as failing with very large acid or base additions or in extreme pH conditions outside their effective range.
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pH calculation after addition
After adding a small amount of strong acid to a buffer, recalculate pH using the new concentrations in the Henderson-Hasselbalch equation.
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Isotonic buffers
Isotonic buffers maintain the same osmotic pressure as body fluids, preventing cell damage while stabilizing pH in medical and biological applications.
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Good's buffers
Good's buffers are synthetic buffers with pKa values between 6 and 8, designed for biological use with minimal interference in experiments.
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HEPES buffer
HEPES buffer, with a pKa of 7.5, is used in cell culture to maintain pH in the physiological range without penetrating cell membranes.
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Succinic acid buffer
A succinic acid buffer, effective around pH 4.2, is used in biochemical pathways like the Krebs cycle to stabilize intermediate pH.
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Error in buffer calculations
A common error in buffer calculations is forgetting to account for the initial pH or the exact amounts added, leading to inaccurate pH predictions.
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Buffers in titration
In titration, the buffer region occurs around the equivalence point for weak acid-strong base titrations, where pH changes gradually.
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Advanced: Multiple buffer systems
In complex systems like blood, multiple buffers work together, such as bicarbonate and phosphate, to provide overlapping pH protection.
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pH stability in buffers
pH stability in buffers relies on the rapid equilibrium between the acid and base forms, quickly counteracting perturbations.
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Example: Making a buffer
To make a pH 5.0 acetate buffer, mix acetic acid and sodium acetate so the ratio [acetate]/[acetic acid] equals 10^(5.0 - 4.76), using the Henderson-Hasselbalch equation.
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Common buffers in labs
Common lab buffers include phosphate and Tris, chosen for their pKa values that match experimental pH needs and chemical stability.