ionic compound

Ionic bonding model.

When metallic and non-metallic atoms form ionic compounds

• Metal atoms lose electrons and become cations.
• Non-metal atoms gain electrons so become anions.
• The ionic bonds form an ionic lattice. (Like the structure of salt.)

The properties of ionic compounds include

• High melting points - this proves the ionic bonds between atoms are very strong.
• Hardness - the lattice structure cannot be scratched easily.
• Brittleness - the lattice structure can shatter.

When an ionic compound in liquid state, the charged ions can move - this means that it can conduct electricity.

The tendency for elements to react so their atoms have 8 valence electrons is the octet rule.

formula writing

Ionic compounds are not charged. So their components must cancel the positive and negative charges.

For example, 1 $\mathrm{A}{\mathrm{l}}^{3+}$ and 3 ${\mathrm{F}}^{-}$ make the ionic compound $\mathrm{AL}{\mathrm{F}}_3$.
For example, 3 $\mathrm{C}{\mathrm{u}}^{2+}$ and 2 ${\mathrm{N}}^{3-}$ make the ionic compound $\mathrm{C}{\mathrm{u}}_3{\mathrm{N}}_2$.

An ion is a charged atom, but a polyatomic ion is a charged molecule.

Common cations list

• caesium, Cs+
• copper (I), Cu+
• gold (I), Au+
• lithium (I), Li+
• potassium, K+
• rubidium, Rb+
• silver, Ag+
• sodium, Na+

Common polyatomic cations list

• ammonium, $\mathrm{N}{\mathrm{H}}_4{}^{+}$
• hydronium, ${\mathrm{H}}_3\mathrm{O}{}^{+}$
• mercury (I), $\mathrm{H}{\mathrm{g}}_2{}^{2+}$
• phosphonium, $\mathrm{P}{\mathrm{H}}_4{}^{+}$

Common anions list

• bromide, Br-
• chloride, Cl-
• fluoride, F-
• iodide, I-
• oxide, ${\mathrm{O}}^{2-}$
• sulfide, ${\mathrm{S}}^{2-}$
• nitride, ${\mathrm{N}}^{3-}$

Common polyatomic anions list

• carbonate, $\mathrm{C}{\mathrm{O}}_3{}^{2-}$
• hydroxide, $\mathrm{O}{\mathrm{H}}^{-}$
• nitrate, $\mathrm{N}{\mathrm{O}}_3{}^{-}$
• nitrite, $\mathrm{N}{\mathrm{O}}_2{}^{-}$
• phosphate, $\mathrm{P}{\mathrm{O}}_4{}^{3-}$
• sulfate, $\mathrm{S}{\mathrm{O}}_4{}^{2-}$
• sulfite, $\mathrm{S}{\mathrm{O}}_3{}^{2-}$

Naming ions and polyatomic ions.

• The name of a cation is the name of the metal plus the word ion. Example, sodium ion.
• The name of an anion is the name of the non-metal plus the suffix “-ide”. Example, chloride.
• The name of a polyatomic cation often has “-ium”.
• The name of a polyatomic anion often has “-ate” or “-ite”.

Writing a chemical formula.

• Cation first.
• Use subscripts to show the number of ions.

Umm here is something might help…

Okay, let’s answer these questions about ionic compounds, precipitation reactions, and spectator ions, providing detailed explanations for each step.

Question 1: Dissolution Equations

This question asks us to write equations showing the dissociation of ionic compounds into their constituent ions when dissolved in water. We’ve already covered this in a previous turn. Here are the solutions:

• a) Potassium Bromide (KBr):

    - KBr(s) → K⁺(aq) + Br⁻(aq)
  

• b) Calcium Nitrate (Ca(NO₃)₂):

    - Ca(NO₃)₂(s) → Ca²⁺(aq) + 2NO₃⁻(aq)
  

• c) Sodium Sulfide (Na₂S):

    - Na₂S(s) → 2Na⁺(aq) + S²⁻(aq)
  

• d) Iron(III) Chloride (FeCl₃):

    - FeCl₃(s) → Fe³⁺(aq) + 3Cl⁻(aq)
  

• e) Aluminium Sulfate (Al₂(SO₄)₃):

    - Al₂(SO₄)₃(s) → 2Al³⁺(aq) + 3SO₄²⁻(aq)
  

Question 2: Precipitation Reactions and Spectator Ions

This question has two parts: predicting precipitates and identifying spectator ions.

2a) Predicting Precipitates

We’ll use solubility rules to determine if a precipitate forms. The general approach is:

1. Identify the ions present in the reactants.
2. Consider the possible combinations of cations and anions to form new compounds.
3. Use solubility rules to determine if any of these new compounds are insoluble (precipitates).

Here’s the breakdown for each reaction:

• i) Ca(NO₃)₂ and K₂CO₃:

    - Ions: Ca²⁺, NO₃⁻, K⁺, CO₃²⁻
    - Possible Products: KNO₃ and CaCO₃
    - Solubility: KNO₃ is soluble (all nitrates are soluble). CaCO₃ is _insoluble_ (most carbonates are insoluble except those of Group 1 and NH₄⁺).
    - **Precipitate Formula: CaCO₃**
  

• ii) MgSO₄ and Cu(NO₃)₂:

    - Ions: Mg²⁺, SO₄²⁻, Cu²⁺, NO₃⁻
    - Possible Products: Mg(NO₃)₂ and CuSO₄
    - Solubility: Both Mg(NO₃)₂ and CuSO₄ are soluble.
    - **Precipitate Formula: No Precipitate (NP)**
  

• iii) Na₂S and MgSO₄:

    - Ions: Na⁺, S²⁻, Mg²⁺, SO₄²⁻
    - Possible Products: Na₂SO₄ and MgS
    - Solubility: Na₂SO₄ is soluble. MgS is generally considered _insoluble_ (most sulfides are insoluble). However, it is slightly soluble and would depend on the concentration. For this purpose we will class is as insoluble.
    - **Precipitate Formula: MgS**
  

• iv) FeCl₃ and NH₄OH:

    - Ions: Fe³⁺, Cl⁻, NH₄⁺, OH⁻
    - Possible Products: NH₄Cl and Fe(OH)₃
    - Solubility: NH₄Cl is soluble (ammonium salts are soluble). Fe(OH)₃ is _insoluble_ (most hydroxides are insoluble).
    - **Precipitate Formula: Fe(OH)₃**
  

• v) Na₃PO₄ and AgNO₃:

    - Ions: Na⁺, PO₄³⁻, Ag⁺, NO₃⁻
    - Possible Products: NaNO₃ and Ag₃PO₄
    - Solubility: NaNO₃ is soluble (all nitrates and Group 1 salts are soluble). Ag₃PO₄ is _insoluble_ (most phosphates are insoluble).
    - **Precipitate Formula: Ag₃PO₄**
  

2b) Spectator Ions

Spectator ions are ions that are present in the solution before and after the reaction but do not participate in the formation of the precipitate. They remain dissolved in solution. To find them:

4. Write the complete ionic equation: Show all dissolved ions as separate species.
5. Identify the ions that appear on both sides of the equation unchanged.

Let’s go through each reaction where a precipitate formed:

• i) Ca(NO₃)₂(aq) + K₂CO₃(aq) → CaCO₃(s) + 2KNO₃(aq)

    - Complete Ionic Equation: Ca²⁺(aq) + 2NO₃⁻(aq) + 2K⁺(aq) + CO₃²⁻(aq) → CaCO₃(s) + 2K⁺(aq) + 2NO₃⁻(aq)
    - **Spectator Ions: K⁺ and NO₃⁻**
  

• iii) Na₂S(aq) + MgSO₄(aq) → MgS(s) + Na₂SO₄(aq)

    * Complete Ionic Equation: 2Na+(aq) + S²⁻(aq) + Mg²⁺(aq) + SO₄²⁻(aq) → MgS(s) + 2Na+(aq) + SO₄²⁻(aq)
    
    * Spectator Ions: Na⁺ and SO₄²⁻
    
  

• iv) FeCl₃(aq) + 3NH₄OH(aq) → Fe(OH)₃(s) + 3NH₄Cl(aq)

    - Complete Ionic Equation: Fe³⁺(aq) + 3Cl⁻(aq) + 3NH₄⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) + 3NH₄⁺(aq) + 3Cl⁻(aq)
    - **Spectator Ions: NH₄⁺ and Cl⁻**
  

• v) Na₃PO₄(aq) + 3AgNO₃(aq) → Ag₃PO₄(s) + 3NaNO₃(aq)

    - Complete Ionic Equation: 3Na⁺(aq) + PO₄³⁻(aq) + 3Ag⁺(aq) + 3NO₃⁻(aq) → Ag₃PO₄(s) + 3Na⁺(aq) + 3NO₃⁻(aq)
    - **Spectator Ions: Na⁺ and NO₃⁻**
  

Question 3: Naming Precipitates and Balanced Equations

3a) Naming Precipitates

• i) K₂S and MgCl₂:

    - Possible Products: KCl and MgS
    - KCl is soluble. MgS is _insoluble_.
    - **Precipitate Name: Magnesium Sulfide**
  

• ii) CuCl₂ and AgNO₃:

    - Possible Products: Cu(NO₃)₂ and AgCl
    - Cu(NO₃)₂ is soluble. AgCl is _insoluble_.
    - **Precipitate Name: Silver Chloride**
  

• iii) KOH and AlCl₃:

    - Possible Products: KCl and Al(OH)₃
    - KCl is soluble. Al(OH)₃ is _insoluble_.
    - **Precipitate Name: Aluminum Hydroxide**
  

• iv) MgSO₄ and NaOH:

    - Possible Products: Mg(OH)₂ and Na₂SO₄
    - Na₂SO₄ is soluble. Mg(OH)₂ is _insoluble_.
    - **Precipitate Name: Magnesium Hydroxide**
  

3b) Balanced Full Chemical Equations

• i) K₂S(aq) + MgCl₂(aq) → MgS(s) + 2KCl(aq)

• ii) CuCl₂(aq) + 2AgNO₃(aq) → 2AgCl(s) + Cu(NO₃)₂(aq)

• iii) 3KOH(aq) + AlCl₃(aq) → Al(OH)₃(s) + 3KCl(aq)

• iv) MgSO₄(aq) + 2NaOH(aq) → Mg(OH)₂(s) + Na₂SO₄(aq)

These equations are all balanced, showing the correct stoichiometric ratios of reactants and products. The states of matter (s = solid, aq = aqueous) are also included.