🎯 By the end of this guide,

You’ll master calculating the overall equilibrium constant, K, for an acid-base reaction using pK_a values. We’ll guide you through each step of the process, from identifying the acids and bases in the reaction to using pK_a values effectively. This will equip you with the tools to assess which side of a reaction is favored confidently. Whether you’re gearing up for exams or just keen to deepen your understanding of acid-base chemistry, this guide is designed for you.

🖐️ Before we begin,

Make sure you’re familiar with:

• Brønsted-Lowry acids and bases

Challenge 1

Question:

Calculate the value of the overall equilibrium constant K for this acid-base reaction

Here’s how we’ll do it

Step 1: Identify the acid on either side of the equilibrium

💡 Hint: You’ll often deal with Brønsted-Lowry acids and bases in these types of questions. Remember, a Brønsted-Lowry acid donates a proton, while a Brønsted-Lowry base accepts one.

In this example,

Acetic acid donates a proton to form acetate, making it the acid and acetate its conjugate base.

Methyl amine accepts a proton to form methyl ammonium, making it the base and methyl ammonium its conjugate acid.

Therefore, the acids are acetic acid and methyl ammonium.

Step 2: Obtain the pKa values for these acids

🗒️ Note: You will typically be provided with a pKa table during exams or you can find them in your textbook or a reliable online resource.

Based on the pKa values, this equilibrium favors the product side. Methyl ammonium, which has a higher pK_a ​, is a weaker acid compared to its counterpart on the reactant side. In acid-base equilibria, the reaction tends to favor the formation of the weaker acid, positioning the equilibrium toward the products in this scenario.

📝 Note: To visually indicate the direction in which the equilibrium favors either the reactants or the products, you can use one of these two arrows.

 

Step 3: Calculate K using the correct formula

K is calculated using the following formula

 

K = 10^−ΔpK_a

 

where ΔpK_a = pK_a of reactant acid− pK_a​ of product acid

Since the pKa​ of acetic acid is 4.76 and the pKa​ of methylammonium is 10.66

ΔpKa = 4.76 – 10.66 = -5.9

K = 10^-(-5.9) = 7.94 × 10⁵

An equilibrium constant, K, greater than 1 means the reaction favors the products. This matches our earlier prediction based on the relative strengths of the acids. By comparing the pKa values, we expected the reaction to favor the weaker acid on the product side, and the calculated K value confirms this.