Analysis of Inorganic Substances

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 $P b^{2 +}$ , mercury $H g^{2 +}$ , cadmium $C d^{2 +}$	Mining, old paint, batteries, industry	Toxic heavy metals; build up in food chains
Phosphate $P O_{4}^{3 -}$ , nitrate $N O_{3}^{-}$	Fertiliser run-off, sewage	Algal blooms (eutrophication) that use up oxygen and kill fish
Salt ( $N a^{+}$ , $C l^{-}$ )	Land clearing, irrigation	Salinity damages crops and soil
Copper $C u^{2 +}$	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 $25 mg \div 0.500 L = 50$ ppm.

Testing for ions (qualitative analysis)

Cations: precipitation tests

Add	Result	Ion present
Dilute HCl	White precipitate	$A g^{+}$ ( $AgCl$ ) or $P b^{2 +}$ ( $PbC l_{2}$ )
Dilute $H_{2} S O_{4}$	White precipitate	$B a^{2 +}$ , $C a^{2 +}$ or $P b^{2 +}$
NaOH	Blue precipitate	$C u^{2 +}$
NaOH	Green precipitate	$F e^{2 +}$
NaOH	Red-brown precipitate	$F e^{3 +}$
NaOH	White precipitate	$M g^{2 +}$ , $C a^{2 +}$ or $P b^{2 +}$

Testing in a sequence (a flow chart) separates the possibilities step by step: add HCl first to remove $A g^{+}$ and $P b^{2 +}$ , then $H_{2} S O_{4}$ , then NaOH, and so on.

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

Ion	Flame colour
$N a^{+}$	Yellow-orange
$K^{+}$	Lilac
$C a^{2 +}$	Brick red
$B a^{2 +}$	Pale green
$C u^{2 +}$	Blue-green
$S r^{2 +}$	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):

$F e^{3 +}$ + thiocyanate $SC N^{-}$ → blood-red $FeSC N^{2 +}$ .
$C u^{2 +}$ + excess ammonia → deep blue $[Cu (N H_{3})_{4}]^{2 +}$ .

Anions

Ion	Test	Result
Carbonate $C O_{3}^{2 -}$	Add dilute acid	Fizzes; gas turns limewater milky
Sulfate $S O_{4}^{2 -}$	Add $B a^{2 +}$ (with HCl present)	White precipitate of $BaS O_{4}$
Chloride $C l^{-}$	Add $AgN O_{3}$ (with $HN O_{3}$ )	White precipitate, dissolves in dilute ammonia
Bromide $B r^{-}$	Add $AgN O_{3}$	Cream precipitate
Iodide $I^{-}$	Add $AgN O_{3}$	Yellow precipitate
Phosphate $P O_{4}^{3 -}$	Add ammonium molybdate (with acid)	Yellow precipitate
Hydroxide $O H^{-}$	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.

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

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, $A g^{+} + C l^{-} \to AgCl (s)$ . A little chromate indicator gives a red-brown precipitate of $A g_{2} Cr O_{4}$ 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.

Choose the wavelength (or filter colour) that the solution absorbs most. This is the complementary colour: a blue solution absorbs orange-red light.
Zero the instrument with a blank (the solvent, usually distilled water).
Measure the absorbance of several standard solutions of known concentration.
Plot a calibration curve (absorbance against concentration). It should be a straight line through the origin.
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.

The sample solution is sprayed into a flame, which turns it into free atoms.
A hollow cathode lamp made of the metal being tested gives out light of exactly the wavelengths that those atoms absorb.
The atoms absorb some of this light. A detector measures how much.
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.

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 12The Beer-Lambert law and harder gravimetric problems

Beer-Lambert law. $A = εl c$ , where $A$ is absorbance, $ε$ is the molar absorptivity, $l$ is the path length (usually 1 cm) and $c$ is the concentration. With $ε = 1.2 \times 1 0^{4}$ L mol⁻¹ cm⁻¹, $l = 1.0$ cm and $A = 0.60$ : $c = 0.60 \div 1.2 \times 1 0^{4} = 5.0 \times 1 0^{- 5}$ mol/L.

Absorbance is also written $A = lo g_{10} \frac{I _{in}}{I _{out}}$ , where $I_{in}$ is the light going in and $I_{out}$ 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 $AgCl$ (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 $AgN O_{3}$ . $n (A g^{+}) = 1.85 \times 1 0^{- 3}$ mol $= n (C l^{-})$ (1 : 1), so $[C l^{-}] = \frac{1.85 \times 1 0 ^{- 3}}{0.0200} = 0.0925$ mol/L, or $0.0925 \times 35.45 = 3.28$ g/L.

Percentage of a metal from its precipitate. 1.50 g of $Pb I_{2}$ (461.0 g/mol) is $3.25 \times 1 0^{- 3}$ mol, which holds $3.25 \times 1 0^{- 3} \times 207.2 = 0.674$ 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 $Q > K_{s p}$ . 5 ppm of lead is far too little for $PbC l_{2}$ (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, $B a^{2 +}$ , $A g^{+}$ 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

14 questions with instant marking. It is free, and your scores are saved to your account so you can carry on from any device.

Turn on JavaScript to sign in and take the quiz.

Try 3 questions now

A concentration of 1 ppm in water equals about:
1. A 1 g/L
2. B 1 mg/L
3. C 1 kg/L
4. D 1 mol/L
Show answer
B: 1 mg/L
One part per million of a litre (1 kg) is 1 mg.
Which ion gives a brick-red flame?
1. A Sodium
2. B Potassium
3. C Calcium
4. D Copper
Show answer
C: Calcium
Calcium's flame is brick red.
Adding NaOH to a solution gives a blue precipitate. The ion is:
1. A $C u^{2 +}$
2. B $F e^{3 +}$
3. C $M g^{2 +}$
4. D $N a^{+}$
Show answer
A: $C u^{2 +}$
Copper(II) hydroxide is blue.