Periodicity

In this section, you will explore the Period 3 elements and their compounds. You will learn about the patterns in their physical and chemical properties, including their hydrides, oxides, hydroxides, and chlorides. By studying atomic number and electron configuration, you will understand what makes each element and compound unique and how they are related.

KEY IDEAS

• Acid strength is a measure of a compound’s ability to donate protons (H+) in solution.

• Amphoteric compounds are compounds that can react with both acids and bases.

• Bond Strength is the energy required to break a bond in a molecule.

• Density is the mass of a substance per unit volume.

• Electrical conductivity is the ability of a material to conduct electricity.

• Hydrolysis is the reaction of a compound with water.

• Melting and boiling points is the temperatures at which a substance changes from solid to liquid (melting) or liquid to gas (boiling).

• Metallic character is a measure of how easily an element can lose electrons.

• Reactivity is how easily an element reacts with others.

• The Periodic Law states that the properties of elements repeat in a regular pattern based on their atomic numbers and electron arrangements.

• Thermal stability is the resistance of a compound to decomposition at high temperatures.

• Volatility is the tendency of a compound to vaporise or form a gas.

Modern Periodic Law

The modern Periodic law states that the physical and chemical properties of elements are related to their atomic numbers. This means elements with similar electron arrangements have similar properties. These properties include:

1. Chemical properties (like reactivity)

2. Physical properties (like melting points, boiling points, and density)

3. Chemical bonding (like ionic, covalent, or metallic bonds) Understanding these properties helps predict how elements react and why they have certain characteristics.

Period 3 Elements

Period 3 elements are the elements in the third row of the periodic table. They include: Sodium (Na), magnesium (Mg), aluminium (Al), silicon (Si), phosphorus (P), sulphur (S), chlorine (Cl) and argon (Ar).

11 23.0Na ₁₂ ^(24.3)Mg ₁₃ ^(27.0)Al ₁₄ ^(28.1)Si ₁₅ ^(31.0)P ₁₆ ^(32.1)S ₁₇ ^(35.5)Cl ₁₈ ^(39.9)Ar These elements have their outer electrons in the third energy level (shell).

1. Metallic Character

Metallic character is how easily an element loses its electrons.

As you move left to right across Period 3:

• Metallic character decreases.

• This happens because, across the period, elements gain electrons more than they lose them in order to exist with a full outer electron shell (increasing electron affinity).

• Ionization energy increases, making it require more energy to lose electrons.

Example: Sodium (Na), magnesium (Mg), and aluminium (Al) are metals with high metallic character.

Silicon (Si) is a metalloid (a mix of metal and nonmetal properties).

Phosphorus (P), sulphur (S), chlorine (Cl), and argon (Ar) are non-metals with low metallic character.

2. Melting and Boiling Points

a. Melting and boiling points of Period 3 elements:

i. Increase from Na to Si (metals to a metalloid).

ii. Decrease from Si to Ar (non-metals).

Reason Metals like Na, Mg, and Al form metallic bonds, which require a lot of energy to break.

Silicon (Si), a metalloid, forms strong covalent bonds, which also require a lot of energy to break.

Trend in the Melting and Boiling Points of P4, S₈, Cl2, and Ar The melting and boiling points of phosphorus (P4), sulphur (S₈), chlorine (Cl2), and argon (Ar) follow the trend:

S₈>P₄>Cl₂>Ar This trend is primarily influenced by London dispersion (van der Waals) forces, the only intermolecular forces present in these non-polar substances.

Factors Influencing the Trend

1. Molecular Size: Larger molecules have more electrons, resulting in stronger dispersion forces.

2. Molecular Shape: More complex structures provide a greater surface area for intermolecular interactions.

Among these substances, S₈is the largest molecule with the most electrons, leading to the strongest van der Waals forces and the highest melting/boiling points. In contrast, Ar, being a monatomic gas, has the weakest intermolecular forces and the lowest melting/boiling points.

Refer to Figure 5.1 for trends in boiling and melting points across Period 3

Figure 5.1: Trends in Boiling and Melting Points Across Period 3 The blue line represents the melting points of Period 3 elements, showing a rise from Na (Sodium) to Si (Silicon), followed by a decrease for the non-metals (P, S, Cl, Ar).

The red line represents the boiling points, which follow a similar trend: an increase up to Si, followed by a decrease across the non-metals.

3. Density For Period 3 elements (Na, Mg, Al, Si, P, S, Cl, Ar), the trend in density is influenced by atomic mass, atomic size, and crystal structure.

Density Trend Across Period 3

Refer to Figure 5.2 for density variation in Period 3 Elements

Figure 5.2: Density Variation in Period 3 Elements

Explanation of the Graph

The graph illustrates the density trend of Period 3 elements:

1. Increase from Na to Al

a. Na → Mg → Al: Density generally increases as atomic mass increases and atomic size decreases.

b. Silicon (Si) has a relatively high density (compared to other covalent molecules) due to its giant covalent structure.

2. Increase from P to Cl

a. P, S, and Cl have increasing densities due the increase in atomic mass as we go across the period. It should be noted that the density of Phosphorous is variable depending on the allotrope being considered (the most common allotrope is the one quoted here as 1.82 g/cm³).

b. Argon (Ar) has a low density as it is a monatomic gas.

3. Electrical Conductivity

Electrical conductivity is the ability of a substance to conduct electricity. It depends on the number of free electrons in the material.

Trend in electrical conductivity across Period 3 The conductivity decreases from left to right.

Na, Mg, and Al are good conductors because they have many free electrons.

Si is a semiconductor: it can conduct electricity under certain conditions, like when it’s modified in electronics.

P, S, Cl, and Ar are poor conductors (insulators) because they have few free electrons.

Element Atomic number Melting Point (°C) Boiling Point (°C) Density (g/cm³) Electrical Conductivity Metallic Character Sodium (Na) 11 97.8 883 0.968 Good conductor High Magnesium (Mg) 12 650 1090 1.738 Good conductor High Aluminium (Al) 13 660 2470 2.70 Good conductor High Silicon (Si) 14 1414 2900 2.33 Semi conductor Metalloid Phosphorus (P) 15 44 280 1.82 Poor conductor Non- metal Sulphur (S) 16 113 444.6 2.07 Poor conductor Non- metal Chlorine (Cl) 17 -101 -34 3.16 Poor conductor Non- metal Argon (Ar) 18 -189 -185 1.40 Poor conductor Non- metal Chemical Properties of Compounds of Period 3 Elements Across Period 3, elements go from metals (that easily lose electrons) to metalloids (with properties of both metals and nonmetals) to nonmetals (that tend to gain electrons). The reactivity of these elements decreases as you move from left to right, meaning the metals are more reactive than the nonmetals.

Reaction with Water

Sodium reacts very violently with water to form sodium hydroxide (NaOH) and hydrogen gas:

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) Magnesium reacts slowly with cold water to form magnesium hydroxide (Mg(OH)2) and hydrogen gas. It reacts more vigorously with hot water (steam) to form magnesium oxide (MgO) Mg(s)+ 2H₂O(l) → Mg(OH)₂(s) + H₂(g) Mg(s) + H₂O(g) → MgO(s) + H₂(g) Reaction with Oxygen Most of the elements in Period 3 react with oxygen:

Sodium forms sodium oxide (Na2O):

4Na(s) + O₂(g) → 2Na₂O(s) Magnesium forms magnesium oxide (MgO): 2Mg(s) + O₂(g) → 2MgO(s) Aluminium forms aluminium oxide (Al2O3): 4Al(s) + 3O₂(g) → 2Al₂O₃(s) Silicon forms silicon dioxide (SiO2): Si(s) + O₂(g) → SiO₂(s) Phosphorus forms phosphorus pentoxide (P4O₁₀): P₄(s) + 5O₂(g) → P₄O₁₀(g) Sulphur forms sulphur dioxide (SO2): S₈(s) + 8O₂(g) → 8SO₂(g) Chlorine does not react with oxygen under normal conditions.

Argon does not react with oxygen because it is a noble gas with a stable electron configuration.

Reaction with Chlorine

Chlorine (Cl2) is a highly reactive gas, and the elements in Period 3 react with it in different ways:

Sodium reacts strongly with chlorine to form sodium chloride (NaCl):

2Na(s) + Cl₂(g) → 2NaCl(s) Magnesium reacts with chlorine to form magnesium chloride (MgCl2):

Mg(s) + Cl₂(g) → MgCl₂(s) Aluminium reacts with chlorine to form aluminium chloride (AlCl₃):

2Al(s) + 3Cl₂(g) → 2AlCl₃(s) Silicon reacts with chlorine to form silicon tetrachloride (SiCl4):

Si(s) + 2Cl₂(g) → SiCl₄(s) Phosphorus reacts with chlorine to form phosphorus trichloride (PCl3):

P(s) + 3Cl₂(g) → PCl₃(s) Sulphur reacts with chlorine to form sulphur dichloride (SCl2):

S(s) + Cl₂(g) → SCl₂(s) Chlorine does not react with itself.

Argon is inert and does not react with chlorine.

Activity 5.1 Period 3 Elements

Fill in the table with the corresponding information for each Period 3 element (Sodium, Magnesium, Aluminium, Silicon, Phosphorus, Sulphur, Chlorine, and Argon).

Element Atomic

number Density (g/cm³) Melting point (°C) Boiling point (°C) Atomic radius (pm) Phase at room temperature Na Mg Al Si P S Cl Ar Use the periodic table and other reference materials to complete the table with the correct values for each property.

Activity 5.2 Exploring Periodic Trends (physical and chemical properties) in Period 3 Elements Material needed: Periodic table Steps

1. Mention what you know about metals and non-metals? ame some metals, metalloids and non-metals from the Periodic Table.

2. a. Explain metallic properties.

a. Discuss how periodic properties such as ionization energy, reactivity are measured.

b. Using a chart of Period 3 elements explain the trend: Metallic property decreases from left to right.

3. Mention physical properties of period 3 elements.

a. State how these physical properties measured.

b. Discuss how density and melting/boiling points vary across the period and down a group.

4. Define chemical properties.

5. Go online and watch videos (or observe these demonstrated in class) illustrating reactions:

a. Sodium with water (vigorous reaction).

https://www.youtube.com/watch?v= dmcfsEEogxs

b. Magnesium with air (combusting to form MgO).

https://www.youtube.com/watch?v= NnFzHt6l4z8

c. Chlorine with hydrogen (formation of HCl gas).

https://www.youtube.com/watch?v= MtygiCwnEzw

6. Discuss how reactivity changes:

a. Metals (Na, Mg, Al): Reactivity decreases from left to right.

b. Non-metals (Si, P, S, Cl): Reactivity increases.

7. In mixed-ability groups research using books, internet access, charts, or visual aids, record the element’s physical and chemical properties of Period 3 elements.

Prepare a presentation to share findings with the class.

Activity 5.3 Interpreting trends in the physical properties of Period 3 elements graphs Material needed: Graph sheet, pencil, ruler Steps

1. Use the data (e.g., atomic radius, melting point, boiling point, density) generated Activity 5.1 to plot graphs of each property against the atomic number of the elements.

Example properties to plot:

a. Atomic radius vs. atomic number.

b. Melting point vs. atomic number.

c. Density vs. Atomic number.

2. Examine the graphs and discuss patterns in small groups.

How do the physical properties change as the atomic number increases?

Explain silicon’s high melting point compared to neighbouring elements

3. What is the relationship between melting point and atomic number?

Activity 5.4 Reactions of Period 3 Elements

Materials needed: Periodic table for reference, diagrams or pictures of chemical reactions, access to the internet or books for research, paper and pencils for note-taking, (Optional) demonstration materials: small pieces of sodium, magnesium, and other Period 3 elements (for teacher-led demonstration), safety goggles (for demonstrations) A chemical reaction is when one substance changes into a new substance with different properties (e.g., when sodium reacts with water to form sodium hydroxide and hydrogen gas) Steps

1. In small group research into how period 3 elements (Na, Mg, Al and Si) react with:

a. Water:

i. Does it react quickly or slowly?

ii. What products are formed?

b. Oxygen:

i. Does it form an oxide?

ii. What type of oxide is produced?

c. Chlorine

i. Does it form a chloride?

ii. How reactive is it with chlorine?

2. Use the internet, or books to gather information on:

a. Reaction equations (e.g., 2Na + 2H₂O → 2NaOH + H₂).

b. Type of bond formed (ionic, covalent, etc.).

c. Reaction observations (e.g., gas released, heat produced).

3. In small groups present your findings, showing the reaction equations and explaining:

a. How your assigned element reacts with water, oxygen, and chlorine.

b. How the reactivity changes as you move across the period.

c. If your element is a metal, metalloid, or nonmetal, and how this affects its reactions.

4. Discuss:

a. Why do metals like sodium and magnesium react more easily with water to non-metals like sulphur and chlorine?

b. Why does argon not react with other substances?

5. Prepare a multimedia presentation to showcase your findings.

Slideshows (e.g., Google Slides, PowerPoint), short videos (e.g., using apps like C or Flipgrid), Posters (physical or digital).

Include visuals (e.g., reaction photos, diagrams), videos of reactions (if online).

Physical and Chemical Properties of Hydrides of

Period 3 Period 3 hydrides are compounds made by combining hydrogen with the elements in the third period of the periodic table. Examples include sodium hydride (NaH), magnesium hydride (MgH₂), aluminium hydride (AlH₃), silicon hydride (SiH₄), phosphorus hydride (PH₃), hydrogen sulphide (H₂S), and hydrogen chloride (HCl).

Physical Properties

1. Bonding and Structure

NaH and MgH₂

These are ionic hydrides where hydrogen forms H– ions.

They have high melting and boiling points because of strong ionic bonds.

AlH₃:

A covalent hydride with some ionic character and a polymeric structure.

SiH₄, PH₃, H₂S, HCl

These are covalent hydrides with simple molecular structures.

They have lower melting and boiling points.

2. Melting and Boiling Points

Ionic hydrides (NaH, MgH₂): Have high melting and boiling points due to strong bonds.

Covalent hydrides (SiH₄, PH₃, H₂S, HCl): Have lower melting and boiling points due to weaker intermolecular forces (Van der Waals forces).

3. Solubility in Water

Ionic hydrides (NaH, MgH₂): React with water to form hydrogen gas and hydroxides.

AlH₃: Reacts with water to form aluminium hydroxide and hydrogen gas.

Covalent hydrides (SiH₄, PH₃, H₂S, HCl): Are generally soluble in water, forming acidic or basic solutions.

4. Thermal Stability

NaH and MgH₂: Thermally stable and decompose only at high temperatures.

AlH₃: Less stable than NaH and MgH₂.

Covalent hydrides (SiH₄, PH₃, H₂S, HCl): Decompose at lower temperatures.

Summary of Physical Properties of Period 3 Hydrides

The properties of Period 3 hydrides change as you move from metals to non- metals:

• Metal hydrides (ionic): High melting points and highly reactive with water.

• Non-metal hydrides (covalent): Lower melting points with varied reactivity.

• This shift reflects the change from ionic bonding in metals to covalent bonding in non-metals.

Chemical Properties

1. Reactivity: Reactivity increases from NaH to AlH₃.

AlH₃is very reactive and can even catch fire when exposed to air!

2. Acid-Base behaviour: Hydrides become more acidic as you go from NaH to AlH₃.

AlH₃acts like a strong acid.

3. Redox Reactions: Hydrides become better at giving away electrons (reducing) NaH to AlH₃.

Physical and Chemical Properties of Oxides of

Period 3 Period 3 oxides are compounds formed between oxygen and the elements in the third period of the periodic table. Examples include Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₃, and Cl₂O₇.

Physical Properties

1. Structure Ionic oxides: (e.g., Na₂O, MgO) have strong ionic bonds between metal and oxygen ions.

Giant covalent oxides: (e.g., SiO₂) form a continuous network of covalent bonds.

Molecular oxides: (e.g., P₄O₁₀, SO₃) consist of small molecules held by weak forces.

2. Melting and Boiling Points

High: In ionic and giant covalent oxides due to strong bonds.

Lower: In molecular oxides because of weak intermolecular forces.

3. Electrical Conductivity

Ionic oxides: Do not conduct electricity when solid but can conduct when molten because the ions are free to move.

Covalent and molecular oxides: Do not conduct electricity.

Types of Oxides

i. Ionic Oxides Examples: Sodium oxide (Na₂O), magnesium oxide (MgO).

Bonding: Strong ionic bonds between metal cations (e.g., Na+, Mg²+) and anions (O²–).

Melting and Boiling Points: High, due to strong attraction between ions.

ii. Giant Covalent Oxides

Example: Silicon dioxide (SiO₂).

Bonding: Each silicon atom is bonded to four oxygen atoms in a 3D network.

Melting and Boiling Points: Very high, because breaking covalent bonds requires lots energy.

iii. Molecular Oxides Examples: Phosphorus pentoxide (P₄O₁₀), sulphur trioxide (SO₃).

Bonding: Made up of small molecules held by weak Van der Waals forces.

Melting and Boiling Points: Low, because less energy is needed to break these forces.

Summary of Trends

Structure: Ionic → Giant covalent → Molecular.

Melting/Boiling Points: High (ionic and covalent) → Low (molecular).

Electrical Conductivity: Only ionic oxides conduct when molten.

Chemical Properties of Period 3 Oxides

1. Acid-Base Character

Basic oxides: Like Na₂O and MgO, react with acids.

Amphoteric oxides: Like Al₂O₃, can react with both acids and bases.

Acidic oxides: Like SiO₂, P₄O₁₀, and SO₃, react with bases.

2. Reactivity with Water

Basic oxides: Form hydroxides in water. Example:

Na₂O + H₂O → 2NaOH (sodium hydroxide).

Acidic oxides: Form acids in water. Example:

SO₃+ H₂O → H₂SO₄(sulphuric acid).

Physical and Chemical Properties of Period 3

Hydroxides Period 3 hydroxides are compounds formed between elements in Period 3 and hydroxide ions (OH–). Examples include NaOH, Mg(OH)₂, Al(OH)₃, Si(OH)₄, P(OH)₃, H₂SO₄, and HClO.

Physical Properties of Period 3 Hydroxides

1. Solubility in Water

NaOH: Highly soluble; forms a strong alkaline solution because it completely in water.

Mg(OH)₂: Sparingly soluble; forms a weak alkaline solution due to its strong energy.

Al(OH)₃: Insoluble in water but dissolves in both acids and bases because it amphoteric.

Si(OH)₄and P(OH)₃: Generally insoluble due to their covalent bonds, which do dissociate easily in water.

2. Electrical Conductivity

NaOH and Mg(OH)₂: Conduct electricity in water because they produce free ions.

Al(OH)₃, Si(OH)₄, and P(OH)₃: Do not conduct electricity as they do not release ions.

3. Melting and Boiling Points

NaOH: High melting point due to strong ionic bonds.

Mg(OH)₂: High melting point but lower than NaOH due to weaker ionic bonds.

Al(OH)₃: Lower melting point; has both ionic and covalent bonds.

Si(OH)₄and P(OH)₃: Very low melting points because they are molecular with weak intermolecular forces.

Summary of Trends

Ionic hydroxides (NaOH, Mg(OH)₂): High solubility, high melting points, and good electrical conductivity.

Covalent hydroxides (Si(OH)₄, P(OH)₃): Low solubility, low melting points, and poor conductivity.

Amphoteric hydroxide (Al(OH)₃): Can act like both acids and bases, with intermediate properties.

The properties shift from ionic compounds (metals) to covalent compounds (non-metals) as you move across Period 3.

Chemical properties of hydroxides of Period 3

1. Acid-Base Behaviour

Metal Hydroxides (Basic Substances)

Metal Hydroxides are substances made by combining metals with hydroxide (OH–). Most are basic.

Examples Sodium Hydroxide (NaOH): A strong base that fully breaks into its parts in water: NaOH(aq) → Na+(aq) + OH–(aq) It reacts with acids to form salt and water: NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l) Magnesium Hydroxide (Mg(OH)₂): A weak base that only partly breaks in water: Mg(OH)₂⇌ Mg²+ + 2OH– It also reacts with acids to make salt and water: Mg(OH)₂(aq)+ 2HCl(aq) → MgCl₂(aq)+ 2H₂O(l) Aluminium Hydroxide (Al(OH)₃): This is special because it can act as both an acid and a base (amphoteric).

With acids: Al(OH)₃(aq)+ 3HCl(aq) → AlCl₃(aq) + 3H₂O(l) With bases: Al(OH)₃(aq)+ NaOH(aq) → NaAl(OH)₄(aq) Non-Metal Hydroxides (Acidic Substances) Non-Metal Hydroxides are substances that have hydroxide but act like acids.

Examples Sulphuric acid (H₂SO₄) It completely dissociates in water: H₂SO₄(aq)→ 2H+ (aq)+ SO₄²–(aq) It reacts with bases to form salt and water:H₂SO₄(aq)+ 2NaOH(aq)→ Na₂SO₄(aq)+ 2H₂O(l) Hypochlorous Acid (HClO): A weak acid that only partly dissociates in water:

HClO(aq) ⇌ H+(aq) + ClO–(aq) It reacts with bases to make salt and water:

HClO(aq) + NaOH(aq) → NaClO(aq) + H₂O(l)

2. Reactivity Metal Hydroxides

Sodium Hydroxide (NaOH) and Magnesium Hydroxide (Mg(OH)₂) are basic and react with acids to form salts and water.

Aluminium Hydroxide (Al(OH)₃) is amphoteric and reacts with both acids and bases.

Non-Metal Hydroxides

Sulphuric Acid (H₂SO₄) and Hypochlorous Acid (HClO) are acidic and react with bases to make salts and water.

3. Thermal Stability

Trend: Thermal stability increases across the period.

NaOH and Mg(OH)₂: Decompose at lower temperatures due to weaker bonds.

Al(OH)₃and Si(OH)₄: More stable at higher temperatures due to stronger lattice covalent bonds.

Patterns in physical and chemical properties of Period 3 chlorides The chlorides of Period 3 elements are compounds formed between these elements and chlorine. Refer to Table 5.1 for the Chlorides of period 3 and Table 5.2 for the Physical Properties of Period 3 Chlorides.

Table 5.1: Chlorides of period 3 Compound Name of Compound Type of Chloride Acidic, Basic, or Neutral NaCl Sodium chloride Ionic Neutral MgCl₂ Magnesium chloride Ionic Neutral AlCl₃ Aluminium chloride Covalent Acidic SiCl₄ Silicon tetrachloride Covalent Acidic PCl₅ Phosphorus pentachloride Covalent Acidic S2Cl₂ Disulphur dichloride Covalent Neutral

Table 5.2: Physical Properties of Period 3 Chlorides

Property Explanation

Structure - Ionic Chlorides (e.g., NaCl, MgCl₂): Have strong, giant ionic lattice structures.

- Covalent Chlorides (e.g., AlCl₃, SiCl₄, PCl₅, S₂Cl₂):

Have covalent structures. AlCl₃bonds with another molecule to form Al₂Cl₆in solid form.

Melting & Boiling Points

- Ionic Chlorides: NaCl and MgCl₂have high melting/ boiling points because of strong ionic bonds.

- Covalent Chlorides: SiCl₄, PCl₅, and S₂Cl₂have lower melting/boiling points due to weaker forces holding their molecules together.

Electrical Conductivity

- Ionic Chlorides: NaCl and MgCl₂can conduct electricity when melted because their ions can move freely.

- Covalent Chlorides: AlCl₃, SiCl₄, PCl₅, and S₂Cl₂ don’t conduct electricity in any state because they don’t have free-moving ions or electrons.

Chemical Properties of Period 3 Chlorides

Refer to Table 5.3 for chemical properties of Period 3 chlorides.

Table 5.3: Chemical Properties of Period 3 Chlorides

Property Explanation

Reactivity with Water

- Ionic Chlorides (e.g., NaCl, MgCl₂): Dissolve in water to form ions. Example: NaCl(aq) → Na+(aq) + Cl–(aq).

- Covalent Chlorides: React with water (hydrolysis) to form acids.

- AlCl3: Forms HCl and Aluminium hydroxide.

- SiCl4: Forms SiO₂and HCl.

- PCl₅: Forms phosphoric acid (H₃PO₄) and HCl.

Property Explanation

Acid-Base Behaviour

- Ionic Chlorides: Neutral in water.

- Covalent Chlorides: Form acidic solutions when reacting with water.

Thermal Stability of Period 2 and 3 Carbonates

See Table 5.4 for thermal stability of Period 3 carbonates

Table 5.4: Thermal Stability of Period 3 Carbonates

Property Explanation

Carbonates - Thermal Decomposition: Carbonates break down when heated to form metal oxides and CO₂gas.

Example: MCO₃→ MO + CO₂.

- Period 3 Carbonates (e.g., MgCO₃): More stable and harder to decompose due to larger cations like Mg²+.

Thermal Stability of Period 3 Nitrates

Refer to Table 5.5 for thermal stability of Period 3 nitrates

Table 5.5: Thermal Stability of Period 3 Nitrates

Property Explanation

Nitrates - Thermal Decomposition: Nitrates break down when heated.

Example: 2M(NO₃)₂→ 2MO + 4NO₂+ O₂.

- Period 3 Nitrates (e.g., NaNO₃, Mg(NO₃)₂): Decompose to form nitrites or oxides and are more stable due to larger cations.

Note

• Small Cations (like Li+): Make compounds less stable and easier to breakdown with heat.

• Big Cations (like K+): Make compounds more stable and harder to decompose.

• Reactions with Water: Ionic chlorides dissolve; covalent chlorides form acids.

• Decomposition: Carbonates and nitrates break down when heated, giving gases and solids.

Activity 5.5 Period 3 Elements and Their Compounds

Objective: Explore the properties, structure, and behaviour of Period 3 elements and their compounds (hydrides, oxides, hydroxides, and chlorides).

1. Use your textbook or provided handouts to learn about Period 3 elements and compounds. Pay attention to:

a. Identify whether each compound exhibits ionic or covalent bonding.

b. Note acid-base behaviour (basic, neutral, or slightly acidic).

c. Understand their behaviour with water and air.

2. Create a table or chart to summarise the properties of each compound type (hydrides, oxides, hydroxides, and chlorides). Use diagrams to illustrate bonding types applicable.

3. Write a brief explanation of the hydrolytic behaviour of chlorides in Period 3 (e.g., AlCl₃is hydrolysed in water while NaCl is not).

4. Write your thoughts on the trends in the acid-base characteristics of Period 3 oxides you move across the period.

5. Complete a quiz or worksheet that includes questions on bonding, behaviour, and reaction equations for at least three compounds.

Activity 5.6 Analysing the Trends in Thermal Stability of Carbonates and Nitrates Objective: Analyse the trends in thermal stability and their relationship with ionic size and charge density.

1. Read about thermal stability of Group 1 and Group 2 carbonates and nitrates. U additional resources like online simulations or diagrams.

2. Draw a concept map explaining ionic size, charge density, and polarising power, how these factors influence thermal stability.

3. Compare the thermal stability of lithium carbonate, potassium carbonate, carbonate. Write a short paragraph explaining the differences.

4. Write balanced decomposition reactions for NaNO₃, and MgCO₃, observations such as gas evolution.

5. Complete a short quiz on trends in thermal stability, ionic size, and charge density.

Activity 5.7 Investigation on Thermal Stability of Carbonates and Nitrates Objective: Conduct and analyse a hands-on experiment to investigate the thermal stability of carbonates.

1. a. Review, together with you teacher or laboratory instructor, the safety guidelines heating chemicals.

b. Materials required: test tubes, boiling tubes, delivery tubes, test tubes, holders, Na₂CO₃, K₂CO₃, CaCO₃, Bunsen burner, limewater, etc.

2. Heat about 10 g of each solid carbonate in a boiling tube and observe changes. Test CO₂gas using limewater.

3. Record your observations in a table (e.g., colour change, gas evolved, changes limewater).

4. Write the balanced equations for the decomposition of the carbonates tested and discuss the differences in thermal stability based on observations in relation to size and polarising power.

5. Write a short conclusion on the factors affecting thermal stability in carbonates nitrates.

6. Reflect on the experiment by answering guiding questions:

a. Why does calcium carbonate decompose at a higher temperature compared sodium carbonate?

b. How does the test for CO₂confirm decomposition

1. What happens to the atomic radius as you move across Period 3?

2. Which Period 3 element has the highest melting point?

3. How does the trend in ionisation energy change across Period 3?

4. Compare the reactivity of sodium (Na) and chlorine (Cl).

5. Explain why the melting points of Period 3 elements first increase and then decrease.

6. Predict how magnesium (Mg) reacts with water compared to sodium (Na) and why.

7. Analyse how the trends in physical properties across Period 3 are related to electron configuration and bonding.

8. Evaluate how the reactivity trends of Period 3 elements influence their applications.

1. List the Period 3 elements and their corresponding hydrides. (Use a periodic table for reference)

2. Compare the thermal stability of Na₂CO₃, K₂CO₃, and CaCO₃.

3. Having gone through the experiment in this lesson, use the knowledge gained to provide concise answers to the following set of questions as experimental review project.

a. List the correct oxides of the Period 3 elements.

b. Classify each oxide as acidic, basic, or amphoteric based on their chemical properties. Provide a hint for your reasons for the classification.

c. Create a table with the headings like the one below to summarise your findings:

Elements Oxide(s) Classification Explanation

d. Write balanced chemical equations for sodium oxide (Na₂O) and sulphur dioxide (SO₂) reacting with water. the differences in the chemical behaviour of the oxides of sodium (Na₂O) and sulphur (SO₂).

e. Write a short paragraph summarising how bonding and structure influence the acid-base behaviour of these oxides.

In this section, you will explore halogens, the elements in Group 17 of the periodic

table. You will examine their physical and chemical properties, focusing on what makes them unique, how they react, and the patterns they follow as you move down the group.

You will also look at the reactions of halide salts (compounds made with halogens) and explore the strength of hydrogen halide acids, discussing their real- life applications in various industries.

Throughout these lessons, you will use digital tools to enhance your learning and ensure that everyone, regardless of gender or background, has the opportunity to understand how chemistry shapes our world.

KEY IDEAS

• Displacement reaction is a reaction where a more reactive halogen displaces a less reactive halide ion from its compound.

• Halogens are elements in Group 17 of the periodic table.

• Ionic halides are formed when halogens gain one electron (e.g., Cl–, Br–).

• Precipitation reaction of halides is a reaction where halide ions react with silver ions to form insoluble silver halides

• Reducing agent is a substance that donates electrons. Iodide ions are the strongest reducing agents among halide ions.

Halogens are elements in Group 17 of the periodic table. They are special because they exist in all three states of matter at room temperature:

1. Fluorine and Chlorine are gases.

2. Bromine is a liquid.

3. Iodine is a solid.

4. Astatine is radioactive, and not much is known about it.

Halogens are non-metals and usually form diatomic molecules (two atoms of the same element, like Cl₂or F₂), except for Astatine. They are very reactive, so they are always found in nature combined with other elements, never on their own.

Halogens have a general electron configuration of ns² np⁵, meaning they almost have a full outer shell of electrons. They only need one more electron to become stable, like noble gases. This makes them the most reactive non-metals. When they gain an electron, they form halide ions (like Cl–), and they can also form covalent compounds by sharing electrons.

1. Physical Properties of Halogens

Table 6.1 helps you to easily compare halogens’ unique features.

Table 6.1: Physical Properties of Halogens

Property Fluorine

(F) Chlorine (Cl) Bromine

(Br) Iodine (I) Astatine

(At) Physical State

Gas Gas Liquid Solid ^(Solid) (Radioactive) Inter- molecular Forces Weak (small size) Weak Moderate Strong Very Strong Colour Pale yellow Yellow green Reddish- brown Shiny purple black Dark (Unknown) Odour Sharp and strong Sharp and strong Strong and unpleasant Slightly metallic Unknown Property Fluorine (F) Chlorine (Cl) Bromine (Br) Iodine (I) Astatine (At) Solubility in Water Slightly soluble Soluble Slightly soluble Not very soluble Unknown Density Very light Light Heavier than water Heavier than water Very heavy Melting Point (°C)

-219.6 -101.5 -7.2 113.7 ~300 Boiling Point (°C)

-188.1 -34.0 58.8 184.3 ~350 Toxicity Highly poisonous Highly poisonous Highly poisonous Highly poisonous Highly poisonous Bond Energy (kJ/ mol) 159 243 193 151 Low (Estimated ~120) Trends in the Table 6.1 Physical State: Changes from gases to liquid to solids as you move down.

Intermolecular Forces: Get stronger due to increasing molecular size.

Colour: Becomes darker down the group.

Density: Increases with atomic mass.

Oxidising Strength: Decreases from fluorine to iodine.

Bond Energy: Decreases due to lone pair repulsion and larger atomic sizes.

Melting/Boiling Points: Increase with stronger intermolecular forces.

Odour & Toxicity: All halogens have strong smells and are highly poisonous.

2. Chemical Properties of Halogens

a. Reactions with Metals Halogens react with metals to form metal halides, which are ionic compounds.

Example: Sodium + Chlorine → Sodium chloride (table salt).

Na + ½Cl₂→ NaCl These compounds have crystal lattice structures.

b. Reactions with non-Metals Halogens combine with other non-metals to form molecular compounds.

Example: Bromine + Fluorine → Bromine Trifluoride (BrF₃).

1__ 2Br₂+ 3F₂→ BrF₃

c. Reactions with hydrogen Halogens react with hydrogen to form hydrogen halides (e.g., HCl, HF).

Fluorine reacts explosively, while iodine reacts very slowly.

Hydrogen halides dissolve in water to make acids:

HF: Weak acid.

HCl, HBr, HI: Strong acids.

d. Displacement reactions A more reactive halogen can replace a less reactive one in a compound.

Example: Chlorine replaces bromine in potassium bromide:

Cl₂+ 2KBr → 2KCl + Br₂.

Reactivity trend: F₂> Cl₂> Br₂> I₂.

e. Precipitation reactions Halogens react with silver ions to form colourful precipitates, helping identify halides:

Chloride (Cl–): White precipitate (AgCl).

Bromide (Br–): Pale yellow precipitate (AgBr).

Iodide (I–): Bright yellow precipitate (AgI).

f. Oxidising Strength Halogens are strong oxidising agents because they gain electrons easily.

Fluorine is the strongest, followed by chlorine, bromine, and iodine.

Halogens as Strong Oxidising Agents (Simplified Table)

Halogen Pair

(X2/2X−) Standard Reduction

Potential (E°) Oxidizing Strength

F₂___ 2F− +2.87 Strongest oxidising agent Cl₂____ 2Cl− +1.36 Can oxidise I− to I₂ Halogen Pair (X2/2X−) Standard Reduction Potential (E°) Oxidizing Strength Br₂____ 2Br− +1.09 Weaker than Cl₂ I₂___ 2I− +0.54 Weakest oxidising agent The higher the reduction potential (E°), the stronger the oxidising power.

Fluorine (F₂) is the strongest oxidising agent, while iodine (I₂) is the weakest.

Example: Cl₂can oxidise I− in solution because it has a higher E°value.

g. Electronegativity and Bonding Halogens are highly electronegative, meaning they attract electrons easily.

Fluorine is the most electronegative and does not form positive oxidation states.

Heavier halogens (chlorine, bromine, iodine) can show multiple positive oxidation states (e.g., +1, +3, +5, +7).

Remember

1. Halogens react with metals, non-metals, and hydrogen to form important compounds.

2. They are very reactive and strong oxidisers, with fluorine being the strongest.

3. Precipitation reactions with silver ions help identify halides based on their unique colours.

4. Reactivity decreases as you move down the group:

Fluorine > Chlorine > Bromine > Iodine.

Activity 6.1 Analysing Physical States of Halogens

Table 6.2: Data on physical properties of halogens.

Halogen Physical

state Boiling point/ °C Melting point/ °C Bond energy/ kjmol⁻¹Fluorine (F) Gas -188.1 -219.6 -158 Chlorine (Cl) Gas -34.4 -101.5 243 Bromine (Br) Liquid 58.8 -7.2 193 Iodine (I) Solid 184.3 113.7 151 Steps

1. a. In small groups examine Table 6.2 with data on halogens’ physical properties.

b. Observe the variations in physical states, melting/boiling points, and bond for fluorine, chlorine, bromine, and iodine.

c. Record how the properties change across the group, focusing on trends.

2. a. Discuss why fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid at room temperature. (Relate this to the strength of intermolecular forces (Van der Waals forces) increasing down the group).

b. Discuss why melting and boiling points increase down the group (from fluorine to iodine). [Connect this trend to the size of the molecules and the strength of intermolecular forces].

c. Compare the bond energy values for F₂, Cl₂, Br₂, and I₂.

d. Discuss why fluorine has low bond energy (due to strong repulsion between lone pairs) and why bond energies generally decrease as you move down the group.

3. a. Select a property to graph (e.g., melting points, boiling points, or bond energies).

Plot the graph:

Horizontal axis: Halogens (F, Cl, Br, I).

Vertical axis: Selected property (e.g., melting point in °C).

b. Label the graph clearly and identify trends (e.g., increasing boiling points or decreasing bond energies).

Activity 6.2 Exploring Halogen Properties

1. a. Write the electron configuration of halogens (F, Cl, Br, I, At).

b. Explain where halogens are located in the periodic table and why their outermost electron shell is important for their chemical behaviour.

c. Discuss how the halogens’ electron configurations (with seven valence electrons) allow them to achieve stable configurations through gaining or sharing electrons.

d. Identify the typical oxidation states for each halogen and how these states vary based on the type of reaction (e.g., -1 in ionic compounds and positive states in covalent interactions).

e. Highlight common traits like high electronegativity, strong oxidising abilities, and trends in reactivity as you move down the group.

2. a. Research and compile data on standard electrode potential values (E°) for fluorine, chlorine, bromine, iodine, and astatine.

b. Discuss how higher electrode potential values indicate stronger oxidising agents and greater reactivity.

c. Rank the halogens by their reactivity, based on electrode potentials and other chemical trends, from most to least reactive.

Activity 6.3 Halogen Reactions with Water and Alkalis

1. Watch video on halogen reactions, using the link https://youtu.

be/JG8HL9BXN40

2. Discuss and record findings on the:

a. Reagents

b. Conditions

c. products formed

d. When the halogens react with:

i. Water

ii. Alkalis

3. Write balanced chemical equations to show how the reactions occur.

4. Write a comprehensive report on halogen reactions following these guidelines:

a. Introduce halogen reactions with water and alkalis.

b. Reagents, Conditions, analyse each reaction.

c. Provide equations for each reaction.

d. Explain industrial, medical, and environmental applications of halogen reactions.

5. Summarise key points and present it to the class for discussion and feedback.

Reactions of Halide Salts

General Behaviour of Halide Salts

1. Halogens (like fluorine, chlorine, bromine, and iodine) are good oxidizing agents, meaning they can gain electrons to form halide ions (F–, Cl–, Br–, I–).

2. Halide ions (e.g., I–) can lose electrons to form their halogen gases (e.g., I₂).

Example: 2I− → I₂+ 2e− Reducing Power of Halide Ions Iodide (I–) is the strongest reducing agent, meaning it easily loses electrons.

Fluoride (F–) is the weakest reducing agent because it holds its electrons tightly.

The reducing power increases down the group because larger ions (like I–) lose electrons more easily.

Reactions with Concentrated Sulfuric Acid (H₂SO₄)

Refer to Table 6.3 showing how halides react differently based on their reducing power:

Table 6.3: Reducing power of Halides Halide Reaction with H₂SO₄ Products Sodium Fluoride (NaF) Forms HF (hydrofluoric acid). No further reaction.

NaF + H₂SO₄→ HF + NaHSO₄ Sodium Chloride (NaCl) Forms HCl (hydrochloric acid). No further reaction.

NaCl + H₂SO₄→ HCl + NaHSO₄ Halide Reaction with H₂SO₄ Products Sodium Bromide (NaBr) Forms HBr, which reduces H₂SO₄ H₂SO₄+ 2H+ + 2Br− → Br₂ + SO₂+ 2H₂O Sodium Iodide (NaI) Forms HI, which reduces H₂SO₄ H₂SO₄+6H+ +6I− →I₂+ S + 4H₂O

Note

• Fluoride and Chloride: Weak reducing agents, only form HF and HCl.

• Bromide: Moderate reducing agent, forms bromine gas (Br₂) and sulphur dioxide (SO₂).

• Iodide: Strongest reducing agent, forms iodine gas (I₂) and multiple sulphur (SO₂, S, H₂S).

Acid Strengths of Hydrogen Halides

Hydrogen halides are molecules made of hydrogen and halogens: HF (hydrofluoric acid), HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid) Acid Strength of Hydrogen Halides Acid strength depends on how easily the hydrogen halide can break apart (dissociate) in water to release hydrogen ions (H+).

Order of Acid Strength: HF << HCl < HBr < HI HF is a weak acid because it does not dissociate easily.

HCl, HBr, HI are strong acids because they dissociate easily in water.

a. Relative Bond Strengths of Hydrogen halides (HX) The strength of the H–X bond affects how easily it breaks:

Strong bonds (like H–F) make it hard to dissociate weak acid.

Weak bonds (like H–I) make it easy to dissociate strong acid.

Refer to Table 6.4 for relative bond strengths of hydrogen halides.

Table 6.4: Relative Bond Strengths of Hydrogen halides Hydrogen Halide (H–X) Bond Strength (kJ/mol) Acid Strength H–F 565 Weak acid (HF) H–Cl 432 Strong acid (HCl) Hydrogen Halide (H–X) Bond Strength (kJ/mol) Acid Strength H–Br 366 Stronger acid (HBr) H–I 298 Strongest acid (HI) As you move down the group Bond strength decreases.

Molecules break apart (dissociate) more easily.

Acids become stronger.

Remember HF is weak because its bond is very strong and hard to break.

Acid strength increases from HF to HI as bond strength decreases.

HCl, HBr, and HI are strong acids because their bonds are weaker and break easily.

b. Ka Values of Hydrogen Halides Kₐis a number that shows how easily an acid splits into its ions in water.

Kₐtells us how strong an acid is.

The bigger the Kₐvalue, the easier the acid breaks apart and the stronger it is.

For hydrogen halides (HX):

HX(aq) → H+(aq) + X−(aq) Larger Kₐvalues = Stronger acids (dissociate more easily).

Table 6.5 give shows how Kₐvalues relate to the strength of hydrogen halides

Table 6.5: KₐValues of Hydrogen Halides

Hydrogen Halide (HX) Kₐ Value Acid Strength

H–F (Hydrofluoric Acid) 6.6×10⁻⁴Weak acid H–Cl (Hydrochloric Acid) 1.3×10⁶Strong acid H–Br (Hydrobromic Acid) 5.0×10⁹Stronger acid H–I (Hydroiodic Acid) 5.0×10¹⁰Strongest acid Reasons why Kₐ Increase Down the Group Atomic Size As the halogen gets larger (from fluorine to iodine), the bond with hydrogen becomes longer and therefore weaker.

Weaker bonds are easier to break, making acids stronger.

Electronegativity Fluorine is very electronegative and holds electrons tightly, forming a strong H–F bond low Kₐ(weak acid).

Iodine is less electronegative, forming a weaker H–I bond → high Kₐ(strong acid).

c. Thermal Stability of Halogen Halides Thermal stability is how well a compound resists breaking apart (decomposing) when heated.

For halogen halides (HX), it depends on how strong the bond is between hydrogen (H) and the halogen (X).

Trend in Thermal Stability

Thermal stability decreases as you move down the group.

o HF (hydrogen fluoride): Strongest bond → Most stable.

o HI (hydrogen iodide): Weakest bond → Least stable.

Bond Energy and Thermal Stability

See Table 6.6 for bond energy and thermal stability

Table 6.6: Bond Energy and Thermal Stability

Hydrogen Halide (HX)

Bond Energy

(kJ/mol) Thermal Stability

H–F 565 Most stable (requires more heat to break).

H–Cl 432 Stable.

H–Br 366 Less stable.

H–I 298 Least stable (breaks easily with heat).

Trend in Thermal Stability

Bond Strength

The H–X bond becomes weaker down the group (from HF to HI).

Weaker bonds are easier to break when heated.

Size of Halogen Atom

Larger halogen atoms (like iodine) form longer, weaker bonds with hydrogen.

Everyday Uses of Halogens

Halogens are highly useful in everyday life because of their unique chemical properties. They help in cleaning, water purification, refrigeration, and sanitation, improving the quality of life. Here are some key applications:

1. Purification of Water

Chlorine is used in water treatment plants to kill harmful microorganisms and water safe for drinking.

Chemical Reaction: Cl₂(g) + H₂O(l) → HCl(aq) + HClO(aq) Hypochlorous acid (HClO) is the main disinfectant. It breaks bacteria by damaging heir cell walls and essential components, stopping them from reproducing.

Chlorine is also used in swimming pools and healthcare settings for disinfection. Refer Figure 6.1 for the water purification steps.

Figure 6.1: Chlorination in water purification

2. Disinfecting Toilets

Chlorine and iodine are strong oxidising agents that kill bacteria and viruses.

Household bleach (sodium hypochlorite, NaClO is commonly used for toilets.

When dissolved in water, it forms hypochlorous acid, which removes stains and germs effectively.

Figure 6.2: Disinfecting hospital-bed.

3. Bleaching Chlorine is used in bleach to remove colours from dyes and stains.

It oxidises coloured molecules, making them colourless.

4. Halogenated Hydrocarbons

These are compounds containing carbon and one or more halogen atoms (e.g., fluorine, chlorine).

Uses Solvents: Halogenated hydrocarbons like chloroform and dichloromethane fats, oils, and waxes.

Refrigerants: Compounds like CFCs and HFCs are used in refrigeration because of their ability to absorb and release heat efficiently.

Aerosols: Halogenated hydrocarbons are used as propellants in spray cans due their stability and low reactivity.

Note

Many halogenated hydrocarbons are now banned due to their role in global warming and ozone layer depletion.

Activity 6.4 Understanding Reducing Power of Halides

Materials needed: Worksheet containing the key questions Steps

1. What do you think reducing power means?

Have you heard the terms chloride, bromide, and iodide? Where might you find these in daily life?

2. How do halides differ in their reducing power?

Provide the trend.

3. What factors influence reducing power?

4. Present your answers to the class.

5. Discuss as a class:

a. Why iodide has the greatest reducing power.

b. How the size and structure of the ions affect their reducing ability.

Activity 6.5 Comparing reducing power of halides Steps

1. a. In small groups, discuss your ideas about the reducing power of halides (Cl–, Br–, and I–).

b. Which halide do you think is the strongest/weakest reducing agent?

Why?

c. Can you think of any real-life applications of halides’ reactions?

2. Write down the differences in reducing power of halides and includes examples.

3. Write and complete the chemical equations for the reactions of halide salts (chlorides, bromides, and iodides) with concentrated tetraoxosulphate

(VI) acid.

For each reaction: balance the equation, identify the products formed and explain the role of the halide ion as a reducing agent and describe any observable changes during the reaction.

Cl− + H₂SO₄→ Cl₂+ SO₂+ H₂O Br− + H₂SO₄→ I− + H₂SO₄→

Activity 6.6 Exploring Halogens

Materials needed: Textbooks, handouts, or supervised online searches Steps

1. In small groups, investigate the following topics on halogens:

a. Acid strengths of halogens

b. Thermal stability of halogen compounds

c. Everyday uses of halogens

2. a. What is the trend in acid strength or thermal stability among halogens?

b. How are halogens used in daily life?

3. Create visual aids (e.g., tables, graphs, concept maps) to illustrate their findings.

4. Present your findings to the class using your visual aids.

5. Reflect on the following: What did you find most interesting about halogens?

Review Questions 6.1

1. What is the physical state of bromine at room temperature?

2. Name the halogen with the highest melting point.

3. Which halogen is the strongest oxidising agent?

4. What type of bond do halogens form when reacting with metals?

5. Why do halogens become less reactive as you move down the group?

6. Explain why fluorine has a lower bond energy than chlorine.

7. Why do halogens have low melting and boiling points compared to other groups?

8. Predict what happens when chlorine gas reacts with potassium bromide

solution. Write the equation.

9. Identify which halide ion (Cl–, Br–, I–) forms a yellow precipitate when reacted with silver nitrate.

10. What would you observe if iodine is added to a solution of sodium chloride? Explain why.

11. Compare the trend in melting points with the trend in oxidising strength of halogens.

12. Analyse why HF is a weak acid while HCl, HBr, and HI are strong acids.

13. Why does chlorine exhibit multiple oxidation states (e.g., +1, +3, +5, +7) but fluorine does not?

Review Questions 6.2

1. Write the balanced equation for the reaction of bromide salts with tetraoxosulphate (VI) acid.

2. Name two common uses of chlorine in daily life.

3. Explain the trend in acid strength of hydrogen halides (HF, HCl, HBr, HI).

4. Describe the observation when iodide salts react with concentrated tetraoxosulphate (VI) acid.

5. Compare the reactions of chloride, bromide, and iodide salts with tetraoxosulphate (VI) acid in terms of reducing power.

6. How does the variation in acid strength of hydrogen halides relate to their structures?

7. Propose an experiment to demonstrate the differences in reactivity of compounds and relate the results to their uses.

8. Evaluate how the acid strength of hydrogen halides affects their industrial uses.