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1. Introduction to Biochemistry4h 34m
2. Water3h 23m
3. Amino Acids8h 10m
4. Protein Structure10h 4m
5. Protein Techniques14h 5m
6. Enzymes and Enzyme Kinetics13h 38m
7. Enzyme Inhibition and Regulation 8h 42m
8. Protein Function 9h 41m
9. Carbohydrates7h 49m
10. Lipids5h 49m
11. Biological Membranes and Transport 6h 37m
12. Biosignaling9h 45m
Review 1: Nucleic Acids, Lipids, & Membranes2h 47m
Review 2: Biosignaling, Glycolysis, Gluconeogenesis, & PP-Pathway3h 12m
Review 3: Pyruvate & Fatty Acid Oxidation, Citric Acid Cycle, & Glycogen Metabolism2h 26m
Review 4: Amino Acid Oxidation, Oxidative Phosphorylation, & Photophosphorylation1h 48m

2. Water

Acid Dissociation Constant

2. Water

Acid Dissociation Constant: Videos & Practice Problems

Video Lessons Practice Worksheet

Topic summary

The acid dissociation constant (K_a) quantifies acid strength, with higher K_a values indicating stronger acids. Each acid can have multiple K_a values, especially polyprotic acids like phosphoric acid (H₃PO₄), which has three acidic hydrogens. To simplify calculations, K_a can be converted to pK_a using the equation pK_a = -log(K_a), where higher pK_a values correspond to weaker acids. Understanding these relationships is crucial for comparing acid strengths and performing related calculations.

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Acid Dissociation Constant (Ka)

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Acid Dissociation Constant (Ka) Video Summary

The acid dissociation constant, denoted as $ K_a $, is a crucial concept in understanding the strength of acids in solution. It quantifies the extent to which an acid donates hydrogen ions ($ H^+ $) to the solution, which directly relates to the pH of that solution. The $ K_a $ value is essentially the equilibrium constant for the dissociation of an acid, similar to other equilibrium constants previously discussed.

In a typical acid dissociation reaction, a conjugate acid ($ HA $) dissociates into a conjugate base ($ A^- $) and a hydrogen ion ($ H^+ $). The expression for the acid dissociation constant can be represented as:

\[K_a = \frac{[A^-][H^+]}{[HA]}\]

Here, the concentrations of the products ($ [A^-] $ and $ [H^+] $) are in the numerator, while the concentration of the reactant ($ [HA] $) is in the denominator. A higher $ K_a $ value indicates a stronger acid, as it reflects a greater tendency for the acid to dissociate and release hydrogen ions.

Some acids, known as polyprotic acids, can donate more than one hydrogen ion. Each acidic hydrogen corresponds to a separate $ K_a $ value. For example, phosphoric acid ($ H_3PO_4 $) has three acidic hydrogens, leading to three distinct $ K_a $ values, each representing the dissociation of one hydrogen ion. The first $ K_a $ value is typically the largest, indicating that the first hydrogen ion is the most easily dissociated, making it the strongest acidic hydrogen among the three.

Understanding these concepts allows for the comparison of different acids based on their $ K_a $ values, enabling the identification of the strongest and weakest acids in various chemical contexts. This foundational knowledge will be essential for further studies in acid-base chemistry.

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Acid Dissociation Constant (Ka)

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Acid Dissociation Constant (Ka) Video Summary

To calculate the acid dissociation constant (Ka) of uric acid, we start with the equilibrium concentrations of the species involved in the dissociation reaction. Uric acid (C₅H₄N₄O₃) dissociates into its conjugate base (A^-) and a hydrogen ion (H⁺). The dissociation reaction can be represented as:

C₅H₄N₄O₃ ⇌ A^- + H⁺

The Ka is defined as the ratio of the concentrations of the products to the reactants at equilibrium:

Ka = \frac{[A^-][H^+]}{[C_5H_4N_4O_3]}

In this case, we know the equilibrium concentration of the conjugate acid (C₅H₄N₄O₃) is 4.07 × 10^-3 M, and the concentration of the conjugate base (A^-) is 7.27 × 10^-4 M. To find the concentration of hydrogen ions (H⁺), we can use an ICE (Initial, Change, Equilibrium) table.

Initially, we assume the concentration of uric acid is X, while the concentrations of the products (A^- and H⁺) are both 0. As the reaction proceeds to equilibrium, the change in concentration for the conjugate base and hydrogen ion is equal to +7.27 × 10^-4 M, while the conjugate acid decreases by the same amount:

At equilibrium:

[C₅H₄N₄O₃] = 4.07 × 10^-3 M - 7.27 × 10^-4 M = 3.34 × 10^-3 M
[A^-] = 7.27 × 10^-4 M
[H⁺] = 7.27 × 10^-4 M

Now, substituting these values into the Ka expression gives:

Ka = \frac{(7.27 × 10^{-4})(7.27 × 10^{-4})}{3.34 × 10^{-3}} = \frac{5.29 × 10^{-7}}{3.34 × 10^{-3}} ≈ 1.58 × 10^{-4}

Thus, the calculated Ka for uric acid is approximately 1.3 × 10^-4, indicating the strength of the acid in solution. This value helps in understanding the acid's behavior in various chemical contexts and its reactivity in biological systems.

Problem

Calculate K_a of proprionic acid (CH ₃CH₂CO₂H) if [CH₃CH₂CO₂H]_eq= 0.2 M & [CH₃CH₂CO₂^-]_eq= 1.62 x 10^-3 M.

K_a = 1.3 x 10^-5

K_a = 7.8 x 10^-10

K_a = 3.9 x 10^-12

K_a = 5.1 x 10^-4

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Acid Dissociation Constant (Ka)

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Acid Dissociation Constant (Ka) Video Summary

The acid dissociation constant, denoted as $ K_a $, serves as a quantitative measure of an acid's strength. A higher $ K_a $ value indicates a stronger acid. However, $ K_a $ values can often be extremely large or small, making calculations cumbersome. To simplify this, we can express $ K_a $ values on a logarithmic scale using $ pK_a $ values. Like $ K_a $, $ pK_a $ also quantifies acid strength, but the relationship is inverse: a higher $ pK_a $ signifies a weaker acid.

The formula for calculating $ pK_a $ is:

$ pK_a = -\log(K_a) $

According to logarithmic rules, the negative log can also be expressed as the positive log of the reciprocal:

$ pK_a = \log\left(\frac{1}{K_a}\right) $

For example, consider acetic acid (a weak acid) and hydrochloric acid (a strong acid). The $ K_a $ value for acetic acid is $ 1.76 \times 10^{-5} $, which is a small number, while hydrochloric acid has a $ K_a $ of $ 1.3 \times 10^{6} $, a significantly larger value. To find the $ pK_a $ for acetic acid, we calculate:

$ pK_a = -\log(1.76 \times 10^{-5}) \approx 4.76 $

For hydrochloric acid, the calculation is:

$ pK_a = -\log(1.3 \times 10^{6}) \approx -6.3 $

This demonstrates that the $ pK_a $ of acetic acid (4.76) is greater than that of hydrochloric acid (-6.3), reinforcing the inverse relationship: as $ pK_a $ increases, acid strength decreases. Understanding these concepts will enhance your ability to analyze and compare the strengths of various acids effectively.

Problem

Which of the following is the strongest acid listed?

Lactic acid, Ka = 1.38 x 10^-4

Formic acid, pKa = 3.75

Acetic acid, Ka = 1.76 x 10^-5

Proprionic acid, pKa = 4.87

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Here’s what students ask on this topic:

What is the acid dissociation constant (Ka) and how is it used to measure acid strength?

The acid dissociation constant (Ka) is a quantitative measure of the strength of an acid in solution. It is defined as the equilibrium constant for the dissociation of an acid into its conjugate base and a hydrogen ion. The equation for Ka is:

$K a = \frac{[H_{+}] [A_{-}]}{[HA]}$

where [H⁺] is the concentration of hydrogen ions, [A^-] is the concentration of the conjugate base, and [HA] is the concentration of the acid. A higher Ka value indicates a stronger acid, as it implies a greater tendency to donate hydrogen ions.

How do you convert Ka to pKa and what is the significance of pKa values?

To convert Ka to pKa, you use the equation:

$p K a = - \log (K a)$

The pKa value is the negative logarithm of the Ka value. It provides a more convenient way to express acid strength, especially when Ka values are very large or very small. The relationship between pKa and acid strength is inverse: a lower pKa value indicates a stronger acid, while a higher pKa value indicates a weaker acid. This inverse relationship helps in comparing the strengths of different acids more easily.

What is the difference between Ka and pKa in terms of acid strength?

Ka and pKa are both measures of acid strength, but they are used differently. Ka is the acid dissociation constant and directly measures the equilibrium concentration of products and reactants in an acid dissociation reaction. A higher Ka value indicates a stronger acid. On the other hand, pKa is the negative logarithm of Ka:

$p K a = - \log (K a)$

pKa provides a more convenient scale for comparing acid strengths. A lower pKa value indicates a stronger acid, while a higher pKa value indicates a weaker acid. The inverse relationship between Ka and pKa helps simplify calculations and comparisons.

What are polyprotic acids and how do their Ka values differ?

Polyprotic acids are acids that contain more than one acidic hydrogen atom, capable of dissociating to form hydrogen ions. Each acidic hydrogen has its own dissociation constant, resulting in multiple Ka values for a single polyprotic acid. For example, phosphoric acid (H₃PO₄) has three acidic hydrogens, leading to three Ka values:

$K_{a1} = [H_{+}] [H_{2} {PO}_{4}] / [H_{3} {PO}_{4}]$

Each successive Ka value is typically smaller than the previous one, indicating that it becomes progressively harder to remove additional hydrogen ions. The first Ka value (Ka1) is usually the largest, indicating the strongest acidic hydrogen, while the subsequent Ka values (Ka2, Ka3) are smaller.

How do you determine the strength of an acid using Ka and pKa values?

The strength of an acid can be determined by comparing its Ka and pKa values. A higher Ka value indicates a stronger acid, as it shows a greater tendency to donate hydrogen ions. Conversely, a lower pKa value indicates a stronger acid, due to the inverse relationship between Ka and pKa:

$p K a = - \log (K a)$

For example, if you have two acids with Ka values of 1.0 × 10^-3 and 1.0 × 10^-5, the acid with the Ka value of 1.0 × 10^-3 is stronger. Similarly, if you convert these to pKa values (3 and 5, respectively), the acid with the pKa value of 3 is stronger.