Identifying Cations in Solution (HSC SSCE Chemistry): Revision Notes

Identifying Cations in Solution

What are cations?

Cations are positively charged ions. Most cations are metal ions, such as barium ($\text{Ba}^{2+}$), calcium ($\text{Ca}^{2+}$), or iron(II) ($\text{Fe}^{2+}$).

In qualitative analysis, you need to identify the presence of specific cations in aqueous solutions. The eight cations covered in this topic are:

• Barium ($\text{Ba}^{2+}$)
• Calcium ($\text{Ca}^{2+}$)
• Magnesium ($\text{Mg}^{2+}$)
• Lead(II) ($\text{Pb}^{2+}$)
• Silver ($\text{Ag}^+$)
• Copper(II) ($\text{Cu}^{2+}$)
• Iron(II) ($\text{Fe}^{2+}$)
• Iron(III) ($\text{Fe}^{3+}$)

Methods for identifying cations

There are three main methods used to identify cations in solution:

1. Flame tests - observing the color produced when metal ions are heated in a flame
2. Precipitation reactions - observing whether precipitates form when specific anions are added
3. Complexation reactions - observing color changes or precipitate dissolution when complex ions form

Flame tests

Understanding the theory

Flame tests work because of electron behavior in atoms. Each element has its own unique arrangement of energy levels where electrons exist. When you heat an atom, the electrons absorb energy and jump up to higher energy levels. This is called the excited state.

Electrons in excited states are unstable and quickly return to their original positions (the ground state) in less than one-millionth of a second. As they fall back down, they release the absorbed energy as light.

When the emitted light falls in the visible region of the electromagnetic spectrum, we see a distinctive color. This color can identify the element.

Conducting flame tests

In a flame test, you place a sample in a high-energy flame from a Bunsen burner or blowtorch. Some metal ions gain electrons to become neutral atoms, and it's the electrons in these neutral atoms that produce the colored flame as they return to ground state.

Flame colors of cations

Here are the characteristic flame colors for the eight cations you need to know:

Ion	Flame Color
$\text{Ca}^{2+}$	Brick-red (orange-red)
$\text{Ba}^{2+}$	Pale green (apple green)
$\text{Cu}^{2+}$	Blue (with halides); green (with other compounds)
$\text{Pb}^{2+}$	Light blue-grey
$\text{Fe}^{2+}$	Gold when very hot; bright blue or green turning to orange-brown
$\text{Fe}^{3+}$	Orange-brown

Limitations of flame tests

Precipitation reactions

Overview of precipitation tests

Precipitation reactions are the primary method for identifying cations. Different cations form precipitates of characteristic colors when specific anions are added to their solutions. By systematically testing with various reagents, you can identify which cation is present.

The main test reagents used are:

• Hydroxide ion ($\text{OH}^-$) - usually as sodium hydroxide (NaOH)
• Chloride ion ($\text{Cl}^-$) - usually as sodium chloride (NaCl)
• Iodide ion ($\text{I}^-$) - usually as potassium iodide (KI)
• Fluoride ion ($\text{F}^-$) - usually as sodium fluoride (NaF)
• Sulfate ion ($\text{SO}_4^{2-}$) - usually as sulfuric acid ($\text{H}_2\text{SO}_4$)
• Permanganate ion ($\text{MnO}_4^-$) - as potassium permanganate
• Thiocyanate ion ($\text{SCN}^-$) - as potassium thiocyanate (KSCN)
• Ammonia solution ($\text{NH}_3$)

Test results for each cation

The following table summarizes the key tests and observations for each cation:

Cation	OH⁻ Test	Cl⁻ Test	SO₄²⁻ Test	Other Tests	Flame Test
$\text{Ag}^+$	Brown (milk coffee colored) precipitate	White precipitate that dissolves in ammonia	No precipitate	With $\text{I}^-$: pale yellow precipitate	No color
$\text{Pb}^{2+}$	White precipitate that dissolves in excess $\text{OH}^-$	White precipitate (doesn't dissolve in ammonia)	White precipitate	With $\text{I}^-$: yellow precipitate	Light blue-grey
$\text{Ba}^{2+}$	No precipitate	No precipitate	White precipitate	No precipitate with $\text{F}^-$	Pale green
$\text{Ca}^{2+}$	White precipitate (if concentrated)	No precipitate	White precipitate (if concentrated)	With $\text{F}^-$: white precipitate	Brick-red
$\text{Mg}^{2+}$	White precipitate	No precipitate	No precipitate	With $\text{F}^-$: white precipitate	No color
$\text{Cu}^{2+}$	Blue precipitate; dissolves in $\text{NH}_3$ to form deep blue solution	No precipitate	No precipitate	-	Blue-green
$\text{Fe}^{2+}$	Green or white precipitate (may turn brown)	No precipitate	No precipitate	Decolorizes acidified $\text{KMnO}_4$	Gold (very hot) or blue-green to orange-brown
$\text{Fe}^{3+}$	Brown precipitate	No precipitate	No precipitate	With $\text{SCN}^-$: deep red solution	Orange-brown

Systematic identification flowchart

When a solution contains only one cation, follow this systematic approach:

The flowchart shows a logical sequence of tests:

Step 1: Start by adding 2 drops of NaOH

• Colored precipitate → test further to distinguish $\text{Fe}^{2+}$ (green), $\text{Fe}^{3+}$ (orange-brown), or $\text{Cu}^{2+}$ (blue)
• White precipitate → add excess NaOH to see if it dissolves ($\text{Pb}^{2+}$) or remains ($\text{Ag}^+$, $\text{Ca}^{2+}$, or $\text{Mg}^{2+}$)
• No precipitate → could be $\text{Ba}^{2+}$ or low concentration $\text{Ca}^{2+}$

Step 2: For colored precipitates, use additional tests:

• Green precipitate + 2 drops NaCl: no change confirms $\text{Fe}^{2+}$; white precipitate indicates $\text{Ag}^+$
• Orange-brown precipitate confirms $\text{Fe}^{3+}$
• Blue precipitate confirms $\text{Cu}^{2+}$

Step 3: For white precipitates, add excess NaOH:

• Precipitate dissolves = $\text{Pb}^{2+}$ (confirm with $\text{H}_2\text{SO}_4$ - white precipitate)
• Precipitate remains = test with NaCl or $\text{H}_2\text{SO}_4$

Step 4: For no precipitate with NaOH, test with:

• NaF: white precipitate = $\text{Ca}^{2+}$; no precipitate = $\text{Ba}^{2+}$
• Or use $\text{H}_2\text{SO}_4$: white precipitate = $\text{Ba}^{2+}$ or $\text{Ca}^{2+}$ (distinguish with flame test); no precipitate = $\text{Mg}^{2+}$

Confirmatory tests

After identifying a probable cation using the flowchart, always confirm with an additional test:

Laboratory investigation tips

When conducting Investigation 14.2 to identify unknown cations:

Complexation reactions

What are complex ions?

A complex ion (or simply a complex) forms when one or more small molecules or ions attach themselves to a central cation. The central cation is usually (but not always) a transition metal ion. The resultant complex has different properties - including different colors and solubility - compared to the simple cation.

The surrounding molecules and ions are called ligands. Ligands must contain at least one lone pair of electrons.

Coordinate covalent bonds

The bond between a metal ion and a ligand is special - it's called a coordinate covalent bond (or dative bond). In this type of bond, both electrons come from the ligand (the electron pair donor), rather than one electron from each atom as in a normal covalent bond.

Coordinate covalent bonds form when one atom (like $\text{H}^+$) doesn't have a complete outer shell and bonds to another atom that has a full outer shell with at least one unshared electron pair.

Copper complexes

When copper(II) salts dissolve in water, they produce a blue solution due to formation of the hexaaquacopper(II) complex, $[\text{Cu(H}_2\text{O)}_6]^{2+}$:

Reaction with hydroxide: When sodium hydroxide is added to a copper(II) solution, a pale blue precipitate forms. Two hydroxide ions replace two water molecules:

$[\text{Cu(H}_2\text{O)}_6]^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow [\text{Cu(H}_2\text{O)}_4(\text{OH})_2](s) + 2\text{H}_2\text{O}(l)$

This is commonly written more simply as:

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

Reaction with ammonia: When ammonia solution is added to an aqueous copper(II) solution, water ligands are replaced by ammonia molecules, forming a deep blue solution:

$[\text{Cu(H}_2\text{O)}_6]^{2+}(aq) + 4\text{NH}_3(aq) \rightleftharpoons [\text{Cu(H}_2\text{O)}_2(\text{NH}_3)_4]^{2+}(aq) + 4\text{H}_2\text{O}(l)$

If ammonia is added to the pale blue copper hydroxide precipitate, the precipitate dissolves and forms the same deep blue complex:

$\text{Cu(OH)}_2(s) + 4\text{NH}_3(aq) \rightleftharpoons [\text{Cu(NH}_3)_4]^{2+}(aq) + 2\text{OH}^-(aq)$

Iron complexes

Iron(II) complexes: When iron(II) salts dissolve in water, they form a pale green solution due to hexaaquairon(II), $[\text{Fe(H}_2\text{O)}_6]^{2+}$.

Adding hydroxide ions produces a green (or sometimes white) precipitate:

$[\text{Fe(H}_2\text{O)}_6]^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow [\text{Fe(H}_2\text{O)}_4(\text{OH})_2](s) + 2\text{H}_2\text{O}(l)$

Simplified:

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

Iron(III) complexes: When iron(III) salts dissolve in water, they form a yellow solution due to hexaaquairon(III), $[\text{Fe(H}_2\text{O)}_6]^{3+}$.

Adding hydroxide ions produces a brown precipitate:

$[\text{Fe(H}_2\text{O)}_6]^{3+}(aq) + 3\text{OH}^-(aq) \rightarrow [\text{Fe(H}_2\text{O)}_3(\text{OH})_3](s) + 3\text{H}_2\text{O}(l)$

Simplified:

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

Thiocyanate test for iron(III): Adding thiocyanate ions ($\text{SCN}^-$) to an iron(III) solution produces a distinctive blood red color due to the complex $\text{Fe(SCN)}^{2+}$:

$[\text{Fe(H}_2\text{O)}_6]^{3+}(aq) + \text{SCN}^-(aq) \rightleftharpoons [\text{Fe(H}_2\text{O)}_5\text{SCN}]^{2+}(aq) + \text{H}_2\text{O}(l)$

Lead and silver complexes

Lead(II): Lead(II) hydroxide precipitate dissolves in excess hydroxide to form a complex ion:

$\text{Pb(OH)}_2(s) + 2\text{OH}^-(aq) \rightleftharpoons [\text{Pb(OH)}_4]^{2-}(aq)$

Silver: Although silver is not a transition metal, it does form complexes. Solid silver chloride dissolves in ammonia solution to produce a linear complex ion:

$\text{AgCl}(s) + 2\text{NH}_3(aq) \rightleftharpoons [\text{Ag(NH}_3)_2]^+(aq) + \text{Cl}^-(aq)$