Topic 8

Topic 8 on a page

Everything from this topic in one place, for the last read before an exam. If anything here is unfamiliar, go back to the full page.

Definitions to learn word for word

• Oxidation number: the charge an atom would have if the substance were made of ions, with the shared electrons given entirely to the more electronegative atom (Oxidation number and the rules).
• Oxidation: loss of electrons, or an increase in oxidation number. Reduction: gain of electrons, or a decrease in oxidation number (Oxidation and reduction).
• Oxidising agent: takes electrons and is itself reduced. Reducing agent: gives electrons and is itself oxidised (Oxidising and reducing agents).
• Disproportionation: the same element is both oxidised and reduced in the same reaction (Disproportionation).
• First ionisation energy: the energy needed to remove one electron from each atom in one mole of gaseous atoms (Ionisation energy).
• Flame colour: the light given out when excited electrons fall back to lower energy levels (Flame colours).
• Standard solution: a solution of accurately known concentration (Core Practical 4).
• End point: the point at which the indicator changes colour, taken as the equivalence point (Concentration and titration).

Oxidation number rules (in priority order)

Write the sign first: +7. A charge has the sign after: 1−. Roman numerals in names are oxidation numbers: manganate(VII) has Mn at +7, iron(III) has Fe at +3.

Equations from half-equations

1. Balance the atom being changed, then O with H₂O, then H with H⁺, then charge with e⁻.
2. Multiply so both half-equations have the same number of electrons.
3. Add and cancel the electrons, and any water or H⁺ that appears on both sides.

Check atoms and charge: left $-1 + 8 + 10 = +17$; right $+2 + 15 = +17$.

Groups 1 and 2

Group 1 lithium is out of line: its carbonate and nitrate decompose on heating, because Li⁺ has a high charge density. Charge density = charge ÷ radius: Mg²⁺ 27.8, Ca²⁺ 20.0, Ba²⁺ 14.7, Li⁺ 13.2, Na⁺ 9.8, K⁺ 7.2 (charge per nm).

Group 1 nitrates (not lithium) give the nitrite and oxygen: $\ce{2NaNO3 -> 2NaNO2 + O2}$.

Flame colours and tests

Silver halides in ammonia: AgCl dissolves in dilute NH₃; AgBr dissolves only in concentrated NH₃; AgI does not dissolve.

Group 7

• Displacement: a more reactive halogen oxidises the halide ion of a less reactive one, so Cl₂ displaces Br⁻ and I⁻, Br₂ displaces I⁻ only.
• Chlorine and water: $\ce{Cl2 + H2O <=> HCl + HClO}$. Chlorine goes 0 to −1 and 0 to +1.
• Chlorine and alkali: cold dilute NaOH gives NaCl, NaClO and water; hot NaOH gives NaCl and NaClO₃:

• Halides with concentrated H₂SO₄: NaCl gives HCl only; NaBr gives HBr, Br₂ and SO₂; NaI gives HI, I₂, SO₂, S and H₂S. Sulfur goes +6 to +4 (SO₂), 0 (S), −2 (H₂S).
• Astatine is predicted to be a black solid with a higher boiling temperature than iodine, and At⁻ the strongest reducing halide.

Titration in six steps

1. Rinse the burette with the solution it will hold, fill it, and remove the air from the tip.
2. Pipette 25.0 cm³ into a clean conical flask, add a few drops of indicator.
3. Do a rough titration, then repeat, adding drop by drop near the end point, until two concordant titres (within 0.10 cm³).
4. Mean of the concordant titres only. Do not include the rough one.
5. Moles of the known solution: $n = cV$ with $V$ in dm³.
6. Use the equation ratio to get moles of the unknown, then $c = n \div V$. To get g dm⁻³ multiply by $M_\text{r}$.

Uncertainty: percentage uncertainty = (absolute uncertainty ÷ measured value) × 100. Burette titre: two readings, so double the single-reading uncertainty (±0.05 each gives ±0.10).

The ten traps

1. Writing the oxidation number with the sign after (2+ for +2), or the ion charge with the sign first.
2. Forgetting that the subscript multiplies (O₂ in SO₂ gives −4 in total).
3. Forgetting that oxygen in peroxides and hydrogen in metal hydrides is −1.
4. Adding half-equations without balancing the electrons first.
5. Leaving H⁺ or H₂O on both sides of a final equation.
6. Calling a reaction a disproportionation without naming the one element that goes both up and down.
7. Explaining thermal stability with "bond strength" instead of charge density and polarisation.
8. Giving the flame colour as "white" for magnesium, or the reds from memory without the mark-scheme wording.
9. Doing the sulfate test without acid first, so carbonate also gives a white precipitate.
10. Including the rough titre in the mean, or forgetting the mole ratio.

Mixed test

1. Work out the oxidation number of Cr in $\ce{Cr2O7^2-}$ and of Mn in $\ce{MnO4-}$.
2. Write the half-equation for the reduction of $\ce{MnO4-}$ to $\ce{Mn^2+}$ in acid, and for the oxidation of $\ce{Fe^2+}$ to $\ce{Fe^3+}$, and combine them.
3. Chlorine reacts with cold dilute sodium hydroxide. Write the equation and show that it is a disproportionation.
4. 25.0 cm³ of 0.100 mol dm⁻³ sodium hydroxide needs 21.40 cm³ of hydrochloric acid for neutralisation. Calculate the concentration of the acid, and the percentage uncertainty in the titre from the burette alone (±0.05 cm³ per reading).
5. Explain why magnesium carbonate decomposes at a lower temperature than barium carbonate.
6. A white solid gives a lilac flame, fizzes with dilute acid and gives a gas that turns limewater milky. Name the solid and justify each observation.
7. Describe what you would see when concentrated sulfuric acid is added to solid sodium iodide, and name the sulfur-containing products with their oxidation numbers.
8. Describe how you would use silver nitrate and ammonia to distinguish chloride, bromide and iodide ions.
9. Explain, using oxidation numbers, why bromine water turns brown when added to potassium iodide solution.
10. Predict the state and colour of astatine at room temperature, and say whether chlorine would displace astatide ions.