Magnesium nitrate decomposes when heated to form magnesium oxide. 2Mg(NO$_3$)$_2$ → 2MgO + 4NO$_2$ + O$_2$ (a) (i) Calculate the mass of oxygen made when 0.45 moles of magnesium nitrate decomposes. Relative atomic mass (A$_r$) O = 16.0. Mass of oxygen = ..................................................... g [3] (ii) Calculate how many molecules of nitrogen dioxide, NO$_2$, are produced from 0.45 moles of magnesium nitrate. The Avogadro constant is 6.02 × 10$^{23}$. Give your answer to 3 significant figures. Number of molecules of NO$_2$ = ........................................... [3]

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Magnesium nitrate decomposes when heated to form magnesium oxide - OCR Gateway - GCSE Chemistry - Question 22 - 2023 - Paper 3

Question 22

Magnesium nitrate decomposes when heated to form magnesium oxide. 2Mg(NO$_3$)$_2$ → 2MgO + 4NO$_2$ + O$_2$ (a) (i) Calculate the mass of oxygen made when 0.45 mole... show full transcript

Worked Solution & Example Answer:Magnesium nitrate decomposes when heated to form magnesium oxide - OCR Gateway - GCSE Chemistry - Question 22 - 2023 - Paper 3

Step 1

Calculate the mass of oxygen made when 0.45 moles of magnesium nitrate decomposes.

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Answer

To determine the mass of oxygen produced, we first need to analyze the decomposition reaction.

From the balanced equation:

$2Mg(NO_3)_2 \rightarrow 2MgO + 4NO_2 + O_2$

We see that 2 moles of magnesium nitrate produce 1 mole of O $_2$ . Therefore, if 0.45 moles of magnesium nitrate decompose, the moles of oxygen produced can be calculated as follows:

$\text{Moles of } O_2 = \frac{0.45 ext{ moles } Mg(NO_3)_2}{2} = 0.225 ext{ moles } O_2$

Next, we find the mass of oxygen using the formula:

$\text{Mass} = \text{Moles} \times \text{Molar mass}$

Here, the molar mass of O $_2 = 2 \times 16.0 = 32.0 ext{ g/mol}$ . Therefore:

$\text{Mass of } O_2 = 0.225 \times 32.0 = 7.2 ext{ g}$

So the mass of oxygen produced is 7.2 g.

Step 2

Calculate how many molecules of nitrogen dioxide, NO₂, are produced from 0.45 moles of magnesium nitrate.

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Answer

From the balanced equation:

$2Mg(NO_3)_2 \rightarrow 2MgO + 4NO_2 + O_2$

2 moles of magnesium nitrate produce 4 moles of NO $_2$ . Therefore, if 0.45 moles of magnesium nitrate decompose, the moles of nitrogen dioxide produced can be calculated as:

$\text{Moles of } NO_2 = 0.45 ext{ moles } Mg(NO_3)_2 \times \frac{4 ext{ moles } NO_2}{2 ext{ moles } Mg(NO_3)_2} = 0.90 ext{ moles } NO_2$

Now, to find the number of molecules of NO $_2$ , we use Avogadro's constant:

$\text{Number of molecules} = \text{Moles} \times 6.02 \times 10^{23}$

So,

$\text{Number of molecules of } NO_2 = 0.90 \times 6.02 \times 10^{23} \approx 5.418 \times 10^{23}$

Rounding to 3 significant figures, we get:

$5.42 \times 10^{23} ext{ molecules of } NO_2$

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Magnesium nitrate decomposes when heated to form magnesium oxide - OCR Gateway - GCSE Chemistry - Question 22 - 2023 - Paper 3

Worked Solution & Example Answer:Magnesium nitrate decomposes when heated to form magnesium oxide - OCR Gateway - GCSE Chemistry - Question 22 - 2023 - Paper 3

Calculate the mass of oxygen made when 0.45 moles of magnesium nitrate decomposes.

Calculate how many molecules of nitrogen dioxide, NO₂, are produced from 0.45 moles of magnesium nitrate.

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