# Analysis of Inorganic Substances

> Why we monitor ions in the environment, tests for cations and anions (precipitates, flame tests and complexes), gravimetric analysis, precipitation titrations, colorimetry, UV-visible spectroscopy and atomic absorption spectroscopy.

From HSC Chemistry on WhizzNotes, free study notes for NSW students: https://whizznotes.com/chemistry/analysing-inorganic-substances/. Updated 3 November 2026.
Year 11 and Year 12 boxes go deeper than the Year 10 text around them.

### Story time: The poisoned town
In the 1950s, people in the fishing town of Minamata in Japan began to fall ill. Cats twitched and fell into the sea; then people lost feeling in their hands and could not walk. It took years to find the cause: a factory had been releasing **mercury** into the bay, and it had built up in the fish.

Today, water authorities test rivers, dams and drinking water for dangerous ions every week. Some tests use simple colour changes in a test tube. Others use machines that can detect a few **parts per million**, or even less. This chapter shows how chemists find out which ions are present, and how much.

**In short:** Inorganic analysis finds which ions are in a sample and how much of each. Precipitate and flame tests identify ions, gravimetric analysis weighs a dried precipitate, and colorimetry, UV-visible spectroscopy and atomic absorption spectroscopy measure concentration with a calibration curve.

## Why monitor ions?

| Ion | Source | Problem |
|---|---|---|
| Lead $\mathrm{Pb^{2+}}$, mercury $\mathrm{Hg^{2+}}$, cadmium $\mathrm{Cd^{2+}}$ | Mining, old paint, batteries, industry | Toxic heavy metals; build up in food chains |
| Phosphate $\mathrm{PO_4^{3-}}$, nitrate $\mathrm{NO_3^-}$ | Fertiliser run-off, sewage | Algal blooms (eutrophication) that use up oxygen and kill fish |
| Salt ($\mathrm{Na^+}$, $\mathrm{Cl^-}$) | Land clearing, irrigation | Salinity damages crops and soil |
| Copper $\mathrm{Cu^{2+}}$ | Mines, pipes | Toxic to fish and algae |

Human activities such as mining, farming and industry can raise the concentration of these ions in waterways.

## Concentration units

- **ppm** (parts per million) = mg per kg, or for dilute water, **mg per L**.
- 0.025 g of an ion in 500 mL is $25 \text{ mg} \div 0.500 \text{ L} = 50$ ppm.

## Testing for ions (qualitative analysis)

**Cations: precipitation tests**

| Add | Result | Ion present |
|---|---|---|
| Dilute HCl | White precipitate | $\mathrm{Ag^+}$ ($\mathrm{AgCl}$) or $\mathrm{Pb^{2+}}$ ($\mathrm{PbCl_2}$) |
| Dilute $\mathrm{H_2SO_4}$ | White precipitate | $\mathrm{Ba^{2+}}$, $\mathrm{Ca^{2+}}$ or $\mathrm{Pb^{2+}}$ |
| NaOH | Blue precipitate | $\mathrm{Cu^{2+}}$ |
| NaOH | Green precipitate | $\mathrm{Fe^{2+}}$ |
| NaOH | Red-brown precipitate | $\mathrm{Fe^{3+}}$ |
| NaOH | White precipitate | $\mathrm{Mg^{2+}}$, $\mathrm{Ca^{2+}}$ or $\mathrm{Pb^{2+}}$ |

Testing in a **sequence** (a flow chart) separates the possibilities step by step: add HCl first to remove $\mathrm{Ag^+}$ and $\mathrm{Pb^{2+}}$, then $\mathrm{H_2SO_4}$, then NaOH, and so on.

**Flame tests.** Heat a sample on a clean wire in a Bunsen flame.

| Ion | Flame colour |
|---|---|
| $\mathrm{Na^+}$ | Yellow-orange |
| $\mathrm{K^+}$ | Lilac |
| $\mathrm{Ca^{2+}}$ | Brick red |
| $\mathrm{Ba^{2+}}$ | Pale green |
| $\mathrm{Cu^{2+}}$ | Blue-green |
| $\mathrm{Sr^{2+}}$ | Crimson |

Heat **excites electrons** to higher energy levels. As they fall back, they give out light of particular wavelengths, which our eyes see as a colour. Mg gives no flame colour.

**Complexes.** Some ions form coloured **complex ions**, where molecules or ions bond to the metal using **coordinate covalent bonds** (both shared electrons come from one atom):

- $\mathrm{Fe^{3+}}$ + thiocyanate $\mathrm{SCN^-}$ → blood-red $\mathrm{FeSCN^{2+}}$.
- $\mathrm{Cu^{2+}}$ + excess ammonia → deep blue $\mathrm{[Cu(NH_3)_4]^{2+}}$.

**Anions**

| Ion | Test | Result |
|---|---|---|
| Carbonate $\mathrm{CO_3^{2-}}$ | Add dilute acid | Fizzes; gas turns limewater milky |
| Sulfate $\mathrm{SO_4^{2-}}$ | Add $\mathrm{Ba^{2+}}$ (with HCl present) | White precipitate of $\mathrm{BaSO_4}$ |
| Chloride $\mathrm{Cl^-}$ | Add $\mathrm{AgNO_3}$ (with $\mathrm{HNO_3}$) | White precipitate, dissolves in dilute ammonia |
| Bromide $\mathrm{Br^-}$ | Add $\mathrm{AgNO_3}$ | Cream precipitate |
| Iodide $\mathrm{I^-}$ | Add $\mathrm{AgNO_3}$ | Yellow precipitate |
| Phosphate $\mathrm{PO_4^{3-}}$ | Add ammonium molybdate (with acid) | Yellow precipitate |
| Hydroxide $\mathrm{OH^-}$ | Red litmus | Turns blue |

Acid is added first in the sulfate and halide tests to **remove carbonate** ions, which would also give a white precipitate.

## Gravimetric analysis

Gravimetric analysis finds the amount of an ion by **weighing a precipitate**.

1. Weigh the sample and dissolve it.
2. Add **excess** of a reagent to precipitate all of the ion.
3. **Filter**, **wash** (to remove soluble impurities) and **dry** the precipitate in an oven.
4. Reweigh, then dry again until the mass stops changing (**constant mass**).
5. Use the mass to calculate the moles of the ion.

### Worked example: Sulfur in fertiliser
1.50 g of fertiliser gives 2.10 g of $\mathrm{BaSO_4}$ (233.4 g/mol).

$n(\mathrm{BaSO_4}) = 2.10 \div 233.4 = 0.00900$ mol $= n(\mathrm{S})$

Mass of S $= 0.00900 \times 32.07 = 0.289$ g

% S $= 0.289 \div 1.50 \times 100 = 19.2\%$

Errors: precipitate not fully dry (result too **high**); some precipitate lost through the filter paper (too **low**); other ions co-precipitating (too high).

## Precipitation titrations

Some ions can be titrated by forming a precipitate. In the **Mohr method**, chloride is titrated with silver nitrate, $\mathrm{Ag^+} + \mathrm{Cl^-} \rightarrow \mathrm{AgCl(s)}$. A little chromate indicator gives a red-brown precipitate of $\mathrm{Ag_2CrO_4}$ once all the chloride is used up.

## Colorimetry and UV-visible spectroscopy

A coloured solution absorbs some wavelengths of light. The more concentrated it is, the more it absorbs.

1. Choose the wavelength (or filter colour) that the solution **absorbs most**. This is the **complementary** colour: a blue solution absorbs orange-red light.
2. Zero the instrument with a **blank** (the solvent, usually distilled water).
3. Measure the absorbance of several **standard solutions** of known concentration.
4. Plot a **calibration curve** (absorbance against concentration). It should be a straight line through the origin.
5. Measure the sample, read off its concentration, and allow for any dilution.

A **UV-visible spectrophotometer** does the same but can choose any exact wavelength, including ultraviolet, so it can also measure colourless substances. Electrons in the molecules or ions absorb energy and jump to higher levels.

## Atomic absorption spectroscopy (AAS)

AAS measures the concentration of a **metal element**, down to parts per billion.

1. The sample solution is sprayed into a **flame**, which turns it into free **atoms**.
2. A **hollow cathode lamp** made of the metal being tested gives out light of exactly the wavelengths that **those atoms** absorb.
3. The atoms absorb some of this light. A detector measures how much.
4. Absorbance is compared with a calibration curve from standards.

Because each element absorbs its **own unique wavelengths**, other metals in the sample do not interfere. Alan Walsh developed AAS at the CSIRO in Australia in the 1950s.

**Read only inside the calibration range.** If a sample's absorbance is higher than the most concentrated standard, **dilute** it and measure again (or prepare more standards). Do not extend the line.

**Mind map**

```text
Inorganic analysis
  Why
    Heavy metals toxic
    Phosphate, nitrate: algal blooms
    ppm = mg/L
  Qualitative
    Precipitates
    Flame tests
    Complexes
    Anion tests
  Gravimetric
    Excess reagent
    Filter, wash, dry
    Constant mass
  Titration
    Mohr: Ag+ and Cl−
  Colorimetry and UV-vis
    Complementary colour
    Blank, standards
    Calibration curve
  AAS
    Flame atomises
    Hollow cathode lamp
    Element specific
```

### Year 11: Calibration curves
A calibration curve is a straight line through the origin, so absorbance is **proportional** to concentration.

**Example.** A 2.0 ppm copper standard has absorbance 0.20. A diluted sample has absorbance 0.34. Its concentration is $0.34 \div 0.20 \times 2.0 = 3.4$ ppm. If the sample was diluted 10 times before measuring, the original was **34 ppm**.

The line is only trustworthy **between** the lowest and highest standards. For example, with standards from 1.0 to 4.0 ppm, readings between 1.0 and 4.0 ppm are the most reliable.

### Year 12: The Beer-Lambert law and harder gravimetric problems
**Beer-Lambert law.** $A = \varepsilon l c$, where $A$ is absorbance, $\varepsilon$ is the molar absorptivity, $l$ is the path length (usually 1 cm) and $c$ is the concentration. With $\varepsilon = 1.2 \times 10^4$ L mol⁻¹ cm⁻¹, $l = 1.0$ cm and $A = 0.60$: $c = 0.60 \div 1.2 \times 10^4 = 5.0 \times 10^{-5}$ mol/L.

Absorbance is also written $A = \log_{10}\dfrac{I_{\text{in}}}{I_{\text{out}}}$, where $I_{\text{in}}$ is the light going in and $I_{\text{out}}$ the light coming out.

**Copper in brass by AAS.** 0.500 g of brass is dissolved and made up to 1.000 L. This solution is diluted 50 times, and AAS reads 6.5 ppm. The original solution is $6.5 \times 50 = 325$ mg/L, so the 1.000 L holds 325 mg $= 0.325$ g of copper: $0.325 \div 0.500 \times 100 = 65\%$ Cu.

**Salt in food.** 2.00 g of food gives 0.287 g of $\mathrm{AgCl}$ (143.32 g/mol): 0.00200 mol, so 0.00200 mol of NaCl, which is $0.00200 \times 58.44 = 0.117$ g, or **5.85%**.

**Mohr titration.** 20.0 mL of water needs 18.5 mL of 0.100 mol/L $\mathrm{AgNO_3}$. $n(\mathrm{Ag^+}) = 1.85 \times 10^{-3}$ mol $= n(\mathrm{Cl^-})$ (1 : 1), so $[\mathrm{Cl^-}] = \dfrac{1.85 \times 10^{-3}}{0.0200} = 0.0925$ mol/L, or $0.0925 \times 35.45 = 3.28$ g/L.

**Percentage of a metal from its precipitate.** 1.50 g of $\mathrm{PbI_2}$ (461.0 g/mol) is $3.25 \times 10^{-3}$ mol, which holds $3.25 \times 10^{-3} \times 207.2 = 0.674$ g of lead. From a 25.0 g sample, that is 2.70%.

**Why a test can miss an ion.** A precipitate only forms if $Q > K_{sp}$. 5 ppm of lead is far too little for $\mathrm{PbCl_2}$ (which is fairly soluble) to come out of solution, but AAS can still detect it.

## Quick summary

- Ions are monitored because heavy metals are toxic and nutrients cause algal blooms. 1 ppm = 1 mg/L.
- Identify cations with precipitates, flame tests and complexes; anions with acid, $\mathrm{Ba^{2+}}$, $\mathrm{Ag^+}$ and molybdate.
- Gravimetric analysis: precipitate with excess reagent, filter, wash, dry to constant mass, weigh.
- Colorimetry and UV-vis: choose the absorbed wavelength, use a blank and standards, read the calibration curve.
- AAS: flame makes atoms, hollow cathode lamp gives element-specific light. Stay inside the calibration range.

## Quiz

Try these questions from the chapter quiz.

1. A concentration of 1 ppm in water equals about:
   - A. 1 g/L
   - B. 1 mg/L
   - C. 1 kg/L
   - D. 1 mol/L
   - Answer: B. 1 mg/L. One part per million of a litre (1 kg) is 1 mg.

2. Which ion gives a brick-red flame?
   - A. Sodium
   - B. Potassium
   - C. Calcium
   - D. Copper
   - Answer: C. Calcium. Calcium's flame is brick red.

3. Adding NaOH to a solution gives a blue precipitate. The ion is:
   - A. $\mathrm{Cu^{2+}}$
   - B. $\mathrm{Fe^{3+}}$
   - C. $\mathrm{Mg^{2+}}$
   - D. $\mathrm{Na^+}$
   - Answer: A. $\mathrm{Cu^{2+}}$. Copper(II) hydroxide is blue.

The whole chapter quiz is on the page, for signed-in students: https://whizznotes.com/chemistry/analysing-inorganic-substances/#quiz