Unit 13: Halogens — Long Questions

Ans. Group 17 (VIIA) Elements (Halogens)
Introduction to Halogens

• Group 17 or VIIA elements are called halogens.
• Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At) and Tennessine (Ts). General Characteristics
• Halogens are highly reactive non-metals.
• Highly electronegative elements of periodic table.
• They show similar chemical properties.
• Exist as diatomic molecules (X₂) in all phases (solid, liquid, gas). Physical States and Appearance at Room Temperature

Element Fluorine (F₂), State: Gas, Color: Pale yellow, Other Properties: Highly reactive
Element: Chlorine (Cl₂), State: Gas, Color: Greenish yellow, Other Properties: Poisonous and irritating
Element: Bromine (Br₂), State: Liquid, Color: Reddish-brown, Other Properties: Volatile, corrosive, toxic fumes
Element: Iodine (I₂), State: Solid, Color: Shiny greyish black, Other Properties: Sublimes to violet vapor

Astatine (At) and Tennessine (Ts) are

• Rare and radioactive
• Less is known due to their unstable nature

S.Q. What is the meaning of halogen? Give importance.
Ans. The name halogen comes from the Greek words "halos", meaning "salt", and "gen", meaning "to make." The first halogen to be isolated and recognized as an element was chlorine. Despite the fact that chlorine is poisonous, small amount is essential to human health and life in the form of chloride.

Ans. Volatility of Chlorine, Bromine and Iodine
Definition of volatility: Volatility refers to the tendency of a substance to vaporize or turn into gas.
Higher volatility means the substance evaporates more easily at room temperature.
Volatility trends in halogens.

Halogen Chlorine (Cl₂), Physical State at Room Temp: Gas, Volatility: Most volatile, Remarks: Spreads quickly in air
Halogen: Bromine (Br₂), Physical State at Room Temp: Liquid, Volatility: Moderately volatile, Remarks: Evaporates easily, releasing toxic fumes
Halogen: Iodine (I₂), Physical State at Room Temp: Solid, Volatility: Least volatile, Remarks: Sublimes slowly; less evaporation

• Chlorine being a gas is most volatile at room temperature, and it disperse quickly in the air.
• Bromine being a liquid is less volatile than chlorine but more volatile than iodine. It evaporates readily releasing toxic fumes at room temperature.
• Iodine is least volatile among the three. Its solid state show volatility or compared to bromine chlorine at room temperature. Volatility decreases from chlorine to iodine Reason for Decreasing Volatility (i) Increasing molecular mass Heavier molecules evaporate less easily. (ii) Increase in atomic size Larger atoms have greater surface area for attraction. (iii) Stronger intermolecular forces London dispersion forces become stronger down the group, making it harder for molecules to escape into gas phase.

TREND IN VOLATILITY OF THE HALOGENS
Trends in volatility of halogens

• Halogen molecules are non-polar
• Volatility is primarily determined by instantaneous dipole-induced dipole (id-id) forces, also known as London dispersion forces.

Table 13.2 Effect of London forces on physical properties
Size: Small, Polarizability: Small, Instantaneous dipole-induced dipole forces: Weak, Boiling Point: Lower, Volatility: Higher
Size: Large, Polarizability: High, Instantaneous dipole-induced dipole forces: Strong, Boiling Point: Higher, Volatility: Lower

The London dispersion forces depends upon factors

• Molecular Size: Larger atoms/molecules have more surface area and stronger forces.
• Shape: Compact shapes can enhance intermolecular attraction.
• Polarizability: More polarizable molecules have more easily distortable electron cloud showing stronger London forces. Substances which weak id-id forces are smaller and less polarizable molecules having lower boiling point and higher volatility. Larger molecules are more polarizable with strong London dispersion forces have higher boiling points and low volatility. Boiling Point and Volatility
• Volatility is inversely related to boiling point.
• High volatility means Low boiling point.
• Low volatility means High boiling point. Stronger inter molecules forces require more energy to separate the molecules from the liquid phases to gas phase.

Interesting Information
S.Q. What is bromine?
Ans. Bromine liquid evaporates easily at room temperatures emitting an orange vapor. Bromine has a very strong and bad odor. It gets its name from the Greek word "bromos" which means "stench."

Ans. Bond strength of Halogen molecules
When we move down Group 17 (halogens), the bond strength decreases from chlorine to iodine.
This is primarily due to: Increasing atomic size, longer bond lengths and weaker overlap of orbitals.
Fluorine is an Exception, moving from top to bottom in group 17, the bond energy of halogen decreases from chlorine to iodine.
Although fluorine is the smallest and most electronegative halogen, the F–F bond is weaker than Cl–Cl bond.

Reason

• Very small atomic radius.
• High electron-electron repulsion between lone pairs on adjacent F atoms.
• This leads to weakened bond strength despite strong attraction.
• Fluorine is an exception to the general trend of decreasing bond strength down the group.

Table 13.3 Bond energies of halogen molecules.
Halogen Molecule: F₂, Bond Energy (kJ/mol⁻¹): 156
Halogen Molecule: Cl₂, Bond Energy (kJ/mol⁻¹): 243
Halogen Molecule: Br₂, Bond Energy (kJ/mol⁻¹): 193
Halogen Molecule: I₂, Bond Energy (kJ/mol⁻¹): 151

Highest bond energy C / ₂
Lowest bond energy: I₂

RELATIVE REACTIVITIES OF THE HALOGENS AS OXIDIZING AGENTS
Relative reactivities of the halogens as oxidizing agents
Oxidizing Nature of Halogens

• Halogens act as strong oxidizing agents, when they react with metals and all the free most of the non-metals. They gain electron and are converted into negative halide ion in ionic bond formation. 2Na₍ₐ₎ + Cl₂₍ₐ₎ ⟶2NaCl₍ₛ₎

• Oxidizing agent: A substance that gains electrons (is reduced) and causes oxidation in others.

Oxidizing Power

• Oxidizing power decreases down the group
• Fluorine is the most powerful oxidizing agent
• Iodine is the least powerful among halogens. The order of decreasing power as an oxidizing agent is: F₂ > Cl₂ > Br₂ > I₂

Explanation of Trend

• Reactivity is directly related to the ability to gain an electron and form X⁻ ions, when they react with other elements.
• Fluorine show high tendency to require electron and form fluoride.
• Fluorine molecule can oxidize and displace all the halide ion from their solutions to free electrons.

Displacement reactions
F₂₍ₐ₎ + 2Cl⁻₍ₐq₎ ⟶2F⁻₍ₐq₎ + Cl₂₍ₐ₎
F₂₍ₐ₎ + 2Br⁻₍ₐq₎ ⟶2F⁻₍ₐq₎ + Br₂₍ₐ₎
F₍ₐ₎ + 2I⁻₍ₐq₎ ⟶2F⁻₍ₐq₎ + I₂₍ₐ₎

Chlorine can oxidize and displace Br⁻ and I⁻ bromine can oxidize or displace iodide. Iodine cannot oxidize any halide ion.

Ans. Reactions of the halogens with hydrogen
All halogens react with hydrogen to form hydrogen halides (HX):
H₂ + X₂ ⟶2HX where X₂ (F₂, C/₂, Br₂ and I₂)

• These hydrogen halides are colourless gases.
• When dissolved in water, they form hydrahalic acids: HF, HCl, HBr, HI.
• Reactivity of halogens with hydrogen decreases down the group: F₂ > C/ ₂ > Br₂ > I₂

Fluorine reacts explosively, while iodine reacts slowly and reversibly at room temperature.
Fluorine (F₂)
Reacts explosively with hydrogen at low temperature and in the dark:
H₂₍ₐ₎ + F₂₍ₐ₎ ⟶2HF₍ₐ₎

HF gas forms hydrofluoric acid when dissolved in water.
Chlorine (Cl₂)
Reacts with hydrogen in presence of light (UV or spark) to produce colourless HCl gas which forms hydrochloric acid with water
H₂₍ₐ₎ + Cl₂₍ₐ₎ →(hv) 2HCl₍ₐ₎

Bromine (Br₂)
Bromine requires heating to react with hydrogen to give hydrogen bromide. HBr gas is less reactive than Hᵣ and HCl.
H₂₍ₐ₎ + Br₂₍ₐ₎ →(∆) 2HBr₍ₐ₎

• HBr gas forms hydrobromic acid in water.
• The reaction is exothermic. Iodine (I₂) Iodine reacts very slowly with hydrogen at high temperature and in presence of a catalyst to form hydrogen iodide gas, which forms hydroiodic acid with water. H₂₍ₐ₎ + I₂₍ₐ₎ ←(catalyst→) 2HI₍ₐ₎

The reaction is reversible occurs slowly and least favorable among halogens.

Halogen F₂, Reaction Conditions: Low temp, dark, Hydrogen Halide: HF (g), Acid Formed in Water: Hydrofluoric acid, Reactivity Level: Explosive
Halogen: Cl₂, Reaction Conditions: UV light or spark, Hydrogen Halide: HCl (g), Acid Formed in Water: Hydrochloric acid, Reactivity Level: High
Halogen: Br₂, Reaction Conditions: Heating, Hydrogen Halide: HBr (g), Acid Formed in Water: Hydrobromic acid, Reactivity Level: Moderate
Halogen: I₂, Reaction Conditions: High temp + catalyst (slow, rev), Hydrogen Halide: HI (g), Acid Formed in Water: Hydroiodic acid, Reactivity Level: Low

Ans. Relative thermal stabilities of hydrogen halides (HX)
Thermal stability trend
When we move down Group 17, the thermal stability of hydrogen halides decreases:
HF > HCl > HBr > HI

Reason
The bond strength (bond dissociation energy) between hydrogen (H) and halogen (X) is the main factor. Bond dissociated energy decreases down the group so thermal stability of hydrogen halide decreases.

• Stronger bond means higher stability
• Weaker bond means lower stability

Hydrogen Fluoride (HF)

• Most thermally stable hydride.
• Small atomic radius of fluorine.
• High electronegativity of fluorine.
• Strong orbital overlap of orbitals produces a very strong H-F bond.
• Bond dissociation energy is 569 kJ/mol, which is the highest among hydrogen halides.

Hydrogen Chloride (HCl)

• Less thermally stable than HF but more stable than HBr and HI
• Chlorine has larger atomic radius than fluorine: H-C/ is weaker
• Bond dissociation energy is 431 kJ/mol which bond than H-F.

Hydrogen Bromide (HBr)

• Weaker than H-Cl and HF is less than H-F. due to even larger atomic radius of bromine
• There is reduced orbital overlap.
• Bond dissociation energy is 366 kJ/mol which is less than H-F and H-Cl.

Hydrogen Iodide (HI)

• Least thermally stable among the common hydrogen halides.
• Iodine has the largest atomic radius, causing poor orbital overlap due to which hydrogen iodide bond is weakest among hydrogen halides.
• The bond dissociation energy is 299 kJ/mol.

Table 13.5 Bond dissociation energies of hydrogen halides (H-X) bonds
Hydrogen halide (HX): H-F, Bond dissociation energy (kJ/mol⁻¹): 569
Hydrogen halide (HX): H-C/, Bond dissociation energy (kJ/mol⁻¹): 431
Hydrogen halide (HX): H-Br, Bond dissociation energy (kJ/mol⁻¹): 366
Hydrogen halide (HX): H-I, Bond dissociation energy (kJ/mol⁻¹): 299

Quick Check 13.3
(a) The reaction between H₂ and F₂ is explosive but that between H₂ and I₂ is slow and reversible. Explain why?
Ans. The H–F bond is very strong, and the reaction releases a large amount of energy, making it explosive. In contrast, the H–I bond is weaker, and the reaction is less exothermic, making it slow and reversible.

(b) Refer to table 13.4 and 13.5 to predicate which of the following reactions would be more exothermic?

• H₂(g) + Cl₂(g) → 2HCl(g)
• H₂(g) + Br₂(g) → 2HBr(g) Ans. The reaction with Cl₂ is more exothermic because the H–Cl bond formed is stronger than the H–Br bond, and more energy is released.

(c) How is thermal stability of hydrogen halides related to their bond dissociation energies?
Ans. Higher bond dissociation energy means greater thermal stability. So, hydrogen halides with stronger H–X bonds are more thermally stable.

(d) HF is the most thermally stable hydrogen halide. Give reasons.
Ans. HF has the strongest H–F bond due to small size and high electronegativity of fluorine, making it the most thermally stable hydrogen halide.

Ans. Relative reactivity of halide ions as reducing agents
The reducing power of halide ions increases as we move down the Group 17 from fluorine to iodine.
Explanation of the trend
(i) Electronegativity decreasing down the group, as it is inversely related to reducing power.

• Lower electronegativity gives Greater tendency to lose electrons and Stronger reducing agent

(ii) Atomic / Ionic Radius

• Increases down the group, as it is directly related to reducing power.
• Larger ions give lower charge density and Easier to donate electrons
• Smaller ions like F⁻ have high charge density F⁻ is the weakest reducing agent.
• Larger ions like I⁻ have low charge density I⁻ is the strongest reducing agent.

The order of decreasing power as a reducing agent is I⁻¹ > Br⁻¹ > C/⁻¹ >F⁻¹

Fluoride The high electronegativity of fluorine and the strong bond between the fluorine atom and its extra electron make it difficult for fluoride ion (F⁻) to donate electrons. The small size of fluoride ion results in a high charge density, further stabilizing the fluoride ion and reducing its tendency to lose an electron.

Chloride Chloride ion (Cl⁻) is a stronger reducing agent than fluoride but still weaker than bromide and iodide. This is due to its larger size and lower electronegativity. Bromide ion: Bromide (Br⁻) is a stronger reducing agent than chloride. Bromine is less electronegative than chlorine. Bromide ion (Br⁻) has a larger ionic radius and lower charge density due to which it easily loses an electron and act as a reducing agent.

Iodide Ion Iodide (I⁻) is much stronger reducing agent than both chloride and bromide. Electronegativity of iodine is much lower than chlorine and bromine. Iodide ion (I⁻) is large in size. This results in a lower charge density and a tendency to donate an electron by the iodide ion, making it strong reducing agent.

REACTIONS OF HALIDES WITH AQUEOUS SILVER ION FOLLOWED BY AQUEOUS AMMONIA

Q7. Give the reactions of halides with aqueous silver ion followed by aqueous ammonia.
Ans. Reactions of halides with aqueous silver ions (Ag⁺)
The Silver Nitrate Test
Halide ions react with aqueous silver nitrate (AgNO₃) to form insoluble silver halides (AgX).
These reactions are used in qualitative analysis to identify halide ions in solution. The general reaction is:
X⁻₍ₐq₎ + 4g⁺₍ₐq₎ ⟶2AgX₍ₛ₎

Where X = Cl⁻, Br⁻, I⁻ (F⁻ does not precipitate)

Reaction with fluoride
No reaction is visible when silver ions and fluoride ions are mixed in the aqueous medium. As silver fluoride is soluble in water, fluoride (F⁻) does not form precipitate as silver fluoride.

Ans. Reactions of halides with aqueous silver ions (Ag⁺)

The Silver Nitrate Test
Halide ions react with aqueous silver nitrate (AgNO₃) to form insoluble silver halides (AgX).
These reactions are used in qualitative analysis to identify halide ions in solution. The general reaction is:
X⁻₍ₐq₎ + 4g⁺₍ₐq₎ ⟶2AgX₍ₛ₎

Where X = Cl⁻, Br⁻, I⁻ (F⁻ does not precipitate)

Reaction with fluoride
No reaction is visible when silver ions and fluoride ions are mixed in the aqueous medium. As silver fluoride is soluble in water, fluoride (F⁻) does not form precipitate as silver fluoride.

Reaction with chloride
When silver ions and chloride ions are mixed in the aqueous medium a white precipitate of silver chloride forms, which is soluble in dilute ammonia.
Ag₍ₐq₎ + Cl⁻₍ₐq₎ ⟶AgCl₍ₛ₎

Reaction with bromide
On mixing of aqueous silver ions with aqueous Bromide ions, a cream-coloured precipitate of silver bromide forms. This precipitate of silver bromide is sparingly soluble in concentrated ammonia.
Ag⁺₍ₐq₎ + Br⁻₍ₐq₎ ⟶AgBr₍ₛ₎

Reaction with iodide
A yellow precipitate of silver iodide forms on mixing aqueous silver ion with aqueous halide solution. The product AgI is insoluble in ammonia.
Ag₍ₐq₎ + I⁻₍ₐq₎ ⟶AgI₍ₛ₎

This type of reactions is used to identify halide ions and is called silver nitrate test.

Reaction of silver halides (AgX) with aqueous ammonia
Addition of aqueous ammonia tests the solubility of initially formed silver halide precipitates. AgF is soluble in water so it does not form precipitate, so ammonia has no effect. Silver chloride dissolves in dilute ammonia to form diamminesiiver (I) complex. The white precipitate of AgCl dissolves in dilute ammonia forming a colourless solution.
Addition of aqueous ammonia tests the solubility of initially formed silver halide precipitates.

AgCl₍ₛ₎ + 2NH₃₍ₐq₎ ⟶[Ag(NH₃)₂]⁺₍ₐq₎ + Cl⁻₍ₐq₎

The cream-coloured precipitate of AgBr dissolves in concentrated ammonia, forming a colourless solution.

AgBr₍ₛ₎ + 2NH₃₍ₐq₎ ⟶[Ag(NH₃)₂]⁺₍ₐq₎ + Br⁻₍ₐq₎

The yellow precipitate of silver iodide (AgI) does not dissolve in both dilute and concentrated ammonia, so there is no change in the presence of ammonia.
The below sequence of reactions in table 13.6 provides a systematic way to differentiate between halide ions using their solubility with aqueous silver ions and ammonia.

Ans. Reactions of halides (X⁻¹) with concentrated sulphuric acid

The reactions of different halide ions with concentrated sulfuric acid are different from one another The nature of the product and nature of reaction changes down the group from Fluoride to Iodide.

Reaction with fluoride
Fumes of hydrogen fluoride gas are produced when concentrated sulfuric acid reacts with NaF. HF is a weak reducing agent, therefore, it does not react with H₂SO₄ further.
NaF₍ₛ₎ + H₂SO₄₍ₗ₎ ⟶NaHSO₄₍ₛ₎ + HF₍ₐ₎

• HF gas evolved
• No further redox (HF is a poor reducing agent)

Reaction with chloride
When concentrated sulfuric acid reacts with sodium chloride, fumes of hydrogen chloride gas are produced.
NaCl₍ₛ₎ + H₂SO₄₍ₗ₎ ⟶NaHSO₄₍ₛ₎ + HCl₍ₐ₎

Reaction with bromide
In case of (Br⁻) ion, steamy fumes of hydrogen bromide (HBr) gas and brown fumes of bromine (Br₂) are produced along with the smell of sulfur dioxide (SO₂). It is a redox reaction. Hydrogen bromide acts as a reducing agent by reducing sulfuric acid to sulfur dioxide and is self-oxidized to bromine.

NaBr₍ₛ₎ + H₂SO₄₍ₗ₎ ⟶NaHSO₄₍ₛ₎ + HBr₍ₐ₎

2HBr₍ₐ₎ + H₂SO₄₍ₗ₎ ⟶Br₂₍ₐ₎ + 2SO₂₍ₐ₎ + 2H₂O₍ₗ₎

Reaction with iodide
When concentrated sulfuric acid reacts with sodium iodide, fumes of hydrogen iodide gas, purple fumes of solid iodine I₂ and smell of H₂S gas produced. HI acts as a strong reducing agent. It reduces H₂SO₄ to H₂S and is itself oxidized to I₂.

NaI₍ₐq₎ + H₂SO₄₍ₗ₎ ⟶NaHSO₄₍ₛ₎ + HI₍ₐ₎

8HI₍ₛ₎ + H₂SO₄₍ₗ₎ ⟶4I₂₍ₛ₎ + H₂S₍ₐ₎ + 4H₂O₍ₗ₎

The trend of increasing reducing power of the halide ions (X⁻) from fluoride to iodide, leads to more complex reactions with concentrated sulfuric acid.

• Purple iodine (I₂) vapors
• Hydrogen sulfide (H₂S) gas (rotten egg smell)
• HI is a strong reducing agent

Reducing power of hydrogen halide, increasing order of reducing ability down the group.

Ans. Reactions of chlorine with cold and hot aqueous sodium hydroxide

Disproportioniation Reactions

Definition Those reactions in which a single element undergoes both oxidation and reduction simultaneously are known as disproportionation reactions.

Example When chlorine reacts with cold and hot aqueous NaOH, it undergoes disproportionation and forms different products according to temperature of reaction.

Reaction with cold aqueous sodium hydroxide
Chlorine undergoes disproportionation when it reacts with cold aqueous sodium hydroxide (NaOH) forming sodium chloride (NaCl) and sodium chlorate (I) (NaClO).

Cl₂₍ₐ₎ + 2NaOH₍ₐq₎ ⟶NaCl₍ₐq₎ + NaClO₍ₐq₎ + H₂O₍ₗ₎

Oxidation states of chlorine in the above reaction are given below:

Chlorine in Cl₂ = 0
Chlorine in NaCl = -1
Chlorine in NaClO = +1

• Chlorine in Cl₂ is oxidized from 0 to +1 in NaOCl.
• Chlorine in Cl₂ is reduced from 0 to -1 in NaCl.
• The above reaction shows the simultaneous oxidation and reduction of chlorine and is example of disproportionation reaction.

Disproportionation occurs because

• One Cl atom is reduced (0 → -1 in NaCl)
• One Cl atom is oxidized (0 → +1 in NaClO)

Reaction with Hot Aqueous Sodium Hydroxide
When chlorine reacts with hot aqueous sodium hydroxide (NaOH), it forms sodium chloride (NaCl) and sodium chlorate (V) NaClO₃). It is another example of disproportionation reaction.

3Cl₂₍ₐ₎ + 6NaOH₍ₐq₎ ⟶5NaCl₍ₐq₎ + NaClO₃₍ₐq₎ + 3H₂O₍ₗ₎

Chlorine in Cl₂ = 0
Chlorine in NaCl = -1
Chlorine in NaClO₃ = +5

• Chlorine is oxidized from 0 to +5 in NaClO₃. Chlorine is reduced from 0 to -1 in NaCl.
• Thus disproportionation reaction shows how the temperature of the reaction influences the products formed, demonstrating the versatility of chlorine in undergoing redox reactins.

• NaClO₃ is a powerful oxidizer used in disinfectants and explosives

Disproportionation occurs because

• Some Cl is reduced (0 → -1 in NaCl)
• Some Cl is oxidized (0 → +5 in NaClO₃)

Ans. Use of chlorine in water purification

• Chlorine gas (Cl₂) is poisonous, but in small amounts, it's safe for humans and lethal to microbes.
• It is widely used in water treatment plants and swimming pools.

Chlorine gas is added to water

• Chlorine gas is very poisonous. However, in small quantities, they are harmless to humans but poisonous to the bacteria which cause diseases.
• Due to its strong disinfectant properties, chlorine is widely used at the treatment plants for water purification.
• The process involves adding chlorine to water, where it forms active species that kill bacteria and other pathogens.
• Water in the swimming pools is also chlorinated with slightly higher concentrations of chlorine because there is likely to be a higher concentration of bacteria in the water.
• The primary active species are chloric (I) acid or hypochlorous acid (HOCl) and the chlorate (I) or hypochlorite ions (OCl⁻).
• Chlorination is a relatively inexpensive method of water disinfections.

Chlorine addition to water
When chlorine gas (Cl₂) is added to water, it undergoes hydrolysis to form a mixture of hydrochloric acid (HCl) and chloric (I) acid (HClO).

Cl₂₍ₐ₎ + H₂O₍ₗ₎ ⟶HCl₍ₐq₎ + HClO₍ₐq₎

Chloic (I) acid (HOCl) is a weak acid and partially dissociate in water to form hydrogen ions (H⁺) and chlorate (I) ion (OCl⁻).

• HOCl = hypochlorous acid (also called chloric(I) acid)
• HOCl is a weak acid, and it partially dissociates: HOCl ⟷ H⁺ + OCl⁻

Disinfection Activity

• HOCl and OCl⁻ are active disinfectants.
• HOCl is more effective because it is neutral and it can easily penetrate bacterial cell walls. It oxidizes essential molecules (proteins, lipids, enzymes).

Both species destroy Cell membranes, enzymes and nucleic acids (DNA & RNA) can be oxidized by HOCl and OC⁻ prevents bacterial replication and vital cellular function.

• HOCl and OCl⁻ are effective disinfectants, but HOCl is more effective due to its neutral charge.
• The neutral charge of HOCl allows to penetrate the cell walls of micro-organisms easily.
• Essential cellular components such as proteins and lipids are oxidized by HOCl and OCl⁻, which disrupt the cell function leading to cell death.
• HOCl and OCl can oxidize and inactivate enzymes that are crucial for survival and replication of bacteria.

Factors affecting disinfection

1. pH

• At pH around 6-7.5, HOCl predominates and makes the disinfection process more effective.
• At higher pH (above 7.5), OCl⁻ predominates. It is less effective but still provides disinfection.

2. Chlorine Dose

• The higher the amount of chlorine, the more effective the disinfection. Sufficient chlorine must be added to get enough HOCl and OCl⁻ to kill bacteria.

3. Contact Time

• Contact time of water with chlorine must be long enough, to allow the disinfectants to penetrate and kill bacteria, viruses and protozoa.