QuestionPractice 8 How much heat is evolved when 266 g of white phosphorus ( $P_4$ ) burns in air?
$P_4(s) + 5O_2(g) \xrightarrow{\Delta} P_4O_{10}(s)$ $\Delta H = -3013$ kJ

Studdy Solution

STEP 1

1. The chemical reaction given is $P_4(s) + 5O_2(g) \rightarrow P_4O_{10}(s)$ .
2. The enthalpy change ( $\Delta H$ ) for the reaction is $-3013$ kJ, indicating that the reaction is exothermic.
3. We need to calculate the heat evolved when 266 g of $P_4$ burns.
4. The molar mass of $P_4$ is needed to convert grams to moles.

STEP 2

1. Calculate the molar mass of $P_4$ .
2. Convert the mass of $P_4$ to moles.
3. Use the stoichiometry of the reaction to find the heat evolved.

STEP 3

Calculate the molar mass of $P_4$ .
The atomic mass of phosphorus (P) is approximately 31 g/mol. Since $P_4$ consists of 4 phosphorus atoms, the molar mass of $P_4$ is:
$\text{Molar mass of } P_4 = 4 \times 31 \, \text{g/mol} = 124 \, \text{g/mol}$

STEP 4

Convert the mass of $P_4$ to moles.
Given mass of $P_4$ is 266 g. Using the molar mass calculated:
$\text{Moles of } P_4 = \frac{266 \, \text{g}}{124 \, \text{g/mol}}$
$\text{Moles of } P_4 = 2.145 \, \text{mol}$

STEP 5

Use the stoichiometry of the reaction to find the heat evolved.
The reaction shows that 1 mole of $P_4$ releases 3013 kJ of heat. Therefore, for 2.145 moles of $P_4$ :
$\text{Heat evolved} = 2.145 \, \text{mol} \times (-3013 \, \text{kJ/mol})$
$\text{Heat evolved} = -6463.785 \, \text{kJ}$
The negative sign indicates that the heat is evolved (released).
The heat evolved when 266 g of $P_4$ burns in air is approximately:
$\boxed{-6463.785 \, \text{kJ}}$

Was this helpful?

Get Studdy

Scan QR code with your device to get Studdy on iOS or Android

QR code

Studdy solves anything!

Download the app

Start learning now

Download Studdy AI Tutor now. Learn with ease and get all help you need to be successful at school.

Parents Influencer program Contact Policy Terms