i. Which of the following salts is soluble in water?
Answer: c) Sodium nitrate (NaNO₃)
Explanation: All sodium (Na⁺), potassium (K⁺), and ammonium (NH₄⁺) salts are soluble in water. Silver chloride, lead carbonate, and lead chloride are insoluble or sparingly soluble.
ii. What type of ions do bases provide for salt formation?
Answer: a) Positive ions (Note: Bases provide hydroxide ions (OH⁻) or accept protons. In the context of salt formation, bases provide the positive ion (cation) for the salt.)
Explanation: In a neutralization reaction, the base provides the positive ion (cation) that combines with the negative ion (anion) from the acid to form the salt.
iii. What is the primary reason salts have high melting points?
Answer: b) Strong electrostatic forces
Explanation: Salts are ionic compounds consisting of a giant lattice structure where oppositely charged ions are held together by strong electrostatic forces of attraction. Breaking these bonds requires a significant amount of energy, resulting in high melting points.
iv. In what state are salts generally good conductors of electricity?
Answer: b) Molten (Note: Aqueous solutions are also good conductors.)
Explanation: In the solid state, ions are locked in fixed positions and cannot move. When molten (or dissolved in water), the ions are free to move and carry an electric charge, making them good conductors of electricity.
v. Which of the following is a general solubility rule for chlorides?
Answer: c) Chlorides are soluble except lead and silver chlorides.
Explanation: Most chloride salts are soluble in water. Notable exceptions include silver chloride (AgCl) and lead(II) chloride (PbCl₂), which are insoluble or sparingly soluble.
vi. What happens to the ions in an ionic compound when it dissolves in water?
Answer: b) They become mobile.
Explanation: When an ionic compound dissolves in water, the water molecules surround the ions and separate them from the lattice. These hydrated ions are free to move throughout the solution.
vii. Which of the following methods is used to prepare a salt by titration?
Answer: d. Acid + Alkali
Explanation: Titration is used when reacting an acid with an alkali (soluble base) because both are in solution and the exact endpoint of neutralization can be determined using an indicator.
viii. Which of these salts is insoluble in water?
Answer: d. Lead chloride (PbCl₂)
Explanation: Potassium, ammonium, and calcium salts are generally soluble. Lead chloride is insoluble in cold water.
ix. What is formed when an acid reacts with an excess of an insoluble base?
Answer: c. Salt and carbon dioxide (Note: This reaction specifically refers to an acid reacting with an insoluble carbonate, which produces salt, water, and carbon dioxide. An insoluble base like CuO would produce salt and water only. Given the options, "c" is the most likely intended answer if "insoluble base" was meant to be "insoluble carbonate.")
x. Which ion is commonly found in soluble nitrates?
Answer: b. NO₃⁻
Explanation: The nitrate ion (NO₃⁻) is the defining anion for nitrates. All common nitrates are soluble.
i. Define what a salt is in terms of ions.
Answer: A salt is an ionic compound composed of a positively charged cation (often a metal ion or ammonium ion) and a negatively charged anion (derived from an acid). It is typically formed when the hydrogen ion of an acid is replaced by a metal ion or an ammonium ion.
ii. Why do salts generally have high melting points?
Answer: Salts have high melting points because they exist as giant ionic lattices. The oppositely charged ions are held together by strong electrostatic forces of attraction. To melt a salt, these strong forces must be overcome, which requires a large amount of thermal energy.
iii. Describe the structure of an ionic compound.
Answer: An ionic compound consists of a three-dimensional giant lattice structure. In this lattice, positive and negative ions are arranged in a regular, repeating pattern. The ions are held in fixed positions by strong electrostatic forces of attraction that act in all directions.
iv. How do salts conduct electricity when molten or in solution?
Answer: In the solid state, ions are fixed in the lattice. When molten or dissolved in water, the lattice breaks down, and the ions become free to move. These mobile ions act as charge carriers, allowing the salt to conduct electricity.
v. List three salts that are generally soluble in water.
Answer: Three examples of soluble salts are:
1. Sodium chloride (NaCl)
2. Potassium nitrate (KNO₃)
3. Ammonium sulphate ((NH₄)₂SO₄)
vi. What happens to the ions in an ionic compound when it dissolves in water?
Answer: When an ionic compound dissolves in water, the polar water molecules surround the individual ions. The positive ends of water molecules are attracted to anions, and the negative ends are attracted to cations. This process, called hydration, separates the ions from the lattice, allowing them to move freely in the solution.
vii. Do you think that most carbonates are soluble in water?
Answer: No, most carbonates are insoluble in water. Exceptions include sodium carbonate (Na₂CO₃), potassium carbonate (K₂CO₃), and ammonium carbonate ((NH₄)₂CO₃).
viii. How can you prepare a soluble salt using an excess of an insoluble base?
Answer: To prepare a soluble salt using an excess of an insoluble base (e.g., copper(II) oxide), add the insoluble base to a warm acid (e.g., sulfuric acid) until no more dissolves. Filter the mixture to remove the excess unreacted base. The filtrate contains the soluble salt. Evaporate the water to crystallize the salt.
ix. Which salts are typically soluble according to the general solubility rules?
Answer: According to general rules, all salts of sodium, potassium, and ammonium are soluble. All nitrates are soluble. Most sulphates are soluble except barium sulphate and lead sulphate. Most chlorides are soluble except silver chloride and lead chloride.
1. Explain the process of preparing a soluble salt using titration. Include a detailed description of the steps involved.
Answer: Titration is used to prepare soluble salts when reacting an acid with an alkali (a soluble base). This method is preferred because both reactants are in solution and the exact point of neutralization can be determined. The process involves the following detailed steps:
Step 1: Preparation of Solutions: A standard solution of the alkali (e.g., sodium hydroxide) is prepared with a known concentration. The acid (e.g., hydrochloric acid) is also prepared.
Step 2: Performing a Trial Titration:
A pipette is used to transfer a precisely known volume of the alkali (e.g., 25.0 cm³) into a conical flask. A few drops of a suitable indicator (e.g., phenolphthalein) are added. The solution turns pink. A burette is filled with the acid. The acid is slowly added to the flask while swirling until the pink color disappears completely. This is the endpoint. The volume of acid used is recorded.
Step 3: Performing an Accurate Titration: The trial titration is repeated several times to obtain concordant results (volumes within 0.1 cm³ of each other). The average volume of acid required to neutralize the alkali is calculated.
Step 4: Preparing the Salt Solution:
A new volume of the alkali is measured into a clean conical flask. The calculated volume of acid (from the accurate titration) is added to the flask. This time, no indicator is used because the indicator would contaminate the salt. The reaction is now complete, and the solution contains only the salt and water.
Step 5: Crystallization:
The resulting solution is transferred to an evaporating dish. The solution is gently heated to evaporate some of the water, allowing the solution to become saturated. The dish is left to cool slowly. Crystals of the salt will form.
Step 6: Filtration and Drying:
The crystals are filtered off from the remaining liquid (mother liquor) using a Buchner funnel or filter paper. The crystals are pressed between filter papers or left to dry in a warm place.
In summary, titration allows for the precise preparation of a soluble salt by accurately neutralizing an acid with an alkali, avoiding contamination from excess reactants.
2. Discuss the properties of salts that make them solid at STP and have high melting points.
Answer: Salts are solid at Standard Temperature and Pressure (STP) and possess high melting points due to their unique structure and bonding.
1. Giant Ionic Lattice Structure: Salts are composed of a vast, three-dimensional network of alternating positive and negative ions. This is known as a giant ionic lattice. In this structure, each ion is surrounded by ions of opposite charge. For example, in sodium chloride (NaCl), each Na⁺ ion is
surrounded by six Cl⁻ ions and vice versa.
2. Strong Electrostatic Forces of Attraction: The oppositely charged ions in the lattice are held together by powerful electrostatic forces of attraction. These forces act in all directions and are very strong. To melt the salt, a large amount of thermal energy is required to overcome these forces and allow the ions to break free from their fixed positions and move around. This results in high melting points.
3. Solid State at STP: Because the electrostatic forces are so strong, the ions vibrate around fixed positions but do not have enough energy at room temperature (STP) to overcome these forces. Therefore, salts exist as rigid, solid crystals at STP.
4. Relationship to Melting Point: The strength of the electrostatic forces depends on the charge and size of the ions. Salts with ions of higher charge (e.g., MgO, 2+ and 2-) have stronger forces and higher melting points than salts with ions of lower charge (e.g., NaCl, 1+ and 1-).
In summary, the combination of a giant ionic lattice and strong electrostatic forces of attraction explains why salts are solid with high melting points at STP.
3. Discuss the solubility rules for salts and explain the exceptions for chlorides, carbonates, and hydroxides.
Answer: Solubility rules are guidelines used to predict whether an ionic compound (salt) will dissolve in water.
1. General Solubility Rules:
All salts of Group 1 metals (Li⁺, Na⁺, K⁺, etc.) are soluble. All salts of ammonium (NH₄⁺) are soluble. All nitrates (NO₃⁻) are soluble.
All acetates (CH₃COO⁻) are soluble. Most sulphates (SO₄²⁻) are soluble.
2. Exceptions for Chlorides (Cl⁻):
Most chlorides are soluble. However, silver chloride (AgCl) and lead(II) chloride (PbCl₂) are insoluble or sparingly soluble. Mercury(I) chloride (Hg₂Cl₂) is also insoluble.
3. Exceptions for Carbonates (CO₃²⁻):
Most carbonates are insoluble. The major exceptions are the carbonates of Group 1 metals (e.g., Na₂CO₃, K₂CO₃) and ammonium carbonate ((NH₄)₂CO₃), which are soluble.
4. Exceptions for Hydroxides (OH⁻):
Most hydroxides are insoluble. The main exceptions are the hydroxides of Group 1 metals (e.g., NaOH, KOH), which are soluble. Barium hydroxide (Ba(OH)₂) and calcium hydroxide (Ca(OH)₂) are moderately soluble.
In summary, while general rules provide a useful guide, exceptions must be memorized for accurate predictions of salt solubility.
4. Explain why ionic compounds have lattice structures and how this influences their physical properties.
Answer: Ionic compounds form lattice structures because of the strong electrostatic attraction between oppositely charged ions. When a cation and an anion come together, they arrange themselves in a repeating three-dimensional pattern that maximizes attraction and minimizes repulsion. This results in a giant ionic lattice.
Influence on Physical Properties:
High Melting and Boiling Points: The strong electrostatic forces throughout the lattice require a large amount of energy
to break, leading to high melting and boiling points. Hard and Brittle: The lattice is rigid and hard. However, if a force is applied, ions of the same charge can be pushed together, causing repulsion and shattering the crystal. Solubility: Polar water molecules can surround and separate the ions, breaking down the lattice. However, if the electrostatic forces in the lattice are too strong (e.g., in high-charge salts), the compound may be insoluble. Electrical Conductivity: In the solid state, ions are fixed in the lattice and cannot move, so ionic compounds do not conduct electricity. When molten or dissolved, the lattice breaks down, and the free-moving ions can conduct electricity.
5. Describe the preparation, separation, and purification of salts using an excess of an insoluble carbonate, providing an example reaction.
Answer: This method is used to prepare soluble salts from insoluble carbonates (e.g., copper(II) carbonate, calcium carbonate).
Step 1: Reaction:
An excess of the insoluble carbonate is added to a warm acid (e.g., dilute sulfuric acid). The mixture is stirred until no more carbonate dissolves and effervescence (CO₂ gas) stops. Adding an excess ensures that all the acid is used up. Example Reaction: CuCO₃(s) + H₂SO₄(aq) [source symbol unavailable] CuSO₄(aq) + H₂O(l) + CO₂(g)
Step 2: Separation (Filtration):
The mixture is filtered using filter paper and a funnel. The unreacted excess carbonate is collected as the residue on the filter paper. The filtrate (the liquid that passes through) contains the soluble salt (copper(II) sulphate) and water.
Step 3: Purification (Crystallization):
The filtrate is transferred to an evaporating dish. The solution is gently heated to evaporate water until it becomes saturated. It is important not to evaporate to dryness, as impurities may be left behind. The saturated solution is allowed to cool slowly. Pure crystals of the salt form as the solubility decreases.
Step 4: Drying:
The crystals are filtered off and dried between filter papers.
In summary, using an excess of an insoluble carbonate is an effective way to prepare soluble salts, as the unreacted solid is easily removed by filtration.
Which of the following salts are Insoluble: Silver chloride, barium sulphate, sodium nitrate, ammonium carbonate, and lead nitrate.
Answer: The insoluble salts are:
1. Silver chloride (AgCl) - Insoluble (Exception to the chloride solubility rule).
2. Barium sulphate (BaSO₄) - Insoluble (Exception to the sulphate solubility rule).
The soluble salts in the list are:
Sodium nitrate (NaNO₃) - Soluble (All sodium salts are soluble). Ammonium carbonate ((NH₄)₂CO₃) - Soluble (All ammonium salts are soluble). Lead nitrate (Pb(NO₃)₂) - Soluble (All nitrates are soluble).
1. Design an experiment to make magnesium sulphate crystals with suitable reactants.
Answer:
Objective: To prepare pure magnesium sulphate crystals. Reactants: Magnesium ribbon (or magnesium oxide/carbonate) and dilute sulphuric acid. Procedure:
A. Add 50 cm³ of 1M sulphuric acid to a beaker.
B. Heat the acid gently (not to boil).
C. Add magnesium ribbon (or MgO/MgCO₃) in small portions until no more dissolves and effervescence stops. This ensures an excess of the solid reactant.
D. Filter the mixture to remove the excess unreacted magnesium. E. Transfer the filtrate to an evaporating dish. F. Heat the filtrate gently to evaporate water until the solution becomes saturated (a small sample should crystallize when cooled). G. Allow the solution to cool slowly. Crystals of magnesium sulphate will form. H. Filter off the crystals and dry them between filter papers. Equation: Mg(s) + H₂SO₄(aq) [source symbol unavailable] MgSO₄(aq) + H₂(g)
2. Design an experiment to make potassium sulphate crystals using titration method.
Answer:
Objective: To prepare potassium sulphate crystals using titration.
Reactants: Potassium hydroxide (KOH) solution and dilute sulphuric acid (H₂SO₄). Procedure:
A. Titration: Pipette 25.0 cm³ of KOH solution into a conical flask. Add a few drops of phenolphthalein indicator.
B. Fill a burette with H₂SO₄. Add the acid to the flask until the pink color just disappears. Record the volume of acid used.
C. Repeat the titration to obtain an average volume.
D. Preparation: Measure a fresh 25.0 cm³ of KOH into a clean flask. Add the exact calculated volume of H₂SO₄ from the burette. Do not add indicator. E. Crystallization: Evaporate the resulting solution to saturation and allow it to cool slowly to form crystals of K₂SO₄. Equation: 2KOH(aq) + H₂SO₄(aq) [source symbol unavailable] K₂SO₄(aq) + 2H₂O(l)
3. Name two types of salts that are always soluble.
Answer: Two types of salts that are always soluble are:
1. Nitrates (e.g., NaNO₃, KNO₃, AgNO₃).
2. Sodium and Potassium salts (e.g., NaCl, KCl, Na₂SO₄, K₂CO₃). (Note: All ammonium salts are also always soluble.)
4. How could you prepare pure dry sample of copper(II) sulphate from copper(II) carbonate.
Answer:
1. Reaction: Add an excess of solid copper(II) carbonate (CuCO₃) to warm dilute sulphuric acid (H₂SO₄). Stir until effervescence stops.
CuCO₃(s) + H₂SO₄(aq) [source symbol unavailable] CuSO₄(aq) + H₂O(l) + CO₂(g)
2. Separation: Filter the mixture to remove the unreacted excess CuCO₃.
3. Crystallization: Transfer the blue filtrate (CuSO₄ solution) to an evaporating dish. Heat gently to evaporate water until the solution is saturated. Allow it to cool slowly to form blue crystals of CuSO₄·5H₂O.
4. Drying: Filter off the crystals and press them between filter papers to dry them. Do not heat the crystals strongly, as they will lose their water of crystallization and turn white.
