Comprehensive Study Notes on Chemistry of Salts: Structure, Properties, Solubility, and Preparation Methods

Chemical Structure and Lattice Organization of Salts

• Definition of Neutralization Reaction:

  • A neutralization reaction occurs when an acid reacts with a base or alkali to produce a salt and water.

  • Example neutralization equation between sodium hydroxide ($\text{NaOH}$) and hydrochloric acid ($\text{HCl}$):   $\text{NaOH}_{(aq)} + \text{HCl}_{(aq)} \rightarrow \text{NaCl}_{(aq)} + \text{H}_2\text{O}_{(l)}$

• Chemical Nature and Bonding of Salts:

  • A salt is a chemical combination of positive ions (cations) and negative ions (anions).

  • Oppositely charged ions are bonded together by electrostatic forces of attraction.

  • The strength of this attraction depends on two key parameters:

  • The magnitude of the ionic charges.

  • The distance between the centers of the ions.

• Crystal Lattice Structure:

  • Under ordinary conditions of temperature and pressure, ionic compounds exist as crystalline solids.

  • Ions arrange themselves in a regular, repeating pattern in three dimensions to form a network structure called a crystal lattice.

  • The characteristic properties shown by ionic compounds are directly caused by the presence of this ordered three-dimensional structure.

• Specific Geometry of Sodium Chloride ($\text{NaCl}$):

  • In a sodium chloride crystal lattice, each sodium ion ($\text{Na}^+$) is surrounded by six chloride ions ($\text{Cl}^-$).

  • Reciprocally, each chloride ion ($\text{Cl}^-$) is surrounded by six sodium ions ($\text{Na}^+$).

  • The arrangement forms a face-centered cubic structure.

  • The larger chloride ions are located at the corners and at the center of each face of the cube.

  • The smaller sodium ions occupy the edges of the cube.

• Historical Value of Common Salt:

  • Common salt was a highly valuable commodity in the past and was historically utilized as currency.

Physical Properties and Conductivity of Ionic Compounds

• Thermal Properties and Melting Points:

  • Ionic compounds generally possess very high melting points because substantial thermal energy is required to break the strong electrostatic forces of attraction holding the crystal lattice together.

  • During the melting process, these interionic attractive forces break down, permitting the ions to move freely in the liquid state.

  • Factors affecting melting point:

  • Higher charges on ions produce stronger electrostatic attractions, leading to higher melting points.

  • Smaller ionic sizes bring charges closer together, strengthening ionic bonds and increasing melting points.

• Electrical Conductivity:

  • Solid State: Ionic compounds do not conduct electricity in the solid state because their ions are held tightly in fixed lattice positions by strong forces of attraction and cannot move.

  • Molten State: Heating an ionic compound to its melting point breaks down the lattice, freeing the ions to move around and conduct electricity.

  • Dissolved State: Dissolving ionic compounds in water breaks down the attractive forces between ions, leaving them free to move independently.

  • Electrolytic Nature: Ionic compounds function as strong electrolytes because they dissociate completely into ions when dissolved in water.

Solubility Rules of Salts in Water

• Key Factors Influencing Solubility:

  • The chemical nature of the salt.

  • The temperature of the solution.

• Rules for Soluble Salts:

  • All salts containing sodium ($\text{Na}^+$), potassium ($\text{K}^+$), and ammonium ($\text{NH}_4^+$) ions are soluble in water.

  • All metallic nitrates ($\text{NO}_3^-$) are soluble in water.

  • Halides: Most chlorides ($\text{Cl}^-$), bromides ($\text{Br}^-$), and iodides ($\text{I}^-$) are soluble in water, except those of silver ($\text{Ag}^+$) and lead(II) ($\text{Pb}^{2+}$).

  • Sulfates: Most sulfates ($\text{SO}_4^{2-}$) are soluble in water, except those of barium ($\text{Ba}^{2+}$), lead(II) ($\text{Pb}^{2+}$), and calcium ($\text{Ca}^{2+}$).

• Rules for Insoluble and Partially Soluble Salts:

  • Carbonates: All carbonates ($\text{CO}_3^{2-}$) are insoluble in water, with the exception of sodium, potassium, and ammonium carbonates.

  • Hydroxides: All hydroxides ($\text{OH}^-$) are insoluble in water, except sodium, potassium, and ammonium hydroxides. Calcium hydroxide ($\text{Ca(OH)}_2$) is partially soluble in water.

• Comprehensive Solubility Summary:

  • Soluble Salt Categories:

  • All salts of sodium, potassium, and ammonium.

  • All nitrate salts.

  • Most chloride, bromide, and iodide salts.

  • Most sulfate salts.

  • Insoluble Salt Categories:

  • Silver chloride/bromide/iodide and lead(II) chloride/bromide/iodide.

  • Barium sulfate, lead(II) sulfate, and calcium sulfate.

  • Most carbonates (except sodium, potassium, and ammonium carbonates).

  • Most hydroxides (except sodium, potassium, ammonium, and partially soluble calcium hydroxide).

Methods for Preparing Soluble Salts

• Method 1: Reaction of Water-Soluble Acid with Water-Soluble Base (Titration Method):

  • Applicability: Used when both the starting acid and base are soluble in water.

  • Principles: Selection of acid and base depends directly on the specific salt desired.

  • Preparation Example (Potassium Nitrate, $\text{KNO}_3$):

  • React appropriate volumes of potassium hydroxide ($\text{KOH}$) and nitric acid ($\text{HNO}_3$) in a conical flask:     $\text{KOH}_{(aq)} + \text{HNO}_{3(aq)} \rightarrow \text{KNO}_{3(aq)} + \text{H}_2\text{O}_{(l)}$

  • General Crystallization Procedure:

  • Transfer the neutralized solution from the conical flask into an evaporating dish.

  • If the product salt is stable towards heat, evaporate the solution to dryness.

  • If heat-sensitive, heat the solution gently until a thin film of crystals forms on the liquid surface.

  • Cool the saturated solution slowly to encourage crystal growth.

  • Filter the mixture to isolate pure crystals, then dry them between folds of filter paper.

• Method 2: Reaction of an Acid with an Insoluble Base or Metal:

  • Applicability: Neutralization reaction between a soluble acid and an insoluble base (metal oxide, metal hydroxide, or metal carbonate) or a reactive metal.

  • General Operational Steps:

  • Mix the selected acid with an excess of the chosen insoluble base or metal while stirring constantly to ensure all acid is fully consumed.

  • Filter the resulting mixture to remove the excess undissolved base or metal.

  • Transfer the filtrate to an evaporating dish and heat gently to form a concentrated, saturated solution.

  • Allow the solution to cool down slowly to yield pure crystals of the salt.

  • Filter the crystalline mixture and dry the crystals carefully.

  • Specific Chemical Reaction Equations:

  • Acid with insoluble hydroxide:     $\text{Cu(OH)}_{2(s)} + 2\text{HCl}_{(aq)} \rightarrow \text{CuCl}_{2(aq)} + 2\text{H}_2\text{O}_{(l)}$

  • Acid with metal:     $\text{Mg}_{(s)} + \text{H}_2\text{SO}_{4(aq)} \rightarrow \text{MgSO}_{4(aq)} + \text{H}_{2(g)}$

  • Acid with insoluble carbonate:     $\text{CaCO}_{3(s)} + 2\text{HCl}_{(aq)} \rightarrow \text{CaCl}_{2(aq)} + \text{CO}_{2(g)} + \text{H}_2\text{O}_{(l)}$

Practical Laboratory Activities

• Activity 18.1: Preparation of Pure Crystals of Sodium Sulphate ($\text{Na}_2\text{SO}_4$) via Titration:

  • Objective: Prepare pure crystals of soluble sodium sulphate.

  • Method Steps:

  • Use a pipette to add $50\,cm^3$ of $1\,mol\,dm^{-3}$ sodium hydroxide ($\text{NaOH}$) solution into a clean conical flask.

  • Add a few drops of phenolphthalein indicator to obtain a pink solution.

  • Run $1\,mol\,dm^{-3}$ sulphuric acid ($\text{H}_2\text{SO}_4$) from a burette into the conical flask until the color changes from light pink to colorless.

  • Record the exact volume of sulphuric acid required for complete neutralization.

  • Chemical Equation:     $2\text{NaOH}_{(aq)} + \text{H}_2\text{SO}_{4(aq)} \rightarrow \text{Na}_2\text{SO}_{4(aq)} + 2\text{H}_2\text{O}_{(l)}$

  • Repeat the procedure using the exact recorded volumes of sodium hydroxide and sulphuric acid without adding any indicator.

  • Transfer the neutralized solution to an evaporating dish and heat gently until one-third of the initial solution volume remains.

  • Dip a glass rod into the hot solution and withdraw it; if the immersed tip turns cloudy, the solution is sufficiently concentrated for crystallization.

  • Cool the solution to allow crystallization to complete.

  • Filter and wash the crystals with a small amount of cold distilled water to remove residual acid and soluble impurities.

  • Dry the crystals between folds of filter paper or in an oven.

• Activity 18.2: Preparation of Pure Crystals of Zinc Sulphate ($\text{ZnSO}_4$) using Metal Granules:

  • Objective: Prepare pure crystals of zinc sulphate.

  • Method Steps:

  • Place approximately $25\,cm^3$ of dilute sulphuric acid ($\text{H}_2\text{SO}_4$) into a beaker and warm it gently.

  • Slowly add small granules of zinc metal while stirring constantly until undissolved metal begins settling at the bottom of the beaker.

  • Chemical Equation:     $\text{Zn}_{(s)} + \text{H}_2\text{SO}_{4(aq)} \rightarrow \text{ZnSO}_{4(aq)} + \text{H}_{2(g)}$

  • Filter the solution to remove unreacted zinc metal and collect the filtrate.

  • Evaporate the filtrate gently in an evaporating dish to concentrate the solution without overheating or boiling.

  • Cool the solution slowly to allow crystals to form.

  • Once crystal formation is complete, filter the mixture carefully to collect the pure crystals.

  • Dry the crystals on filter paper.   

    

Uses and Applications of Salts

• Agricultural and Biological Applications:

  • Copper sulphate ($\text{CuSO}_4$) is used as a fungicide.

• Industrial Applications:

  • Various salts are utilized in water softening procedures.

  • Salts are essential raw materials in glass manufacture.

Questions & Discussion

• Section 18.1 Quick Check:

  • Question 1: How do the ions present in $\text{NaCl}$ arrange themselves to form a crystal lattice?

  • Response: Sodium ($\text{Na}^+$) and chloride ($\text{Cl}^-$) ions organize into a three-dimensional face-centered cubic crystal lattice. Each $\text{Na}^+$ ion is surrounded by six $\text{Cl}^-$ ions, and each $\text{Cl}^-$ ion is surrounded by six $\text{Na}^+$ ions. The larger chloride ions sit at the corners and face centers, while the smaller sodium ions occupy the edge positions.

  • Question 2: Do you expect the melting point of $\text{KCl}$ to be higher or lower than $\text{NaCl}$?

  • Response: Lower. The potassium ion ($\text{K}^+$) is larger in size than the sodium ion ($\text{Na}^+$). The larger ionic radius increases the distance between oppositely charged ions in the lattice, weakening the electrostatic attraction forces and resulting in a lower melting point.

• Section 18.2 Quick Check:

  • Question 1: Which salts of barium and calcium are soluble in water?

  • Response: Soluble barium salts include barium nitrate ($\text{Ba(NO}_3\text{)}_2$) and barium chloride ($\text{BaCl}_2$). Soluble calcium salts include calcium nitrate ($\text{Ca(NO}_3\text{)}_2$) and calcium chloride ($\text{CaCl}_2$), with calcium hydroxide ($\text{Ca(OH)}_2$) being partially soluble.

  • Question 2: How is barium sulphate prepared in the laboratory?

  • Response: Barium sulphate ($\text{BaSO}_4$) is an insoluble salt and is prepared by a precipitation method involving the reaction of two soluble salts (such as aqueous barium chloride and aqueous sodium sulphate). The precipitated barium sulphate is filtered, washed with distilled water, and dried.