Reactions Between Ionic Compounds & Precipitation Reactions

Reactions Between Ionic Compounds in Solution

  • In some cases, a reaction occurs when two solutions of ionic compounds are mixed.

  • For example, when a colorless solution of silver nitrate (AgNO_3) is mixed with a colorless solution of sodium chloride (NaCl), a cloudy solution results due to the formation of a fine-powdered white solid (precipitate).

Identifying Ions in Reactant Solutions

  • Silver nitrate solution contains dissolved silver ions (Ag^+) and nitrate ions (NO_3^-).

  • Sodium chloride solution contains dissolved sodium ions (Na^+) and chloride ions (Cl^-).

When the solutions are mixed, the resulting mixture will contain all four ions.

Ion Interactions and Precipitate Formation

  • All ions move independently within the mixture.

  • Ions collide, and if positive and negative ions join to form an insoluble compound, a precipitate may form.

  • In the mixture of sodium chloride and silver nitrate, two new combinations are possible:

    • Sodium and nitrate ions (forming sodium nitrate, NaNO_3)

    • Silver and chloride ions (forming silver chloride, AgCl)

  • Solubility tables are used to determine which compound is the precipitate.

Solubility Tables

  • Table 5.3.1 lists soluble ionic compounds – those where >0.1 mol dissolves per liter at 25°C.

    • Most chlorides (Cl^-), bromides (Br^-), and iodides (I^-) are soluble, except AgCl, AgBr, AgI, and PbI_2.

    • All nitrates (NO3^-), ammonium salts (NH4^+), sodium salts (Na^+), potassium salts (K^+), and ethanoates (CH_3COO^-) are soluble with no exceptions.

    • Most sulfates (SO4^{2-}) are soluble, except for SrSO4, BaSO4, and PbSO4.

  • Table 5.3.2 lists insoluble ionic compounds.

    • Most hydroxides (OH^-) are insoluble, except for NaOH, KOH, Ba(OH)2, Ca(OH)2, and Sr(OH)2. Slight solubility is noted for NH4OH & AgOH.

    • Most carbonates (CO3^{2-}), phosphates (PO4^{3-}), and sulfides (S^{2-}) are insoluble, except for Na2CO3,
      K2CO3, (NH4)2CO3, Na3PO4, K3PO4, (NH4)3PO4, Na2S, K2S, and (NH4)2S.

Determining the Precipitate

  • All nitrate and sodium-containing compounds are generally soluble in water.

  • Most chloride-containing compounds are soluble, but silver chloride (AgCl) is an exception.

  • Therefore, the precipitate in the reaction between silver nitrate and sodium chloride is silver chloride.

Pictorial Representation of Precipitation Reaction

  • When hydrated Ag^+ and Cl^- ions come into contact, an ionic lattice of AgCl forms.

  • Na^+ and NO_3^- ions remain soluble.

  • The attraction between Ag^+ and Cl^- ions form an ionic lattice of AgCl precipitate.

Determining the Precipitate: A Step-by-Step Approach

  • List the positive and negative ions of each compound: e.g., Ag^+, NO_3^-, Na^+, Cl^-.

  • Draw lines connecting the positive ion of one solution to the negative ion of the other.

  • Use solubility tables to determine which new combination of ions results in an insoluble compound (the precipitate).

    • The other ions remain in solution.

Predicting Products of Precipitation Reactions (Worked Example 5.3.1)

  • Question: What precipitate(s) will form if solutions of potassium hydroxide (KOH) and lead(II) nitrate (Pb(NO3)2) are mixed?

  • Ions Present: K^+(aq), OH^- (aq), Pb^{2+} (aq), NO_3^- (aq)

  • Possible Combinations: K^+ (aq) and NO_3^- (aq); Pb^{2+} (aq) and OH^- (aq)

  • Solubility Check:

    • Potassium nitrate is usually soluble (no precipitate).

    • Lead(II) hydroxide is usually insoluble (precipitate will form).

Writing Equations for Precipitation Reactions

  • The reaction between silver nitrate and sodium chloride can be summarized as:

    • silver nitrate solution + sodium chloride solution → silver chloride solid + sodium nitrate solution.

  • The full equation includes state symbols for each species:

    • AgNO3(aq) + NaCl(aq) ightharpoonup AgCl(s) + NaNO3(aq)

Spectator Ions

  • Sodium and nitrate ions (Na^+ and NO_3^-) are not combined in the product; they move freely through the solution.

  • They are present at the start and end of the reaction as separate ions.

  • Spectator ions undergo no chemical change; they start and remain as aqueous ions.

Writing Full Equations for Precipitation Reactions (Worked Example 5.3.2)

  • Question: Write a balanced equation for the reaction of iron(III) nitrate and sodium sulfide, where the precipitate is iron(III) sulfide. Identify the spectator ions.

  • Unbalanced Equation: Fe(NO3)3(aq) + Na2S(aq) ightharpoonup Fe2S_3(s)

  • Add Other Compound: Fe(NO3)3(aq) + Na2S(aq) ightharpoonup Fe2S3(s) + NaNO3(aq)

  • Balanced Equation: 2Fe(NO3)3(aq) + 3Na2S(aq) ightharpoonup Fe2S3(s) + 6NaNO3(aq)

  • Spectator Ions: Na^+(aq) and NO_3^-(aq)

Ionic Equations for Precipitation Reactions

  • The essential feature of the reaction is the combination of silver and chloride ions to form a precipitate.

  • This reaction can be summarized in an ionic equation.

  • Spectator ions are not included in an ionic equation.

  • Only the species that combine to form the precipitate are included.

  • An ionic equation can be thought of as a full equation with the spectator ions removed.

  • Full equation: AgNO3(aq) + NaCl(aq) ightharpoonup AgCl(s) + NaNO3(aq)

  • Ionic equation: Ag^+(aq) + Cl^-(aq)
    ightharpoonup AgCl(s)

Writing Ionic Equations (Worked Example 5.3.3)

  • Question: Write an ionic equation for the reaction of aluminum nitrate and sodium sulfide, producing aluminum sulfide.

  • Precipitate:
    ightharpoonup Al2S3

  • Add Ions: 2Al^{3+} + 3S^{2-}
    ightharpoonup Al2S3

  • Add Symbols of State and Balance: $$2Al^{3+}(aq) + 3S^{2-}(aq)
    ightharpoon


KEY TERMS

(R): Ratio of solute distance to solvent distance in chromatography.

(branch): Side chain off the main carbon chain.

Avogadro's constant: 6.022×10236.022 \times 10^{23} particles in one mole.

abundance (relative isotopic): % of an isotope in a natural element sample.

acid (weak): Partially ionises in water.

addition reaction: Reaction where atoms are added to a molecule.

adsorption: Particles sticking to a surface.

alcohol: Compound with a hydroxyl (–OH) group.

alkane: Saturated hydrocarbon, only single bonds.

alkene: Unsaturated hydrocarbon, has at least one double bond.

alkyl group: Carbon and hydrogen chain (e.g. –CH₃) from an alkane.

alkyl side chain (branch): Extra carbon chain attached to main one.

aqueous: Dissolved in water.

aqueous solution: Water is the solvent.

bio-derived: Made from living/biological sources.

bioethanol: Ethanol produced from plant matter.

boiling point: Temperature at which a liquid becomes a gas.

carboxyl group: Functional group –COOH.

carboxylate ion: –COO⁻, formed when a carboxyl group loses H⁺.

carboxylic acid: Organic acid with –COOH group.

chemical property: How a substance reacts.

chromatogram: Output from chromatography showing separated parts.

chromatography: Technique to separate mixture components.

combustion: Reaction with oxygen producing heat and light.

complete combustion: Combustion making only CO₂ and H₂O.

component: Part of a mixture.

condensed structural formula: Molecule written without full bonds (e.g. CH₃CH₂OH).

crude oil: Natural mix of hydrocarbons from underground.

desorption: Removal of particles from a surface.

dispersion forces: Weak forces between molecules (especially non-polar).

dimer: Molecule made from 2 identical parts.

dissociation: Ions separating from a compound in water.

dissolution: Process of a solute dissolving in a solvent.

empirical formula: Simplest whole-number ratio of atoms in a compound.

fractional distillation: Separating liquids by boiling points.

functional group: Group of atoms giving a molecule its reactivity.

general formula: Formula that applies to a whole group (e.g. alkanes = CnH2n+2).

haloalkane: Alkane with a halogen atom replacing hydrogen.

halogen: Group 17 elements (F, Cl, Br, I).

homogeneous: Evenly mixed throughout.

homologous series: Group of compounds with same functional group and pattern.

hydrated: Contains water molecules.

hydrocarbon: Compound made of hydrogen and carbon only.

hydrogen bonds: Attraction between H and O/N/F in different molecules.

hydronium ion: H₃O⁺, formed when H⁺ bonds to water.

hydroxyl group: –OH group in alcohols.

immiscible: Liquids that do not mix.

incomplete combustion: Burning with limited oxygen, makes CO or C (soot).

ion-dipole attractions: Forces between ions and polar molecules.

ionisation: Gaining or losing electrons to form ions.

ionised: Turned into ions.

isotope: Atoms of the same element with different numbers of neutrons.

locant: Number showing atom position in a molecule.

mass spectrometer: Machine to measure mass of atoms/molecules.

mass spectrum: Output graph from a mass spectrometer.

melting point: Temp at which solid becomes liquid.

miscible: Liquids that mix fully.

mobile phase: Moving part in chromatography (solvent).

molar mass: Mass of 1 mole of a substance (g/mol).

mole: 1 mole = 6.022×10236.022 \times 10^{23} particles.

molecular formula: Actual number of atoms of each element in a molecule.

molecule: Group of atoms bonded together.

non-polar: Molecules with no charge separation.

non-renewable: Resource that can’t be replaced quickly (e.g. fossil fuels).

organic chemistry: Study of carbon compounds.

organic compound: Compound mainly made of carbon.

origin (chromatography): Starting point where sample is placed.

paper chromatography: Uses paper as stationary phase.

percentage composition: % by mass of each element in a compound.

photosynthesis: Plants make glucose + O₂ from CO₂ + H₂O.

physical property: Can be measured without changing substance (e.g. melting point).

plant-sourced biomass: Plant material used as fuel or raw chemical source.

polar: Molecule with uneven charge distribution.

polymer: Long molecule made of repeating smaller units (monomers).

purity: How much of a substance is one single compound.

qualitative analysis: Tells what substances are present.

quantitative analysis: Tells how much of each substance is present.

relative atomic mass: Weighted average mass of an atom’s isotopes.

relative isotopic mass: Mass of an isotope compared to carbon-12.

relative mass: Mass of one particle compared to another.

renewable: Resource that is naturally replaced quickly.

retardation factor (R): See (R) – same thing.

saturated: Only single bonds between carbon atoms.

semi-structural formula: Shows connections, not all bonds (e.g. CH₃CH₂OH).

solute: Substance being dissolved.

solution: Solute dissolved in a solvent.

solvent: Substance doing the dissolving.

standards: Known samples used for comparison.

stationary phase: Non-moving part in chromatography.

stem name: Base name showing main carbon chain length (e.g. "meth-", "eth-").

structural formula: Shows full structure and bonding of a molecule.

structural isomer: Molecules with same formula but different structure.

substitution reaction: Atom/group is swapped for another in a molecule.

terminal carbon: End carbon of a chain.

thin-layer chromatography: Uses coated glass/plastic as stationary phase.

unsaturated: Has double or triple carbon-carbon bonds.