Chemistry review

Acids, Bases, and Salts

Preparation of Acids and Bases

  • Acids: Release hydrogen ions (H^+) when dissolved in water.

    • Example: Hydrochloric acid (HCl) dissociates into H^+ and Cl^- ions.

  • Bases: Release hydroxide ions (OH^-) when dissolved in water.

    • Example: Sodium hydroxide (NaOH) dissociates into Na^+ and OH^- ions.

  • pH: Represents the "phy-power of hydrogen," indicating the acidity or basicity of a solution.

Salt Formation

  • Salt Definition: A salt is an ionic compound formed by the reaction between an acid and a base.

  • General Reaction: Acid + Base \rightarrow Salt + Water

    • Example: HCl + NaOH \rightarrow NaCl + H_2O

  • Ionic Composition: A salt consists of the positive ion (cation) from the base and the negative ion (anion) from the acid.

Solubility Rules for Salts

  • Soluble Salts:

    • All sodium, potassium, and ammonium salts are soluble.

    • All nitrates are soluble.

    • Chlorides are generally soluble, except for silver chloride (AgCl) and lead(II) chloride (PbCl_2).

    • Sulfates are generally soluble, except for calcium sulfate (CaSO4), barium sulfate (BaSO4), and lead(II) sulfate (PbSO_4).

  • Insoluble Salts:

    • All other carbonates are insoluble, except for sodium, potassium, and ammonium carbonates.

    • All other hydroxides are insoluble, except for sodium, potassium, and ammonium hydroxides.

Neutralization Reaction

  • Definition: The reaction between an acid and a base to form a salt and water.

  • General Equation: Acid + Base \rightarrow Salt + Water

    • Example: HCl + NaOH \rightarrow NaCl + H_2O

  • Alkalis: Bases that contain hydroxide ions and metals.

  • Metallic Oxides and Hydroxides: All metallic oxides and hydroxides are bases.

  • Neutral Substances: Substances that neither contain hydrogen ions (H^+) nor hydroxide ions (OH^-).

Preparation of Insoluble Salts

  • General Method: Insoluble salts are prepared by mixing solutions of soluble salts to form a precipitate.

  • Reaction Type:

    • Soluble Salt + Soluble Salt → Insoluble Salt (Precipitate) + Soluble Salt

  • Example: Preparation of Silver Chloride (AgCl)

    • Reactants: Silver nitrate (AgNO_3) and sodium chloride (NaCl)

    • Equation: AgNO3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO3(aq)

    • Procedure:

      1. Dissolve silver nitrate in water (Beaker 1).

      2. Dissolve sodium chloride in water (Beaker 2).

      3. Mix the two solutions.

      4. Silver ions (Ag^+) and chloride ions (Cl^-) combine to form insoluble silver chloride (AgCl), which precipitates out as a white solid.

      5. Sodium nitrate (NaNO_3) remains dissolved in the solution.

      6. Filter the mixture to separate the silver chloride precipitate from the sodium nitrate solution.

  • Polyatomic Ions:

    • Cations: NH_4^+ (Ammonium)

    • Anions: PO4^{3-} (Phosphate), NO3^- (Nitrate), SO_4^{2-} (Sulfate)

Steps for Preparing Insoluble Salts

  1. Mixing Reactants: Combine solutions containing the desired ions.

  2. Precipitation: The insoluble salt forms as a precipitate.

  3. Filtration: Separate the precipitate from the solution using filter paper and a funnel.

  4. Washing: Wash the precipitate to remove any remaining soluble ions.

  5. Drying: Dry the precipitate.

Examples of Insoluble Salt Preparation

  1. Barium sulfate (BaSO_4):

    • Ba(NO3)2(aq) + Na2SO4(aq) \rightarrow BaSO4(s) + 2NaNO3(aq)

  2. Calcium sulfate (CaSO_4):

    • Ca(NO3)2(aq) + Na2SO4(aq) \rightarrow CaSO4(s) + 2NaNO3(aq)

  3. Magnesium carbonate (MgCO_3):

    • Mg(NO3)2(aq) + Na2CO3(aq) \rightarrow MgCO3(s) + 2NaNO3(aq)

  4. Calcium hydroxide (Ca(OH)_2):

    • Ca(NO3)2(aq) + 2NaOH(aq) \rightarrow Ca(OH)2(s) + 2NaNO3(aq)

  5. Lead(II) sulfate (PbSO_4):

    • Pb(NO3)2(aq) + Na2SO4(aq) \rightarrow PbSO4(s) + 2NaNO3(aq)

Preparation of Soluble Salts from Insoluble Salts

  1. Reactants: Insoluble salt (e.g., zinc) and dilute acid (e.g., sulfuric acid).

  2. Procedure:

    1. Add excess insoluble salt to the dilute acid until the reaction stops.

    2. Filter off the excess solid.

    3. Evaporate most of the water using a hot water bath.

    4. Leave the remaining solution to cool and crystallize.

    5. Filter the crystals and dry them on filter paper.

  3. Example:

    • Zn + H2SO4(aq) \rightarrow ZnSO4(aq) + H2(g)

Titration

  • Definition: A method of analyzing the concentration of a solution.

  • Applications:

    • Determining solution concentration.

    • Preparing soluble salts.

  • Acid-Base Titration: A specific type of titration involving the reaction of an acid with a base.

Apparatus for Titration

  • Conical flask

  • Burette

  • Beaker

Preparation of Soluble Salt (Sodium Chloride) via Titration

  1. Fill a burette with dilute hydrochloric acid (HCl) to the zero mark.

  2. Put 25 cm³ of dilute sodium hydroxide (NaOH) into a conical flask.

  3. Add a few drops of a suitable indicator, such as universal indicator.

  4. Slowly add acid from the burette into the flask containing the alkali until the indicator changes color (e.g., from pink to colorless).

  5. Repeat the process without the indicator, using the recorded volume of acid and the same volume of alkali.

  6. Evaporate most of the water from the solution using a Bunsen burner and evaporating dish.

  7. After evaporation, let the solution cool to form crystals. Filter the crystals and dry them.

Preparation of Soluble Salts from Insoluble Solids

  • Insoluble Solid Types: Metal, base, or carbonate.

  • General Reaction: Insoluble Solid + Acid \rightarrow Salt + Other Products

    • Example: Zn + H2SO4 \rightarrow ZnSO4 + H2

Displacement Reactions

  • Carbon can displace zinc, iron, tin, and lead from their ores.

  • Historical Context: Approximately 3500 years ago, it was discovered that heating iron ore with charcoal at high temperatures produces molten iron.

    • Iron Ore + Charcoal \rightarrow Iron + Carbon Dioxide

  • Charcoal: A form of coal.

Carbonic Acid and Carbonates

  • Carbonic Acid: A weak acid formed when carbon dioxide (CO_2) reacts with water.

  • Carbonates: Salts derived from carbonic acid.

Citric Acid and Citrates

  • Citric Acid: Found in fruits such as oranges and lemons.

  • Citrates: Salts formed using citric acid.

Reaction of Metals with Dilute Acids

  • General Reaction: Metal + Acid \rightarrow Salt + Hydrogen

    • Example: Zn + 2HCl \rightarrow ZnCl2 + H2

Unreactive Metals

  • Some metals, like silver and copper, are too unreactive to displace hydrogen from acids.

Reaction of Metal Oxides with Acids

  • General Reaction: Metal Oxide + Acid \rightarrow Salt + Water

    • Example: CuO + H2SO4 \rightarrow CuSO4 + H2O

Metal Carbonates and Acids

  • Carbonates: Such as calcium carbonate, are salts that can be formed by the reaction of a metal with carbonic acid.

  • Examples:

    • The skeletons of coral are made of calcium carbonate and react with acids.

    • Blue-green colors in rocks (e.g., in the Atacama Desert) indicate the presence of copper salts, such as malachite (copper carbonate).

Forming Salts by Neutralization

  • Neutralization Reaction: Alkalis react with acids to neutralize them, forming a salt and water.

    • Example: NaOH + HCl \rightarrow NaCl + H_2O

  • Alkalis and Bases: Metal oxides that dissolve in water form alkaline solutions and are considered bases.

  • Soluble Metal Bases: Form alkalis when dissolved in water.

    • Example: Sodium Oxide + Water \rightarrow Sodium Hydroxide

  • Insoluble Metal Oxides: React with acids to form salts.

    • Example: CuO + H2SO4 \rightarrow CuSO4 + H2O

Ways to Make Soluble Salts

  1. Metal + Acid → Salt + Hydrogen Gas

  2. Metal Oxide + Acid → Salt + Water

  3. Metal Carbonate + Acid → Salt + Water + Carbon Dioxide

  4. Acid + Alkali → Salt + Water

Preparation of Soluble Salts by Neutralization

  1. Acid with Excess Insoluble Oxide:

    1. Pour dilute acid into a beaker.

    2. Add small amounts of insoluble metal oxide while stirring gently (reactions are faster at higher temperatures).

    3. Continue adding metal oxide until no more dissolves (ensuring all acid is used).

    4. Filter excess solid.

    5. Transfer the solution to an evaporating dish.

    6. Warm the mixture to evaporate some water (too much heat can drive off water of crystallization).

    7. Let the solution cool to allow crystallization.

    8. Filter off the crystals, wash with cold water, and dry with filter paper.

    9. Allow the crystals to dry naturally in air (heat can cause decomposition).

    • Example: CaO + HNO3 \rightarrow Ca(NO3)2 + H2O

  2. Acid with Excess Insoluble Carbonate:

    • Same procedure as with excess insoluble oxide, except:

      1. Do not heat, as the reaction is fast.

      2. Effervescence (CO_2 release) is observed.

    • Example: MgCO3 + H2SO4 \rightarrow MgSO4 + CO2 + H2O

  3. Acid with Excess Metal:

    • Not all metals are suitable (some are too reactive or do not react).

      1. Place dilute acid in a beaker.

      2. Carefully add small amounts of metal while stirring (avoid excessive effervescence).

      3. Continue adding metal until no more dissolves.

      4. Filter the solution to remove excess metal.

      5. Transfer the mixture to a tripod and warm it to remove some water.

      6. Let the solution cool to form crystals.

      7. Filter off crystals, wash with cold water, and dry with filter paper.

      8. Allow crystals to dry naturally in air.

Summary of Salt Preparation Methods

  • Insoluble Salts: Soluble Salt + Soluble Salt → Insoluble Salt (Precipitate)

  • Soluble Salts:

    • Metal + Acid → Salt + Hydrogen

    • Metal Oxide + Acid → Salt + Water

    • Metal Carbonate + Acid → Salt + Water + CO_2

    • Alkali + Acid → Salt + Water

Law of Conservation of Mass

  • In a chemical reaction, atoms are rearranged into new combinations.

  • Law of Conservation of Mass: The total mass at the end of the reaction is the same as at the start.