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 ($NO<em>3^-$), ammonium salts ($NH</em>4^+$), sodium salts ($Na^+$), potassium salts ($K^+$), and ethanoates ($CH_3COO^-$) are soluble with no exceptions.

  • Most sulfates ($SO<em>4^{2-}$) are soluble, except for $SrSO</em>4$, $BaSO<em>4$, and $PbSO</em>4$.

• Table 5.3.2 lists insoluble ionic compounds.

  • Most hydroxides ($OH^-$) are insoluble, except for $NaOH$, $KOH$, $Ba(OH)<em>2$, $Ca(OH)</em>2$, and $Sr(OH)<em>2$. Slight solubility is noted for $NH</em>4OH$ & $AgOH$.

  • Most carbonates ($CO<em>3^{2-}$), phosphates ($PO</em>4^{3-}$), and sulfides ($S^{2-}$) are insoluble, except for $Na<em>2CO</em>3$,
    $K<em>2CO</em>3$, $(NH<em>4)</em>2CO<em>3$, $Na</em>3PO<em>4$, $K</em>3PO<em>4$, $(NH</em>4)<em>3PO</em>4$, $Na<em>2S$, $K</em>2S$, and $(NH<em>4)</em>2S$.

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 ($Pb(NO<em>3)</em>2$) 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:

  • $AgNO<em>3(aq) + NaCl(aq) ightharpoonup AgCl(s) + NaNO</em>3(aq)$

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: $Fe(NO<em>3)</em>3(aq) + Na<em>2S(aq) ightharpoonup Fe</em>2S_3(s)$

• Add Other Compound: $Fe(NO<em>3)</em>3(aq) + Na<em>2S(aq) ightharpoonup Fe</em>2S<em>3(s) + NaNO</em>3(aq)$

• Balanced Equation: $2Fe(NO<em>3)</em>3(aq) + 3Na<em>2S(aq) ightharpoonup Fe</em>2S<em>3(s) + 6NaNO</em>3(aq)$

• 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: $AgNO<em>3(aq) + NaCl(aq) ightharpoonup AgCl(s) + NaNO</em>3(aq)$

• Ionic equation: $Ag^+(aq) + Cl^-(aq) <br>ightharpoonup AgCl(s)$

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: $<br>ightharpoonup Al<em>2S</em>3$

• Add Ions: $2Al^{3+} + 3S^{2-} <br>ightharpoonup Al<em>2S</em>3$

• 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.