Applying Chemical Ideas

Module 8 is worth about 26 marks a year. Analysis of organic substances (spectroscopy, about 10 marks a year) and analysis of inorganic substances (ion tests and quantitative analysis, about 9 marks) are the two big inquiry questions. Chemical synthesis and design is worth about 3–4 marks, usually as an extended "evaluate" response.

• Analysis of inorganic substances: why we monitor the environment; precipitation, complexation and flame tests for cations and anions; gravimetric analysis, precipitation titrations, colorimetry, UV-visible spectrophotometry and atomic absorption spectroscopy (AAS).
• Analysis of organic substances: chemical tests for unsaturation, hydroxyl groups and carboxylic acids; infrared (IR), carbon-13 and proton NMR, and mass spectrometry; combining evidence to identify a compound.
• Chemical synthesis and design: reaction conditions, yield, purity, rate and cost in industrial processes; environmental, social and economic issues.

The data sheet gives IR absorption ranges, NMR chemical shift data, $K_{sp}$ values, relative atomic masses and $A=\epsilon lc={\log }_{10}(I_0/I)$. Know how to use them; don't memorise every number.

Chemists monitor water, soil and air because human activity adds substances that are harmful even at low concentrations:

• Heavy metals (lead, mercury, cadmium) are toxic and bioaccumulate up food chains.
• Excess nutrients (nitrate and phosphate from fertilisers) cause algal blooms and eutrophication.
• Salinity and pH affect crops, aquatic life and drinking water.

Results are compared with guideline values (for example in the Australian Drinking Water Guidelines), often in mg per litre or parts per million (ppm). For dilute water, $1\text{ppm}\approx 1{\text{mg L}}^{-1}$.

Cations

The deep blue copper–ammonia and blood-red iron(III)–thiocyanate colours are complexation reactions: the metal ion bonds to surrounding ligands to form a complex ion.

Order matters. To separate a mixture, add reagents that precipitate one ion at a time: chloride first (removes ${\text{Ag}}^{+}$ and ${\text{Pb}}^{2+}$), filter, then sulfate (removes ${\text{Ba}}^{2+}$, and ${\text{Ca}}^{2+}$ if concentrated), then hydroxide. A flowchart is the clearest way to present a sequence.

Flame tests work because heat excites electrons to higher energy levels; as they fall back they emit light of specific wavelengths, so each element has a characteristic colour. Sodium contamination (strong yellow) can mask other colours, and colours like pale green and blue-green are easy to confuse.

Anions

Test for carbonate first, because carbonate also precipitates with ${\text{Ba}}^{2+}$ and ${\text{Ag}}^{+}$; acidifying removes it as ${\text{CO}}_2$. Use nitric acid (not HCl or ${\text{H}}_2{\text{SO}}_4$) to acidify before the silver nitrate test, so you don't add the ion you're testing for.

Gravimetric analysis

The ion is precipitated as a compound of known formula with very low $K_{sp}$, using excess precipitating agent (common ion effect). The precipitate is filtered, washed with distilled water to remove soluble impurities, then dried and weighed repeatedly to constant mass.

Sources of error: precipitate not fully dried (mass too high), unwashed precipitate (soluble salts add mass), some precipitate passing through the filter paper or slightly soluble (mass too low), and co-precipitation of other ions.

Worked example. Four iron supplement tablets were dissolved and all the iron was precipitated and heated to give 0.2150 g of ${\text{Fe}}_2{\text{O}}_3$. Find the mass of iron per tablet. ($M(\text{Fe})=55.85$, $M(\text{O})=16.00$.)

Each ${\text{Fe}}_2{\text{O}}_3$ contains 2 Fe:

Per tablet: $0.1504/4=0.03759$ g, which is 37.6 mg of iron.

Precipitation titrations

A titrant forms a precipitate with the analyte, and an indicator shows when the analyte is used up. In the Mohr method, chloride is titrated with standard silver nitrate using chromate as indicator. White AgCl forms first; once the chloride is gone, the next drop of ${\text{Ag}}^{+}$ forms red-brown silver chromate, which is the end point. The calculation is the same as an acid–base titration: moles of titrant, mole ratio, then concentration.

Colorimetry and UV-visible spectrophotometry

A coloured solution absorbs light of particular wavelengths. Choose the wavelength (or filter) the solution absorbs most strongly, usually the complementary colour. Beer–Lambert law (data sheet):

$A$ is absorbance (no units), $\epsilon$ is the molar absorptivity (${\text{L mol}}^{-1}{\text{cm}}^{-1}$), $l$ is the path length (cm) and $c$ is the concentration (${\text{mol L}}^{-1}$). Absorbance is directly proportional to concentration, so a calibration curve of standard solutions gives a straight line, and an unknown is read off it. Colourless ions can be analysed by first reacting them to form a coloured complex (for example iron(III) with thiocyanate). UV-visible spectrophotometers work the same way but scan wavelengths, including ultraviolet, so they also suit organic compounds that absorb UV.

Worked example. Standards of an iron(III)–thiocyanate complex gave:

10.0 mL of bore water was diluted to 50.0 mL and treated the same way; its absorbance was 0.302. Find the iron concentration in the bore water.

The line of best fit has gradient $0.119{\text{L mg}}^{-1}$ and passes through the origin (to three decimal places). Reading off:

The sample was diluted by a factor of $50.0/10.0=5$, so the bore water contains $2.54\times 5=12.7{\text{mg L}}^{-1}$ of iron.

Two things earn marks: using the line of best fit (not one data point) and undoing the dilution. A sample whose absorbance is above the highest standard must be diluted and re-measured, because the calibration is only reliable within its range.

Atomic absorption spectroscopy

AAS was developed in Australia by Alan Walsh at CSIRO in the 1950s. It measures the concentration of a specific metal element, often at ppm or ppb levels:

1. The sample solution is sprayed into a flame (or furnace), which breaks compounds into free gaseous atoms.
2. A hollow cathode lamp made of the element being measured emits exactly the wavelengths that element's atoms absorb.
3. Ground-state atoms in the flame absorb some of this light; the detector measures how much.
4. Absorbance is compared with a calibration curve of standards to find the concentration.

AAS is highly selective (other elements don't absorb the lamp's wavelengths) and sensitive, which is why it transformed monitoring of heavy metals in water, soil, food and blood. Its limitation is one element per lamp.

Chemical tests

Tertiary alcohols give no colour change with dichromate, and alcohols don't fizz with carbonate, so the two tests together separate alcohols from acids.

Infrared spectroscopy

Covalent bonds absorb IR radiation at frequencies that make them vibrate, so each bond type absorbs in a characteristic wavenumber range (${\text{cm}}^{-1}$). Use the data-sheet ranges to look for:

• O–H in alcohols: broad, strong band around 3230–3550.
• O–H in carboxylic acids: very broad band, about 2500–3000, overlapping the C–H peaks.
• C=O: strong, sharp band around 1680–1750 (aldehydes, ketones, acids, esters, amides).
• N–H: 3300–3500 in amines and amides.
• C–O: 1000–1300 (alcohols, esters, ethers).
• C=C: 1620–1680; C–H just below and above 3000.

Below about 1500 ${\text{cm}}^{-1}$ is the fingerprint region: too complex to interpret bond by bond, but unique to each compound, so it can be matched against a database. IR identifies functional groups; it rarely gives the full structure.

Carbon-13 NMR

Each signal corresponds to a set of carbon atoms in a different chemical environment; carbons related by symmetry give one signal. The chemical shift (ppm) shows what the carbon is attached to: alkyl carbons about 5–40 ppm, C–O carbons about 50–90 ppm, C=C about 100–150, and C=O carbons 160–220 ppm (acids and esters at the lower end, aldehydes and ketones at the upper end).

Proton NMR

Proton (¹H) NMR gives four pieces of information:

1. Number of signals: number of different hydrogen environments.
2. Chemical shift: the neighbouring groups (alkyl H about 0.9–1.5 ppm; H next to C=O about 2.0–2.7; H on C–O about 3.3–4.5; aldehyde H about 9–10; carboxylic acid H about 10–13).
3. Relative peak area (integration): the ratio of hydrogens in each environment.
4. Splitting (n + 1 rule): a signal is split into $n+1$ peaks by $n$ hydrogens on adjacent carbons. Singlet means no neighbours, a doublet one, a triplet two, a quartet three. H on O or N usually shows as a singlet.

An ethyl group (${\text{CH}}_3{\text{CH}}_2$–) always shows a 3H triplet coupled to a 2H quartet.

Mass spectrometry

Molecules are ionised and fragmented, and the ions are sorted by mass-to-charge ratio ($m/z$). The molecular ion (${\text{M}}^{+}$, usually the peak with the highest $m/z$) gives the molar mass. The tallest peak (base peak) is the most abundant fragment. Fragment losses identify groups: a loss of 15 suggests ${\text{CH}}_3$, 29 suggests ${\text{C}}_2{\text{H}}_5$ or CHO, and 45 suggests COOH. Peaks one unit higher than ${\text{M}}^{+}$ come from carbon-13.

Worked example: combining spectra

A compound ${\text{C}}_4{\text{H}}_8{\text{O}}_2$ gives:

• IR: strong sharp peak at 1740 ${\text{cm}}^{-1}$, strong peak near 1200 ${\text{cm}}^{-1}$, nothing broad above 3000.
• ¹³C NMR: four signals, near 174, 51, 27 and 9 ppm.
• ¹H NMR: 3.7 ppm (3H, singlet), 2.3 ppm (2H, quartet), 1.1 ppm (3H, triplet).
• Mass spectrum: ${\text{M}}^{+}$ at $m/z$ 88; peaks at 59 and 57.

Reasoning, one piece of evidence at a time:

1. Molar mass of ${\text{C}}_4{\text{H}}_8{\text{O}}_2$ is 88, matching ${\text{M}}^{+}$.
2. C=O present (1740) but no O–H, so it is not an acid; with C–O at 1200 it is an ester.
3. Four ¹³C signals: four different carbons, no symmetry. The 174 ppm carbon is the ester C=O; 51 ppm is a carbon bonded to O.
4. ¹H: the 2H quartet and 3H triplet are an ethyl group; the quartet at 2.3 ppm means the ${\text{CH}}_2$ is next to C=O, not O. The 3H singlet at 3.7 ppm is a ${\text{CH}}_3$ on oxygen with no neighbouring H.
5. Fragments: 57 is ${\text{CH}}_3{\text{CH}}_2{\text{CO}}^{+}$ and 59 is ${\text{COOCH}}_3^{+}$.

The compound is methyl propanoate, ${\text{CH}}_3{\text{CH}}_2{\text{COOCH}}_3$. Its isomer ethyl ethanoate would instead show a 3H singlet near 2.0 ppm and the quartet near 4.1 ppm, because there the ${\text{CH}}_2$ is bonded to oxygen. Full-mark answers use every spectrum and say what each one shows.

Industrial chemists choose conditions to balance:

• Yield (equilibrium position) against rate: for exothermic equilibria, low temperature gives more product but slowly, so a compromise temperature with a catalyst is used.
• Purity: separation and purification steps add cost; by-products may need treatment.
• Cost and energy: high pressure and temperature need expensive equipment and energy.
• Raw materials, safety, waste and environmental impact, plus social factors (jobs, location, community risk).

Example: sulfuric acid (Contact process). The key step is

which is exothermic. A temperature of about 400–450 °C with a vanadium(V) oxide catalyst gives a fast rate with high conversion. Higher pressure would favour ${\text{SO}}_3$ (fewer gas moles), but conversion is already very high, so only slightly above atmospheric pressure is used to save cost. ${\text{SO}}_3$ is absorbed into concentrated sulfuric acid rather than water, because reacting it directly with water produces a corrosive mist. Sulfur dioxide emissions must be controlled to prevent acid rain; heat from the exothermic steps is recovered to generate electricity.

Worked example: yield with a limiting reagent

5.00 g of propanoic acid ($M=74.08{\text{g mol}}^{-1}$) reacts with 4.00 g of propan-1-ol ($M=60.09{\text{g mol}}^{-1}$), giving 5.20 g of propyl propanoate ($M=116.16{\text{g mol}}^{-1}$). Calculate the percentage yield.

The ratio is 1 : 1, so propan-1-ol is limiting.

The yield is below 100% because esterification is an equilibrium and some product is lost during washing and distillation.

• Testing for halides after adding HCl, or for sulfate after adding ${\text{H}}_2{\text{SO}}_4$: you've added the ion you're looking for.
• Listing tests without observations. Every test needs reagent and expected result and conclusion.
• Forgetting the dilution factor in colorimetry or AAS, or reading the calibration graph off a single point.
• Saying AAS detects compounds or anions; it measures metal elements.
• Counting hydrogens on the same carbon when using the n + 1 rule; only neighbouring carbons count.
• Confusing the broad acid O–H band (2500–3000) with the alcohol O–H band (3230–3550).
• In gravimetric calculations, forgetting the mole ratio between the precipitate and the ion (two Fe per ${\text{Fe}}_2{\text{O}}_3$).
• Calling the base peak the molecular ion.

• Identify a compound from spectra: state the conclusion and link each spectrum to a structural feature. Marks are usually given per spectrum interpreted, plus the structure.
• Describe a sequence of tests: use a flowchart or numbered steps with reagent, observation and conclusion for each ion, in an order that avoids interference.
• Explain how AAS or colorimetry works: include why the wavelength is specific, how the calibration curve is made and used, and why the method suits low concentrations.
• Evaluate an industrial process: discuss rate, yield, cost, safety and environmental issues with specific chemistry, and make a judgement. Generic statements ("it's bad for the environment") earn little.
• Calculations here combine moles, mass and dilution; show every step with units and keep extra figures until the end.

• Module 5: $K_{sp}$ and the common ion effect underpin precipitation tests and gravimetric analysis; equilibrium drives industrial condition choices.
• Module 6: titration technique and calculations carry over to precipitation titrations; acid–base tests identify carbonates and acids.
• Module 7: the functional groups, reactions and naming you identify here come from Module 7.
• Year 11 Module 1: electron energy levels explain flame tests and atomic absorption.