Thermochemistry worksheet with problems on enthalpy changes and energy calculations in chemical reactions.

Let's solve each problem step by step from the worksheet titled "Enthalpy Changes in Chemical Reactions".

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Problem 1:

> How much heat will be released if 7.24 g of hydrogen peroxide (H₂O₂) decomposes according to the following reaction?
>
> 2 H₂O₂(l) → 2 H₂O(l) + O₂(g)  ΔH = –190 kJ

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#### Step 1: Understand what’s given
- The reaction releases 190 kJ of heat when 2 moles of H₂O₂ decompose.
- We are given 7.24 grams of H₂O₂.
- We need to find how much heat is released for this amount.

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#### Step 2: Molar mass of H₂O₂
- H = 1.01 g/mol × 2 = 2.02
- O = 16.00 g/mol × 2 = 32.00
- Molar mass of H₂O₂ = 34.02 g/mol

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#### Step 3: Moles of H₂O₂
$$
\text{Moles} = \frac{\text{mass}}{\text{molar mass}} = \frac{7.24\ \text{g}}{34.02\ \text{g/mol}} = 0.2128\ \text{mol}
$$

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#### Step 4: Use stoichiometry
From the reaction:
- 2 moles H₂O₂ → release 190 kJ
- So, 1 mole H₂O₂ → releases $ \frac{190}{2} = 95 $ kJ

So for 0.2128 mol:
$$
\text{Heat released} = 0.2128\ \text{mol} \times 95\ \text{kJ/mol} = 20.216\ \text{kJ}
$$

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#### ✔ Answer:
> 20.2 kJ of heat is released (rounded to 3 significant figures).

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Problem 2:

> How much heat is transferred when 6.22 grams of glucose (C₆H₁₂O₆) in your body reacts with oxygen gas during cellular respiration?
>
> C₆H₁₂O₆(s) + 6 O₂(g) → 6 CO₂(g) + 6 H₂O(l)  ΔH = –2803 kJ

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#### Step 1: Molar mass of glucose (C₆H₁₂O₆)
- C = 12.01 × 6 = 72.06
- H = 1.01 × 12 = 12.12
- O = 16.00 × 6 = 96.00
- Total = 180.18 g/mol

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#### Step 2: Moles of glucose
$$
\text{Moles} = \frac{6.22\ \text{g}}{180.18\ \text{g/mol}} = 0.03452\ \text{mol}
$$

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#### Step 3: Heat released per mole
- ΔH = –2803 kJ per mole of glucose
- So, heat released = $ 0.03452 \times 2803 = 96.75\ \text{kJ} $

(We take absolute value since it asks "how much heat is transferred"—but we note it’s exothermic.)

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#### ✔ Answer:
> 96.8 kJ of heat is transferred (rounded to 3 significant figures).

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Problem 3:

> Determine the amount of energy released when 147 grams of NO₂ is produced in the reaction:
>
> 3 NO₂(g) + H₂O(l) → 2 HNO₃(l) + NO(g)  ΔH = –140 kJ

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Wait — let's analyze this carefully.

The reaction shows:
- 3 moles of NO₂ react → ΔH = –140 kJ (so 140 kJ released)

But the question says: "when 147 grams of NO₂ is produced"

Hold on — is NO₂ being produced or consumed?

Looking at the reaction:
- NO₂ is a reactant, not a product.
- But the question says "produced", which contradicts the reaction as written.

Wait — maybe there’s a typo in interpretation.

But let’s double-check:

Reaction:
3 NO₂(g) + H₂O(l) → 2 HNO₃(l) + NO(g) ΔH = –140 kJ

So NO₂ is consumed, not produced.

But the question says: *"when 147 grams of NO₂ is produced"*

This implies that NO₂ is a product, but in this reaction, it's a reactant.

So unless this is a reverse reaction, something is off.

But perhaps the reaction is meant to be interpreted differently?

Alternatively, maybe the reaction is not producing NO₂ — so the question might have a mistake?

But wait — maybe the question is asking about the production of NO₂ in a different context?

Wait — no, the reaction is given, and ΔH is for that reaction.

So if NO₂ is being produced, then we must be considering the reverse reaction.

But the ΔH given is for the forward reaction (where NO₂ is consumed).

So if NO₂ is being produced, that would be the reverse reaction, and ΔH would be +140 kJ.

But the question says: “energy released” — so it can’t be endothermic.

Thus, contradiction.

Wait — perhaps the reaction is miswritten?

Wait — actually, look again:

Is it possible that the reaction is written correctly, but the question means that NO₂ is consumed, and the phrasing “produced” is a mistake?

But the question clearly says: “when 147 grams of NO₂ is produced”

So unless the reaction is reversed, this doesn't make sense.

Alternatively, could it be a typo in the reaction?

Wait — let’s consider the possibility that the reaction is:

Maybe the reaction is supposed to be one where NO₂ is produced?

But as written, it’s not.

Alternatively, perhaps the reaction is:

2 HNO₃(l) + NO(g) → 3 NO₂(g) + H₂O(l) — reverse, with ΔH = +140 kJ

Then, producing NO₂ would require absorbing energy — but the question says “energy released”, so that can’t be.

So there’s an inconsistency.

Wait — unless the reaction is not consuming NO₂, but producing it?

But in the given reaction, NO₂ is a reactant.

So either:
- The reaction is wrong,
- Or the question has a typo.

But let’s suppose the question meant 147 grams of NO₂ is consumed, because otherwise it doesn’t make sense.

Alternatively, maybe the reaction is not the one shown — but let’s assume it’s correct.

Wait — another idea: maybe the reaction is reversed, but then ΔH = +140 kJ, and energy is absorbed, not released.

But the question says “energy released”.

So unless the sign is wrong, this is inconsistent.

But let’s check standard reactions.

Actually, the reaction:
3 NO₂ + H₂O → 2 HNO₃ + NO is well-known and exothermic.

So NO₂ is consumed.

Therefore, NO₂ cannot be produced in this reaction.

So likely, the question meant: when 147 grams of NO₂ is consumed, not produced.

Let’s proceed assuming that — it’s a common typo.

So assume: 147 grams of NO₂ is consumed.

Then we can calculate.

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#### Step 1: Molar mass of NO₂
- N = 14.01
- O₂ = 32.00
- Total = 46.01 g/mol

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#### Step 2: Moles of NO₂
$$
\frac{147\ \text{g}}{46.01\ \text{g/mol}} = 3.195\ \text{mol}
$$

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#### Step 3: Stoichiometry
From the reaction:
- 3 moles NO₂ → release 140 kJ
- So, per mole NO₂ → $ \frac{140}{3} = 46.67\ \text{kJ/mol} $

So for 3.195 mol:
$$
3.195 \times 46.67 = 149.0\ \text{kJ}
$$

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#### ✔ Answer (assuming NO₂ is consumed):
> 149 kJ of energy is released (rounded to 3 sig figs).

But since the question says “produced”, this may be incorrect.

Alternatively, if NO₂ is produced, then we need to reverse the reaction.

But then:
- Reverse reaction: ΔH = +140 kJ (endothermic)
- Energy is absorbed, not released

But the question says “energy released”

So contradiction.

Therefore, likely a typo in the question — it should say "consumed", not "produced".

So we’ll go with 149 kJ released.

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Problem 4:

> Calculate the amount of energy involved if 1.38 grams of H₂ are consumed in the following reaction:
>
> SO₂(g) + H₂(g) → S(s) + H₂O(l)  ΔH = –277 kJ/mol

Wait — ΔH = –277 kJ/mol — per mole of what?

It says ΔH = –277 kJ/mol — but mol of what?

Typically, ΔH is given for the reaction as written.

So the reaction is:
SO₂ + H₂ → S + H₂O ΔH = –277 kJ

That means: –277 kJ is released when 1 mole of H₂ is consumed.

Because H₂ has coefficient 1.

So yes, ΔH = –277 kJ per mole of H₂ consumed.

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#### Step 1: Molar mass of H₂
= 2.016 g/mol ≈ 2.02 g/mol

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#### Step 2: Moles of H₂
$$
\frac{1.38\ \text{g}}{2.02\ \text{g/mol}} = 0.6832\ \text{mol}
$$

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#### Step 3: Energy released
Since 1 mol H₂ → 277 kJ released
Then:
$$
0.6832 \times 277 = 189.4\ \text{kJ}
$$

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#### ✔ Answer:
> 189 kJ of energy is released (rounded to 3 significant figures).

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✔ Final Answers:

1. 20.2 kJ of heat is released.
2. 96.8 kJ of heat is transferred.
3. 149 kJ of energy is released (assuming NO₂ is consumed; likely typo in "produced").
4. 189 kJ of energy is released.

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