Chemistry 9701/21 — October/November 2018

Background Concept

Iron has atomic number 26, so a neutral iron atom has 26 electrons. The electron configuration of neutral Fe is $1\text{s}^2\ 2\text{s}^2\ 2\text{p}^6\ 3\text{s}^2\ 3\text{p}^6\ 4\text{s}^2\ 3\text{d}^6$. When a transition metal forms a positive ion, electrons are removed from the outermost shell first — the 4s subshell — before the 3d subshell. This is because the 4s orbital is higher in energy than 3d once electrons occupy both.

Understanding the Question

The question asks for the full electronic configuration of $\text{Fe}^{2+}$, which means removing two electrons from neutral iron. The command word is "give", so a single line answer is expected.

Approach

Start from the configuration of neutral Fe (26 electrons), then remove two electrons from the 4s subshell to form Fe²⁺ (24 electrons).

Step-by-Step Reasoning

1. Neutral Fe: $1\text{s}^2\ 2\text{s}^2\ 2\text{p}^6\ 3\text{s}^2\ 3\text{p}^6\ 4\text{s}^2\ 3\text{d}^6$ (26 electrons total)
2. Fe²⁺ has lost 2 electrons. For transition metals, the 4s electrons are removed before the 3d electrons.
3. Removing both 4s electrons gives: $1\text{s}^2\ 2\text{s}^2\ 2\text{p}^6\ 3\text{s}^2\ 3\text{p}^6\ 3\text{d}^6$ (24 electrons total)
4. Note that $4\text{s}^0$ is sometimes written for clarity but is not required.

Key Takeaways

• Transition metal ions lose 4s electrons before 3d electrons.
• Fe²⁺ has 24 electrons; Fe³⁺ would have 23 (one more from 3d).

Common Mistakes

• Writing $4\text{s}^2$ at the end (forgetting that 4s electrons are removed first when forming the ion).
• Removing electrons from 3d instead of 4s, giving $1\text{s}^2\ 2\text{s}^2\ 2\text{p}^6\ 3\text{s}^2\ 3\text{p}^6\ 4\text{s}^2\ 3\text{d}^4$.
• Writing the configuration of neutral Fe instead of the ion.

Things to Be Careful About

• The question says "full" electronic configuration, so you must write all subshells explicitly rather than using the $[\text{Ar}]$ shorthand.
• The order of writing can be either $3\text{d}^6\ 4\text{s}^0$ or $4\text{s}^0\ 3\text{d}^6$; both are acceptable, but the key point is that 4s is empty.

Background Concept

The oxidation number (or oxidation state) is a formal charge assigned to an atom in a compound or ion, based on a set of rules. The most important rule for this question is: the sum of the oxidation numbers of all atoms in a polyatomic ion equals the overall charge of that ion.

Understanding the Question

We are told that $\text{S}_2^{2-}$ has an overall charge of $-2$ and that each sulfur atom has the same oxidation number. We need to find that common oxidation number.

Approach

Let the oxidation number of each S atom be $x$. Since there are two sulfur atoms and the total charge is $-2$:
$2x = -2$
$x = -1$

Step-by-Step Reasoning

1. The ion $\text{S}_2^{2-}$ has a total charge of $-2$.
2. There are two identical sulfur atoms, each with oxidation number $x$.
3. Sum of oxidation numbers = overall charge: $2x = -2$
4. Solving: $x = -1$

Key Takeaways

• The sum of oxidation numbers in a polyatomic ion equals the charge on the ion.
• In $\text{S}_2^{2-}$, sulfur has an unusual oxidation state of $-1$ (similar to the peroxide ion $\text{O}_2^{2-}$ where oxygen is $-1$).

Common Mistakes

• Assuming sulfur must have oxidation number $-2$ (its common state in simple sulfides like $\text{FeS}$). The $\text{S}_2^{2-}$ ion is a disulfide, analogous to peroxide.
• Forgetting that the charge is $-2$ for the whole ion, not per atom.

Things to Be Careful About

• The question explicitly states to assume each sulfur has the same oxidation number, so no need to consider different values.
• The answer must include the negative sign.

Background Concept

Sulfur dioxide is a major atmospheric pollutant released from the combustion of sulfur-containing fossil fuels and from the roasting of sulfide ores. When released, $\text{SO}_2$ reacts with water and oxygen in the atmosphere to form sulfurous acid ($\text{H}_2\text{SO}_3$) and sulfuric acid ($\text{H}_2\text{SO}_4$), which fall as acid rain. Acid rain has a pH of about 4–5 (compared to normal rainwater at about 5.6 due to dissolved $\text{CO}_2$). The increased acidity has widespread environmental effects.

Understanding the Question

The command words are "state and explain". We need to name one environmental consequence (M1) and explain how it causes damage (M2). This is a 2-mark question.

Approach

Choose the most well-known consequence — acid rain — and then select one specific effect from the mark scheme's list of acceptable explanations.

Step-by-Step Reasoning

1. M1: State that $\text{SO}_2$ causes acid rain (or contributes to acid rain).
2. M2: Explain one consequence. The mark scheme accepts any of:
   • Destroys/damages buildings, statues (limestone/calcium carbonate reacts with acid)
   • Kills or harms fish, coral, plants, crops, trees (deforestation)
   • Leaches salts/ions (e.g. aluminium) from soil into rivers/lakes
   • Leaches away soil nutrients
   • Causes breathing difficulties
   • Lowers pH / increases acidity of soil, rivers, oceans, seas

Any one of these explanations is sufficient for M2.

Key Takeaways

• $\text{SO}_2$ is the primary cause of acid rain.
• The environmental effects are diverse but all stem from increased acidity.
• "State and explain" questions require both the identification AND a reason/mechanism.

Common Mistakes

• Stating "pollution" without specifying acid rain or a named effect.
• Giving only the consequence without explaining the mechanism (e.g. saying "damages buildings" without mentioning the acidic nature).
• Saying "global warming" — this is caused by $\text{CO}_2$, not $\text{SO}_2$.

Things to Be Careful About

• M1 and M2 must be linked: the explanation must follow from the stated consequence.
• "Acid rain" alone without an explanation only earns M1.

Background Concept

Balancing a chemical equation requires that the number of atoms of each element is the same on both sides, satisfying the law of conservation of mass. When some coefficients are given, we use them to deduce the unknown coefficients systematically.

Understanding the Question

We are given the skeleton equation:
$4\text{FeS}_2 + \text{?O}_2 \rightarrow 2\text{Fe}_2\text{O}_3 + \text{?SO}_2$
We must find the coefficients of $\text{O}_2$ and $\text{SO}_2$ to balance the equation. This is a 2-mark question: M1 for identifying the correct species ($\text{O}_2$ and $\text{SO}_2$) and M2 for the correct coefficients (11 and 8).

Approach

Use the given coefficients (4 for $\text{FeS}_2$ and 2 for $\text{Fe}_2\text{O}_3$) to determine the number of each atom on each side, then solve for the unknown coefficients.

Step-by-Step Reasoning

1. Balance Fe: Left side: $4 \times 1 = 4$ Fe. Right side: $2 \times 2 = 4$ Fe. ✓ (already balanced)
2. Balance S: Left side: $4 \times 2 = 8$ S. Right side: coefficient of $\text{SO}_2$ must give 8 S, so it is 8.
3. Balance O: Right side: $(2 \times 3) + (8 \times 2) = 6 + 16 = 22$ O atoms. Left side: coefficient of $\text{O}_2$ must give 22 O, so it is $22/2 =$ 11.
4. Final check: $4\text{FeS}_2 + 11\text{O}_2 \rightarrow 2\text{Fe}_2\text{O}_3 + 8\text{SO}_2$
   • Fe: 4 = 4 ✓
   • S: 8 = 8 ✓
   • O: 22 = 22 ✓

Key Takeaways

• When some coefficients are given, use them to fix the atom counts and solve for the unknowns.
• Balance elements other than oxygen last (or use oxygen to find the final coefficient).

Common Mistakes

• Writing $\text{SO}_3$ instead of $\text{SO}_2$ (the question states sulfur dioxide is produced).
• Getting the coefficient of $\text{O}_2$ wrong by miscounting oxygen atoms on the right side.
• Writing the wrong species in the blanks (e.g. putting the coefficient of $\text{SO}_2$ in the $\text{O}_2$ blank).

Things to Be Careful About

• The mark scheme requires BOTH correct species AND correct coefficients. Getting the species right but coefficients wrong (or vice versa) loses a mark.
• The equation must be fully balanced — check all three elements.

Background Concept

Stoichiometric calculations link the masses of reactants and products through the mole concept and the balanced equation. The general pathway is: mass → moles (divide by $M_r$) → use mole ratio from equation → moles of unknown → mass of unknown (multiply by $M_r$).

Understanding the Question

We are given the mass of $\text{Fe}_2\text{O}_3$ produced (33.18 g) and must find the mass of $\text{FeS}_2$ that was roasted. We are told all $\text{FeS}_2$ was converted to $\text{Fe}_2\text{O}_3$, and we are given the $M_r$ values. The balanced equation from part (c)(ii) provides the mole ratio.

Approach

1. Convert mass of $\text{Fe}_2\text{O}_3$ to moles.
2. Use the stoichiometric ratio from the balanced equation ($4\text{FeS}_2 : 2\text{Fe}_2\text{O}_3$) to find moles of $\text{FeS}_2$.
3. Convert moles of $\text{FeS}_2$ to mass.

Step-by-Step Reasoning

1. Moles of $\text{Fe}_2\text{O}_3$:
   $n = \frac{m}{M_r} = \frac{33.18}{159.6} = 0.2079\text{ mol}$
   This is M1 in the mark scheme.

2. Mole ratio: From $4\text{FeS}_2 + 11\text{O}_2 \rightarrow 2\text{Fe}_2\text{O}_3 + 8\text{SO}_2$, the ratio is $\text{FeS}_2 : \text{Fe}_2\text{O}_3 = 4 : 2 = 2 : 1$. So for every mole of $\text{Fe}_2\text{O}_3$ produced, 2 moles of $\text{FeS}_2$ were consumed.

3. Moles of $\text{FeS}_2$:
   $n(\text{FeS}_2) = 0.2079 \times \frac{4}{2} = 0.4158\text{ mol}$

4. Mass of $\text{FeS}_2$:
   $m = n \times M_r = 0.4158 \times 120.0 = 49.9\text{ g}$
   This is M2 in the mark scheme (correct use of stoichiometry and 120.0).

Key Takeaways

• Always use the balanced equation to determine the mole ratio.
• The pathway is: mass → moles → (ratio) → moles → mass.
• Error carried forward (ecf) is allowed: if you used 55.00 g from the alternative instruction, the method would still earn M2.

Common Mistakes

• Using the wrong mole ratio (e.g. 1:1 instead of 4:2, or inverting the ratio).
• Dividing by 120.0 instead of multiplying at the final step.
• Forgetting to use the $M_r$ values given in the question.

Things to Be Careful About

• The mark scheme awards M1 for the mole calculation of $\text{Fe}_2\text{O}_3$ and M2 for the correct application of the ratio AND the multiplication by 120.0.
• The answer should be given to at least 3 significant figures (49.9 g). The mark scheme shows 49.89 g as acceptable.

Background Concept

Percentage by mass (mass percent) of a component in a mixture is calculated as:
$\text{mass percent} = \frac{\text{mass of component}}{\text{total mass of mixture}} \times 100\%$
Here, the impure sample consists only of $\text{FeS}_2$ and Au (as stated in the question), so the total mass is the sum of these two.

Understanding the Question

We need to find the percentage by mass of gold in the original impure iron pyrite sample. We know:

• Mass of Au = 0.37 g
• Mass of $\text{FeS}_2$ = 49.9 g (from part iii)
• Gold is the only impurity, so total mass = mass of $\text{FeS}_2$ + mass of Au.

The answer must be given to two significant figures.

Approach

Add the two masses to get the total, then divide the mass of gold by the total and multiply by 100.

Step-by-Step Reasoning

1. Total mass of impure sample = $49.9 + 0.37 = 50.27\text{ g}$ (using 49.89 gives 50.26 g)
2. Percentage of Au = $\frac{0.37}{50.27} \times 100 = 0.736\%$
3. Rounded to two significant figures: $0.74\%$

Note: If using the alternative value of 55.00 g for $\text{FeS}_2$ (as instructed for students who could not complete part iii):
$\frac{0.37}{55.00 + 0.37} \times 100 = \frac{0.37}{55.37} \times 100 = 0.67\%$
This would earn the mark as ecf.

Key Takeaways

• Percentage by mass requires dividing by the TOTAL mass of the mixture, not just one component.
• Always check the required number of significant figures.

Common Mistakes

• Dividing by only the mass of $\text{FeS}_2$ (49.9) instead of the total (50.27). This gives 0.74% by coincidence here but is methodologically wrong.
• Dividing by the mass of $\text{Fe}_2\text{O}_3$ (33.18) instead of the original $\text{FeS}_2$ mass.
• Giving the answer to the wrong number of significant figures.

Things to Be Careful About

• The question specifies two significant figures. The answer 0.74% has exactly two.
• The mark scheme shows $(0.37/(0.37 + 49.89)) = 0.74$, confirming the denominator is the total mass of the original impure sample.