A solution of potassium cyanide, KCN, is made by dissolving 3.2 g of KCN in water. When this solution reacts with excess hydrochloric acid, HCl, it produces 0.30 g of poisonous hydrogen cyanide gas, HCN.

KCN + HCl - -> KCl + HCN

(a) Calculate the theoretical yield of HCN.

(b) Calculate the percent yield of HCN.

22.7%

Explanation:

We must first put down the equation of reaction to guide our work while solving the problem.

KCN (aq) + HCl (aq) - -> KCl (aq) + HCN (aq)

The questions specifically state that HCl is in excess so KCN is the limiting reactant.

Number of moles of KCN reacted = mass of KCN reacted / molar mass of KCN

Mass of KCN reacted = 3.2 g

Molar mass of KCN = 65.12 g/mol

Number of moles of KCN = 3.2/65.12 g/mol = 0.049 moles

Theoretical yield of HCN is obtained thus;

From the reaction equation;

1 mol of KCN produced 1 mole of HCN thus 0.049 moles of KCN will produce 0.049 moles of HCN.

Mass of HCN = number of moles * molar mass

Molar mass of HCN = 27.0253 g/mol

Hence mass of HCN produced = 0.049mol * 27.0253 g/mol = 1.32g of HCN

Actual yield of HCN = 0.30g

% yield = actual yield / theoretical yield * 100

% yield = 0.30/1.32 * 100

%yield = 22.7%

The theoretical yield HCN is 1.33 grams

The percent yield HCN is 22.56 %

Explanation:

Step 1: Data given

Mass of KCN = 3.2 grams

Molar mass KCN = 65.12 g/mol

Mass of HCN produced = 0.30 grams

Step 2: The balanced equation

KCN + HCl - -> KCl + HCN

Step 3: Calculate moles KCN

Moles KCN = mass KCN / moalr mass KCN

Moles KCN = 3.2 grams / 65.12 g/mol

Moles KCN = 0.0491 moles

Step 4: Calculate moles HCN

For 1 mol KCN we need 1 mol HCl to produce 1 mol KCl and 1 mol HCN

For 0.0491 moles HCN we'll have 0.0491 moles HCN

Step 5: Calculate mass HCN

Mass HCN = moles HCN * molar mass HCN

Mass HCN = 0.0491 moles * 27.03 g/mol

Mass HCN = 1.33 grams = the theoretical yield

Step 6: Calculate the percent yield HCN

Percent yield = (actual yiel / theoretical yield) * 100 %

Percent yield = (0.30 grams / 1.33 grams)

Percent yield = 22.56 %