Identifying the products of chemical reactions

In chemistry, reactions often happen invisibly. Two clear liquids are mixed, a change occurs, and we need to figure out exactly what new substances have been made. Whether you are identifying a gas released in a test tube, identifying dissolved ions in water, or interpreting a chart from a high-tech instrument, you are acting as a chemical detective.

In these notes, you will learn the qualitative chemical tests used to identify gases and ions, how to carry out flame tests, and how modern instrumental techniques are used to identify compounds.

What you'll learn

How to test for and identify common gases (oxygen, hydrogen, carbon dioxide, and chlorine)
How to use flame tests and precipitation reactions to identify metal cations
How to carry out chemical tests for carbonates, sulfates, and halide anions
How to evaluate the advantages of modern instrumental analysis and interpret a mass spectrum chart

Prerequisites & Core Concepts

Before diving into the tests, we need to clarify what we are trying to identify. Many of the reactions we use to identify substances take place in aqueous solutions.

Definition

Cations and Anions

Cation: A positively charged ion (usually a metal ion, such as $Na+\text{Na}^+$ or $Ca2+\text{Ca}^{2+}$ , or the ammonium ion $NH4+\text{NH}_4^+$ ).
Anion: A negatively charged ion (such as a halide ion $Cl−\text{Cl}^-$ , carbonate $CO32−\text{CO}_3^{2-}$ , or sulfate $SO42−\text{SO}_4^{2-}$ ).

Many of our chemical tests rely on forming a precipitate.

Definition

Precipitate

An insoluble solid that forms and settles out of a solution during a chemical reaction between two soluble substances.

Common Mistake

Mass vs Density in State Changes

When a precipitate forms, the mixture looks thicker and heavier, but the mass does not change! Total mass is always conserved in chemical reactions. Similarly, if a reaction produces a gas that escapes, the reading on a balance will go down, but the "lost" mass has simply escaped as gas into the air. The total mass of all reactants and products remains identical.

1. Identifying Gases

Gases are often colourless and escape quickly. You must know the standard laboratory tests for these four key gases. These tests are required for both Combined Science and separate GCSE Chemistry.

Oxygen ( $O2\text{O}_2$ )

Test: Insert a glowing splint (a wooden splint that has been lit and then blown out so it is only glowing red) into a test tube of the gas.
Result: The splint relights.

Hydrogen ( $H2\text{H}_2$ )

Test: Hold a lit splint at the open end of a test tube of the gas.
Result: The gas burns rapidly with a "squeaky pop" sound.

Carbon Dioxide ( $CO2\text{CO}_2$ )

Test: Bubble the gas through limewater (aqueous calcium hydroxide, $Ca(OH)2\text{Ca(OH)}_2$ ).
Result: The limewater turns cloudy (milky) because a white precipitate of insoluble calcium carbonate ( $CaCO3\text{CaCO}_3$ ) forms.

Chlorine ( $Cl2\text{Cl}_2$ )

Test: Hold a piece of damp blue litmus paper in the mouth of the test tube.
Result: The paper turns red momentarily (because chlorine gas dissolves to form an acidic solution) and is then bleached white.

2. Identifying Cations: Flame Tests

Note: The following sections on ion testing, flame tests, and instrumental methods are specific to the separate GCSE Chemistry course (J248) and are not tested in Combined Science.

Metal ions (cations) in solid salts or in solution can be identified by the unique colour they emit when heated in a flame.

How to Perform a Flame Test

To get a clear, unpolluted colour, you must follow a precise practical method:

Dip a clean nichrome or platinum wire loop into concentrated hydrochloric acid ( $HCl\text{HCl}$ ).
Heat the wire in the hot, blue flame of a Bunsen burner to clean off any impurities. Repeat this step until the wire no longer colours the flame.
Dip the clean wire back into the acid, and then touch it to a small sample of the solid metal salt so some sticks.
Place the wire loop back into the non-luminous (blue) Bunsen flame and observe the colour change.

Flame Colours to Memorise

Cation	Formula	Flame Colour
Lithium	$Li+\text{Li}^+$	Red (Crimson)
Sodium	$Na+\text{Na}^+$	Yellow
Potassium	$K+\text{K}^+$	Lilac
Calcium	$Ca2+\text{Ca}^{2+}$	Orange-red
Copper	$Cu2+\text{Cu}^{2+}$	Blue-green

Tip

Remembering potassium

A helpful way to remember potassium's flame is Potassium = Purple (Lilac).

3. Identifying Cations: Sodium Hydroxide Tests

If the metal salt is already in solution, we can identify the cation by adding sodium hydroxide solution ( $NaOH(aq)\text{NaOH(aq)}$ ). The hydroxide ions ( $OH−\text{OH}^-$ ) react with metal ions to form insoluble metal hydroxides, which appear as coloured precipitates.

The Metal Hydroxide Precipitates

Copper(II) ( $Cu2+\text{Cu}^{2+}$ ): Forms a blue precipitate.

\text{Cu}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Cu(OH)}_2\text{(s)}

Iron(II) ( $Fe2+\text{Fe}^{2+}$ ): Forms a slimy green precipitate.

\text{Fe}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_2\text{(s)}

Iron(III) ( $Fe3+\text{Fe}^{3+}$ ): Forms a red-brown (rust-like) precipitate.

\text{Fe}^{3+}\text{(aq)} + 3\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_3\text{(s)}

Precipitation of copper(II) hydroxide

The "White Precipitate" Problem

Two of the cations you need to know form white precipitates with sodium hydroxide: Calcium ( $Ca2+\text{Ca}^{2+}$ ) and Zinc ( $Zn2+\text{Zn}^{2+}$ ).

To tell them apart, you must add excess sodium hydroxide solution and shake:

Zinc hydroxide precipitate: Redissolves in excess sodium hydroxide to form a clear, colourless solution.
Calcium hydroxide precipitate: Does not redissolve; it remains as a white solid.

Key Idea

Zinc is amphoteric

Only zinc hydroxide dissolves in excess sodium hydroxide. If your white precipitate disappears when you add more alkali, the cation is $Zn2+\text{Zn}^{2+}$ . If it stays white and solid, it is $Ca2+\text{Ca}^{2+}$ .

Example

Identifying an unknown cation from sodium hydroxide tests

An unknown solution contains a single metal cation. A student adds a few drops of sodium hydroxide solution and observes a white precipitate. They continue to add sodium hydroxide until it is in excess, and the precipitate remains. Explain how the student can determine the exact identity of this metal ion.

Analyze the first observation: Adding a few drops of $NaOH(aq)\text{NaOH(aq)}$ produced a white precipitate. This narrows the identity of the cation down to either calcium ( $Ca2+\text{Ca}^{2+}$ ) or zinc ( $Zn2+\text{Zn}^{2+}$ ).
Analyze the second observation: The white precipitate did not dissolve when excess sodium hydroxide was added. Because zinc hydroxide dissolves in excess sodium hydroxide to form a colourless solution, the solid must be calcium hydroxide, which is insoluble in excess.
Select a confirmatory test: To confirm that the cation is calcium ( $Ca2+\text{Ca}^{2+}$ ) rather than another ion that forms an insoluble white hydroxide, the student should perform a flame test on the sample. A positive result for calcium will yield an orange-red flame.

4. Identifying Anions (Negative Ions)

Anions require different reagents to identify them. You must learn the specific tests for carbonates, sulfates, and halides.

Carbonate Ions ( $CO32−\text{CO}_3^{2-}$ )

Test: Add dilute acid (such as dilute hydrochloric acid) to the solid or solution.
Observation: The mixture will fizz (effervesce) as carbon dioxide gas is produced.
Confirmation: Bubble the gas produced through limewater to confirm it is $CO2\text{CO}_2$ (the limewater turns cloudy).

\text{CO}_3^{2-}\text{(aq)} + 2\text{H}^+\text{(aq)} \rightarrow \text{CO}_2\text{(g)} + \text{H}_2\text{O}\text{(l)}

Sulfate Ions ( $SO42−\text{SO}_4^{2-}$ )

Test: Add dilute hydrochloric acid, followed by barium chloride solution ( $BaCl2(aq)\text{BaCl}_2\text{(aq)}$ ).
Observation: A white precipitate of barium sulfate ( $BaSO4\text{BaSO}_4$ ) forms.

\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}

Halide Ions ( $Cl−\text{Cl}^-$ , $Br−\text{Br}^-$ , $I−\text{I}^-$ )

Test: Add dilute nitric acid ( $HNO3\text{HNO}_3$ ), followed by silver nitrate solution ( $AgNO3(aq)\text{AgNO}_3\text{(aq)}$ ).
Observations:
- Chloride ( $Cl−\text{Cl}^-$ ): Forms a white precipitate of silver chloride ( $AgCl\text{AgCl}$ ).
- Bromide ( $Br−\text{Br}^-$ ): Forms a cream precipitate of silver bromide ( $AgBr\text{AgBr}$ ).
- Iodide ( $I−\text{I}^-$ ): Forms a yellow precipitate of silver iodide ( $AgI\text{AgI}$ ).

Tip

Halide colour mnemonic

Halide precipitate colours get darker as you go down Group 7:

Chlorine is at the top $→\rightarrow$ White (like milk)
Bromine is in the middle $→\rightarrow$ Cream (like butter)
Iodine is further down $→\rightarrow$ Yellow (like cheese)

Common Mistake

Why do we add acid first?

You must add dilute acid (hydrochloric acid for the sulfate test, nitric acid for the halide test) before adding the testing reagent. The acid reacts with and removes any carbonate impurities ( $CO32−\text{CO}_3^{2-}$ ) that might be present. If you do not add acid, the carbonate ions will react with the barium or silver ions to form barium carbonate or silver carbonate, both of which are white solids. This would give you a false-positive result!

Example

Deducing the chemical formula of an unknown salt

A solid white salt, $Y$ , was dissolved in distilled water. The following tests were carried out on the resulting solution:

Test 1: A flame test produced a blue-green flame.
Test 2: Dilute nitric acid was added to the solution, followed by silver nitrate solution. A cream-coloured precipitate formed. Deduce the chemical formula of salt $Y$ .

Identify the cation: The blue-green flame in Test 1 is characteristic of copper(II) ions ( $Cu2+\text{Cu}^{2+}$ ).
Identify the anion: In Test 2, acidification with nitric acid followed by silver nitrate solution yielded a cream precipitate. A cream precipitate with silver nitrate is the diagnostic test for bromide ions ( $Br−\text{Br}^-$ ).
Combine the ions to write the formula: To form a neutral ionic compound, the positive and negative charges must balance. Copper(II) has a charge of $+ 2$ , and bromide has a charge of $- 1$ . Therefore, two bromide ions are needed for every one copper(II) ion:

\text{Charges: } (+2) + 2 \times (-1) = 0

The chemical formula of salt $Y$ is $CuBr2\text{CuBr}_2$ (copper(II) bromide).

5. Instrumental Methods of Analysis

In modern laboratories, chemists often use machines (instruments) rather than manual test tubes to identify substances.

Definition

Instrumental Analysis

The use of scientific instruments to detect, identify, and quantify chemical substances in a sample.

Advantages of Instrumental Methods

Compared to manual chemical tests, instrumental methods offer three main advantages:

Sensitivity: They can detect and identify extremely tiny amounts of a substance (even parts per billion).
Accuracy: They are highly precise and less prone to human error, providing very reliable identification.
Speed: They can run tests rapidly and are easily automated, allowing hundreds of samples to be processed in a short time.

Interpreting Mass Spectrometry Charts

One common instrumental method is mass spectrometry. It is used to determine the relative molecular mass ( $M_r$ ) of a compound and help map out its structure.

A mass spectrometer vaporises and ionises molecules, then accelerates the ions through a magnetic field. The ions are separated based on their mass-to-charge ratio ( $m / z$ ). Because most ions detected have a $+ 1$ charge, the $m / z$ value on the x-axis directly corresponds to the relative mass of the fragment.

A typical mass spectrum chart

When interpreting a mass spectrum chart, focus on these key features:

The peaks: Each vertical line represents a fragment of the molecule that has hit the detector. The height of the peak shows its relative abundance.
The Molecular Ion Peak ( $M^+$ ): This is the peak furthest to the right on the spectrum (ignoring any tiny, negligible isotope peaks). This peak represents the intact molecule that has lost a single electron. The $m / z$ value of this peak gives the relative molecular mass ( $M_r$ ) of the compound.

Exam technique

In the exam

Learn the reagents precisely: Do not just write "add barium" or "add silver". You must specify the full reagent name, e.g., "acidified silver nitrate solution".
State the acid used: Remember to state that you acidify before testing for anions. Use nitric acid with silver nitrate, and hydrochloric acid with barium chloride. Mixing them up (e.g., adding hydrochloric acid to a halide test) introduces chloride ions, which will instantly ruin your test by forming a white precipitate of silver chloride!
State observations clearly: Write "forms a blue precipitate", not just "turns blue". A precipitate is a solid, and marking schemes require you to specify that a solid has formed.
Spot the zinc exception: If a question describes a white precipitate with sodium hydroxide that "redissolves in excess", it is always zinc ( $Zn2+\text{Zn}^{2+}$ ). If it does not redissolve, it is calcium ( $Ca2+\text{Ca}^{2+}$ ).
Read the molecular ion peak correctly: When asked to find the relative molecular mass from a mass spectrum, look for the peak furthest to the right and read its value on the x-axis ( $m / z$ ).

Self review

Check yourself

Which gas turns damp blue litmus paper red and then bleaches it white?
A flame test produces a lilac flame. What metal ion is present?
Why must you add dilute hydrochloric acid before adding barium chloride when testing for sulfate ions?
If you have an unknown sample that could be either calcium chloride or zinc chloride, what single chemical test can you do to tell them apart, and what would you observe for each?

Identifying the products of chemical reactions

In these notes, you will learn the qualitative chemical tests used to identify gases and ions, how to carry out flame tests, and how modern instrumental techniques are used to identify compounds.

What you'll learn

How to test for and identify common gases (oxygen, hydrogen, carbon dioxide, and chlorine)
How to use flame tests and precipitation reactions to identify metal cations
How to carry out chemical tests for carbonates, sulfates, and halide anions
How to evaluate the advantages of modern instrumental analysis and interpret a mass spectrum chart

Prerequisites & Core Concepts

Before diving into the tests, we need to clarify what we are trying to identify. Many of the reactions we use to identify substances take place in aqueous solutions.

Definition

Cations and Anions

Cation: A positively charged ion (usually a metal ion, such as $Na+\text{Na}^+$ or $Ca2+\text{Ca}^{2+}$ , or the ammonium ion $NH4+\text{NH}_4^+$ ).
Anion: A negatively charged ion (such as a halide ion $Cl−\text{Cl}^-$ , carbonate $CO32−\text{CO}_3^{2-}$ , or sulfate $SO42−\text{SO}_4^{2-}$ ).

Many of our chemical tests rely on forming a precipitate.

Definition

Precipitate

An insoluble solid that forms and settles out of a solution during a chemical reaction between two soluble substances.

Common Mistake

Mass vs Density in State Changes

1. Identifying Gases

Gases are often colourless and escape quickly. You must know the standard laboratory tests for these four key gases. These tests are required for both Combined Science and separate GCSE Chemistry.