Grade 10 chemisty Physical Chemistry – Introduction to Salts Notes
Physical Chemistry — Introduction to Salts
- Classify different salts based on their properties (soluble and insoluble).
- Prepare salts using appropriate laboratory methods and write balanced chemical equations for the reactions involved. (Ionic equations limited to precipitation reactions.)
- Describe the behaviour of salts when exposed to air.
- Outline and appreciate applications of salts in day-to-day life (agriculture, food industry, medicine, paper, paints, glass, laundry, etc.).
- Evaluate environmental effects of applications of salts such as inorganic fertilisers (water pollution, eutrophication, soil and air pollution).
- Discuss measures to address challenges of using inorganic fertilisers in agriculture.
What is a salt?
A salt is an ionic compound made of positive ions (cations) and negative ions (anions). Many salts form when an acid reacts with a base (neutralisation) or when two soluble salts react and an insoluble salt precipitates. Common example: sodium chloride (table salt), NaCl.
Classification by solubility
Salts are often grouped as soluble or insoluble in water. Solubility matters for uses and how salts are prepared in the lab.
- Sodium, potassium and ammonium salts (Na+, K+, NH4+) — e.g., NaCl, KNO3.
- All nitrates (NO3–) are soluble — e.g., KNO3 (used in fertilizers and fireworks).
- Most chlorides are soluble (exceptions: AgCl, PbCl2).
- Most sulfates are soluble (exceptions: BaSO4, PbSO4).
- Silver chloride, AgCl (white precipitate).
- Lead(II) iodide, PbI2 (yellow precipitate).
- Barium sulfate, BaSO4 (used as a radiocontrast agent and insoluble).
- Calcium carbonate, CaCO3 (limestone, chalk — insoluble).
Note: Solubility rules are handy for predicting precipitation when two solutions are mixed.
How to prepare salts (lab methods)
- Method: React exactly the right amounts of an acid and a base (titration), then evaporate the water to crystallise the salt.
Example (make sodium chloride solution): HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l). After neutralisation, evaporate the water to get solid NaCl.
- Titrate HCl with NaOH to the neutral point (using indicator).
- Evaporate the neutral solution gently to crystallise the salt.
- Collect and dry crystals.
- Method: Mix two solutions containing ions that form an insoluble salt. The insoluble salt appears as a precipitate and is filtered off, washed and dried.
Molecular equation: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Ionic equation (precipitation only): Ag+(aq) + Cl–(aq) → AgCl(s)
Molecular equation: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)
Ionic equation: Pb2+(aq) + 2I–(aq) → PbI2(s)
Used to obtain pure crystals from a solution (e.g., obtain NaCl from seawater/brine by evaporating water).
- Wear goggles, gloves and a lab coat.
- Handle concentrated acids and bases carefully — neutralise spills and wash skin immediately.
- Dispose of heavy metal salts (e.g., lead, silver) as hazardous waste.
Behaviour of salts when exposed to air
Salts can change when left in air. Common behaviours:
- Hygroscopic — absorb water from air (but do not necessarily dissolve). Example: CaCl2 is hygroscopic.
- Deliquescent — absorb enough water to dissolve and form a liquid solution. Example: Calcium chloride (CaCl2) and some salts used as drying agents.
- Efflorescent — hydrated salts that lose water to air and become powdery. Example: CuSO4·5H2O gradually losing water to become anhydrous CuSO4.
- Reaction with gases — some salts react with gases in the air. Example: NaOH/alkali may absorb CO2 → forms carbonate on the surface (Na2CO3).
Applications of salts in day-to-day life (Kenyan context)
- Nutrient salts: CAN (calcium ammonium nitrate), DAP (diammonium phosphate), NPK blends, ammonium sulfate.
- Used widely by Kenyan farmers to increase crop yields (maize, tea, horticulture).
- Table salt (NaCl) for cooking and preservation.
- Baking soda (NaHCO3) — in baking.
- Saline (NaCl) for rehydration and wound cleaning.
- MgSO4 (Epsom salt) used as a laxative and bath additive.
- Glass: sodium carbonate (soda ash) and calcium carbonate.
- Papers and paints: various salts used in processing and pigments.
- Laundry: washing soda (Na2CO3) and other salts aid cleaning.
Environmental effects of inorganic fertilisers (examples & evaluation)
Use of inorganic fertilisers (salts) improves yields but can cause environmental problems if misused:
- Water pollution / eutrophication: Runoff carries nitrates and phosphates into rivers and lakes. Excess nutrients cause algal blooms (e.g., risks for Lake Victoria and local water bodies).
- Soil degradation: Overuse can cause salinisation, loss of soil structure, or acidification, reducing long-term fertility.
- Air pollution and greenhouse gases: Nitrous oxide (N2O) from nitrogen fertilisers contributes to global warming; ammonia volatilisation can cause local air pollution.
- Human health risks: High nitrate levels in drinking water can affect infants (blue baby syndrome) and livestock.
Measures to address challenges of inorganic fertilisers
Practical strategies farmers, communities and policymakers can use:
- Soil testing before application — apply the right amount and type of fertiliser.
- Integrated soil fertility management — combine organic manures (compost, farmyard manure) with inorganic fertilisers.
- Precision application — place fertiliser near plant roots, avoid broadcasting; use appropriate timing (avoid rainy season at application time).
- Controlled-release and nano-fertilisers to reduce leaching and emissions.
- Buffer strips and vegetative barriers along rivers to reduce runoff.
- Use of cover crops and crop rotations to reduce nutrient loss and improve soil health.
- Farmer education, extension services and regulations to ensure correct use and handling.
Suggested learning experiences (classroom & field)
- Carry out a precipitation experiment: mix AgNO3(aq) and NaCl(aq) to make AgCl(s). Write molecular and ionic equations.
- Prepare a soluble salt by titration (acid + base) and crystallise it. Record mass of crystals obtained.
- Investigate behaviour of hydrated salts: observe CuSO4·5H2O losing water on heating; observe a deliquescent salt like CaCl2 absorbing moisture.
- Field visit: visit a local farm to see fertiliser use. Interview farmer about types (DAP, CAN, NPK) and methods used.
- Class debate: “Are inorganic fertilisers more helpful than harmful?” — students research environmental impacts and propose solutions.
Salts are ionic compounds important in daily life and industry. Learn to classify salts by solubility, prepare them using neutralisation, precipitation and crystallisation, and consider both benefits and environmental risks — especially in agriculture. Use good laboratory practice and sustainable farming methods to reduce harm.
Short class quiz / tasks
- Name a soluble salt and an insoluble salt and give one use for each.
- Write the ionic equation for the precipitation of AgCl.
- List two ways farmers can reduce fertiliser runoff into rivers.
- Plan a simple experiment to make and dry crystals of a salt in the school laboratory (include safety steps).