Equilibrium Constant Calculations - Online Tutor, Practice Problems & Exam Prep

Understanding the equilibrium constant (K_c) is crucial in chemical reactions. It can be calculated using the concentrations of reactants and products at equilibrium. For gaseous reactions, K_c is derived from the ratio of the products' concentrations to the reactants', each raised to the power of their stoichiometric coefficients. In a given example, the K_c is determined by plugging in the equilibrium concentrations of methane, carbon dioxide, carbon monoxide, and hydrogen gas into the expression:

K_c=[CO]²[H₂]²[CH₄][CO₂] The resulting value, after considering significant figures, gives us the equilibrium constant for the reaction, which is dimensionless and in this case, is 0.0335. This figure reflects the ratio of product and reactant concentrations at equilibrium, a fundamental concept in chemical thermodynamics.

Equilibrium Constant Calculations

Here we can say that when all equilibrium concentrations of reaction are known then our equilibrium constant K and be calculated. In addition to this we can also say that your equilibrium constant can also be used to calculate 1 missing equilibrium concentration.

If we take a look here at this example question, it says what is the value of the equilibrium constant KC for the reaction below if the equilibrium mixture contains 0.255 molar of methane, 1.10 molar of carbon dioxide, 0.388 more of carbon monoxide and 0.250 molar of hydrogen gas. Now here they're telling us this is an equilibrium mixture. So these are all equilibrium concentrations.

We say that K equals products over reactants. Remember, we exclude solids and liquids from this expression. Since everything is a gas, everything is going to be included. We're going to have here products overreacted. So that's CO2. So if it's coefficient is a 2, that's going to be squared times H2. Coefficient is 2, so that's also squared divided by methane, which has a coefficient of 1, so we don't need to include it times the concentration of CO2.

Now we're going to plug in the values given to us for each of the compounds. We have carbon monoxide as being 0.388 molar which is going to be squared times 0.250 molar for H2 which is also squared divided by. Here we have 0.255 molar for methane and we have 1.10 molar for carbon dioxide. We plug this into our calculators and what we will get initially, and I'm going to give a long string of numbers and then we're going to round it down.

So we're going to say here we have 0.033543672. Our equilibrium constant K has no unit, so we're not going to plug in molarity or anything. It's just this value Here. We're going to say within our question, all the values given to us have three significant figures. So we're going to give K 3 significant figures. This means that our final answer is going to be 0.0335. This represents the answer for our equilibrium constant K from the given example question.

The reaction: 2 NO(g) + Br₂(g) ⇌ 2 NOBr(g), has a K_p of 2.5 x 10² at 35°C. Calculate the equilibrium concentration of NOBr, if equilibrium concentrations of NO and Br₂ are 0.2 atm and 0.050 atm, respectively.

• A. 2 atm
• B. 0.7 atm
• C. 1 atm
• D. 0.5 atm

For the reaction below, K_c = 1.5 at a constant temperature. A 3.2 L flask contains an equilibrium mixture of 3 compounds: 3.7 g of NH₄HS, 70. g of NH₃ and unknown amount of H₂S. What is the mass (grams) of H₂S produced at equilibrium?

NH₄HS(s) ⇌ NH₃(g) + H₂S(g)

• A. 12.4 g H₂S
• B. 433 g H₂S
• C. 130 g H₂S
• D. 210 g H₂S

Here’s what students ask on this topic:

How do you calculate the equilibrium constant?

To calculate the equilibrium constant (K), you first need to understand that it is a numerical value that expresses the ratio of the concentrations of products to reactants at equilibrium, each raised to the power of their respective coefficients from the balanced chemical equation.

Here's a step-by-step guide:

1. Write down the balanced chemical equation for the reaction.
2. Identify the reactants and products along with their coefficients.
3. At equilibrium, measure the concentrations of each reactant and product in moles per liter (M).
4. Use the expression for the equilibrium constant, K, which is given by:

K = ^{[C]c [D]d} / _{[A]^a [B]^b}

where:

• [A] and [B] are the molar concentrations of the reactants,
• [C] and [D] are the molar concentrations of the products,
• a, b, c, and d are the coefficients of the reactants and products from the balanced equation.

5. Plug in the equilibrium concentrations into the expression and calculate the value of K.

Remember that for gases, partial pressures can be used instead of concentrations, leading to the equilibrium constant for pressure, K_p. Also, solids and pure liquids do not appear in the expression as their concentrations do not change.

How can the equilibrium constant be calculated without concentrations?

The equilibrium constant (K) for a chemical reaction can be calculated without directly using concentrations if you have access to the standard Gibbs free energy change (ΔG°) for the reaction. This is because ΔG° and K are related through the equation:

$\mathrm{\Delta G°}=-RTln\left(K\right)$

Where:

• ΔG° is the standard Gibbs free energy change in joules per mole (J/mol)
• R is the universal gas constant (8.314 J/mol·K)
• T is the temperature in Kelvin (K)
• K is the equilibrium constant

To find K without concentrations, you would rearrange the equation to solve for K:

$K=e^{\frac{-\mathrm{\Delta G°}}{RT}}$

You'll need to know the value of ΔG° at a given temperature to use this equation. This information can often be found in thermodynamic tables or calculated from standard enthalpy (ΔH°) and entropy (ΔS°) changes using the equation:

How do you calculate the equilibrium constant for a reaction?

To calculate the equilibrium constant (K) for a reaction, you need to know the balanced chemical equation and the concentrations of the reactants and products at equilibrium. The general formula for the equilibrium constant of a reaction is:

$K=\frac{{\left[C\right]}^c{\left[D\right]}^d}{{\left[A\right]}^a{\left[B\right]}^b}$

Here, [A], [B], [C], and [D] represent the molar concentrations of the reactants and products at equilibrium, and a, b, c, and d are the stoichiometric coefficients from the balanced chemical equation.

For a generic reaction:

$aA+bB\rightleftharpoons cC+dD$

You would raise the concentration

How do you calculate the equilibrium constant, K_c?

To calculate the equilibrium constant, K_c, for a chemical reaction, you need to know the balanced chemical equation and the concentrations of the reactants and products at equilibrium. The equilibrium constant is a ratio of the concentrations of the products raised to the power of their coefficients to the concentrations of the reactants raised to the power of their coefficients.

Here's the general formula for a reaction aA + bB → cC + dD:

$K{c}_{}=\frac{\left[C\right]{}^c\left[D\right]{}^d}{\left[A\right]{}^a\left[B\right]{}^b}$

Where:

• [A], [B], [C], and [D] are the molar concentrations of the reactants and products at equilibrium.
• a, b, c, and d are the stoichiometric coefficients of the reactants and products from the balanced equation.

Remember that only gases and aqueous solutions are included in the expression. Pure solids and liquids do not appear in the K_c expression. Make sure to use equilibrium concentrations, not initial concentrations, and ensure that all concentrations are