Chapter 13: Halogens

Group 17 (VIIA) Elements (Halogens) Introduction to Halogens • Group 17 or VIIA elements are called halogens. • Fluorine (F), Chlorine (CI), 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 State Color Gas Fluorine (F2) Pale yellow Gas Chlorine (Clz) Greenish yellow Reddish-brown Bromine (Br2) Liquid Solid Iodine (12) Shiny greyish black Astatine (At) and Tennessine (Ts) are • Rare and radioactive • Less is known due to their unstable nature enterine Greenish Purple Orange Yellow Eloding) Fig: Greenish yellow chlorine (left), orange bromine (middle) and purple iodine (right) - Other Properties Highly reactive Poisonous and irritating Volatile, corrosive, toxic fumes Sublimes to violet vapor Keep in mind! 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. Trend in Colour • Fluorine (F2) and chlorine (Cl2) are gases of pale yellow and greenish yellow colors respectively at room temperature and pressure. • Bromine (Br2) is a volatile liquid of reddish brown color having toxic fumes and corrosive. • Iodine is shiny greyish black solid at room temperature. It sublimes directly from solid to violet vapors. • The color darkens from chlorine to iodine, because changes in light absorption caused by electron transitions in the diatomic molecules. Table 13.1 Atomic and physical properties of the common halogens. Fluorine Element 9 Proton number Electron shell structure 2, 7 shell Outer electron 252p$ configuration 19.0 Relative atomic mass (4) gas Physical state at of 20°C Colour pale yellow 220 Melting point/°C 188 Boiling point/°C change Enthalpy + 3.3 vaporization/kJ mol-1 Solubility/g per 100 g of Reacts readily with water water at 20°C VOLATILITY OF CHLORINE, BROMINE AND IODINE

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. Physical State Halogen Volatility at Room Temp Gas Most volatile Chlorine (Clz) Liquid Bromine (Brz) Moderately volatile Least volatile Solid Iodine (12) Interesting Information S.Q. What is bromine? Bromine liquid evaporates easily at room temperatures emitting an orange vapor Ans. Bromine has a very strong and bad odor. It gets its name from the Greek word "bromos" which means "stench." Bromine Chlorine iodine 53 35 17 2,8,18, 18, 7 2,8,7 2,8,18,7 555p5 454p 35 ps 126.9 79.9 35.5 solid liquid gas red-brown dark gray pale green 113 - 7 - 101 183 59 - 35 + 30 + 15 + 10.2 0.59 (reacts 0.018 3.6 slightly) Remarks Spreads quickly in air Evaporates easily, releasing toxic fumes . 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 Instantaneous dipole- Size Polarizability induced dipole forces Small Weak Small High Large Strong 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. Volatility Boiling Point Lower Higher Lower Higher State of Halogens at room temperature Size Halogen Polarizability Small Low Fluorine (F2) Small Chlorine (Clz) Slightly higher Medium Moderate Liquid Bromine (Brz) Solid Large High iodine (I2) • The first two halogens fluorine and chlorine are gases due to weaker id-id forces. • Bromine is a liquid as its size is bigger and it possesses stronger id-id forces than chlorine and fluorine. • Iodine has stronger forces among the group and it is solid at room temperature. Quick Check 13.1 (a) Which halogens elements are radioactive? Ans. Astatine (At) is the only radioactive halogen. (b) What is the reason behind the different colours of halogens? Ans. The different colours of halogens are due to the absorption of light in the visible region, which causes electronic excitation. Each halogen absorbs different wavelengths, and the remaining light gives them their characteristic colour (c) Why chlorine is more volatile than bromine and iodine? Ans. Chlorine is more volatile than bromine and iodine because it has weaker van der Waals forces due to its smaller molecular size and lower molecular mass THE BOND STRENGTH OF HALOGEN MOLECULES

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 CI-Cl bond. Intermolecular Physical Volatility State Forces Gas Weak id-id High (most volatile) forces Gas Weak id-id High forces Stronger id-id Moderate forces Strongest id-id Low (least volatile) forces lp - lp repulsion F F Fig: F2 molecule with lone pairs of electrons 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 Cl2 Br2 Highest bond energy: Cl Lowest bond energy: 12 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) + C/2(g) →2 NaCk(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: 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(g) + 2C (ag) - F2(g) + 2 Br(ag) F2(g) +21 (ag) Chlorine can oxidize and displace Br and I. bromine can oxidizè or displace iodide. iodine cannot oxidize any halide ion. Bond Energy (kJ mol) 156 243 193 151 →→ 2F (a4) + C/2(g) →→ 2F (a9) + Br2(g) → 2F (a4) + 12(g) Table 13.4 Standard electrode potential E° (X2/X) Halogen Molecule (X2) F2 Ck2 Br2 12 The oxidizing power of halogens can be related to standard electrode potential (E°) values. Fluorine the most reactive halogen and most powerful oxidizing agent. Standard reduction potential of fluorine is highest among all elements, so it is the strongest oxidizing agent from periodic table. The standard electrode potential E° become less positive from fluorine to iodine. Decrease in oxidizing power of halogens down the group. Factors affecting oxidizing power: Bond dissociation energy (lower = better oxidizer) (li) Electron affinity (higher = better oxidizer) (iii) Hydration energy of halide ions (higher = better oxidizer) (iv) Heat of vaporization (for Br2, l2) Quick Check 13.2 (a) The F-F bond is weaker than CI-Cl bond although fluorine is the most electronegative element. Explain. Ans. The F-F bond is weaker because of strong repulsion between lone pairs of electrons on the small fluorine atoms, which weakens the bond despite high electronegativity. (b) Is the reaction between NaClaq) and Fz gas possible? Give reason whether yes or no. Ans. Yes, the reaction is possible because fluorine is more reactive than chlorine and can displace it from its salt. (ii) If yes, write the equation for this reaction. Ans. F2(g) + 2NaCl(aq) → 2NaF(aq) + Cl2(g) (c) What is the relationship between the oxidizing power of halogens and their standard reduction potential values? Ans. The higher the standard reduction potential, the greater the oxidizing power of the halogen. Fluorine has the highest value and is the strongest oxidizing agent. REACTIONS OF THE HALOGENS WITH HYDROGEN

Reactions of the halogens with hydrogen All halogens react with hydrogen to form hydrogen halides (HX): H,+ X, →→2HX where X, (F2, Cl. 2, Br, and 12) • These hydrogen halides are colourless gases. When dissolved in water, they form hydrohalic acids: HE, HCL, HBr, HI • Reactivity of halogens with hydrogen decreases down the group: F, > Cl2 > Br, > -2 Standard Reduction Potential, E° (V) + 2.87 + 1.36 + 1.07 + 0.54 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: 12(g) + F2(g) - HF gas forms hydrofluoric acid when dissolved in water. Chlorine (Clz) Reacts with hydrogen in presence of light (UV or spark) to produce colorless Hcl gas which forms hydrochloric acid with water. 2(g) + C/(8) "→2HCL(g) Bromine (Br2) Bromine requires heating to react with hydrogen to give hydrogen bromide. HBr gas is less reactive than Hr and Hcl. H 2(g) + Br2(g) → >2HBr(g) • HBr gas forms hydrobromic acid in water. • The reaction is exothermic. Iodine (12) 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. The reaction is reversible occurs slowly and least favorable among halogens. Reaction Hydrogen Halogen Conditions Halide F2 HF (g) Low temp, dark Cl2 Hcl (g) UV light or spark HBr (g) Br2 Heating 12 HI (g) High temp + catalyst (slow, rev) RELATIVE THERMAL STABILITIES OF HYDROGEN HALIDES IN TERMS OF THEIR BOND STRENGTH

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 →2.HF (g) Acid Formed in Reactivity Level Water Explosive Hydrofluoric acid High Hydrochloric acid Moderate Hydrobromic acid Low Hydroiodic acid 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-Cl is weaker • Bond dissociation energy is 431 kJ/mol which bond than H-F. Hydrogen Bromide (HBr) • Weaker than H-CI 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-CI. 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 H-Cl H-Br H-1 Quick Check 13.3 (a) The reaction between H2 and Fz is explosive but that between H2 and Iz 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. Refer to table 13.4 and 13.5 to predicate which of the following reactions would be more exothermic? • H2(g) + Cl(g) → 2HCl(g) • H2(g) + Brz(g) → 2HBr(g Ans. The reaction with Cl 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. 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. Bond dissociation energy (kJmol) 569 431 366 299 RELATIVE REACTIVITY OF HALIDE IONS AS REDUCING AGENTS Discuss the relative reactivity of halide ions as reducing agents.

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 (Ags). These reactions are used in qualitative analysis to identify halide ions in solution. The general reaction Is: →2AgX(s) (ag) + A8 (ag) Where X = Cl, Bri, 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. A8 (ag) + 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. A8 (ag) + Brag) AgBI(s) Reaction with iodide A yellow precipitate of silver iodide forms on mixing aqueous silver ion with aqueous halide solution. The product Agl is insoluble in ammonia. Ag(a9) + (aq) → Ag(s) This type of reactions is used to identify halide ions and is called silver nitrate test Halide Reaction with Ag Ion *Ke Precipitate No visible reaction (AgF is F- soluble) CI Ag* + CF → AgCI(s) Br Ag+ + Br → AgBr(s) Ag+ + T → Agi(s) 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 diamminesilver (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) + 2NH 3(aq) The cream-coloured precipitate of AgBr dissolves in concentrated ammonia, forming a colourless solution. AgBI(s) + 2NH 3(ag) The yellow precipitate of silver iodide (Agl) 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. White Yellow Cream Fig: Colours of the silver halide precipitates: silver chloride (left, silver bromide (middle) and sılver iodide (right) Colour of Solubility in Ammonia No precipitate Not applicable Soluble in dilute NH3 White precipitate Slightly soluble in Cream precipitate conc. NH3 Insoluble in NH3 Yellow precipitate →[Ag(NH,) 2](ag) + Cl (ag) →(Ag(NH,) Trag) + Brag) Quick Check 13.4 (a) F is a weaker reducing agent than CI. Explain why. Ans. F is more stable and less likely to lose an electron due to the high electronegativity and small size of fluorine, making it a weaker reducing agent than Cl (b) What is the cause of the different solubilities of silver halides in ammonia? Ans. The difference in solubility is due to the varying lattice energies and ability to form complex ions with ammonia. For example, AgCl dissolves in NH3 forming a soluble complex, while AgI does not. (c) Write down the equation for the reaction of KI with Ag followed by NH3. What would you observe at the completion of this reaction? Ans. KI + AgNO3 → Agl (yellow ppt.) + KNO3 Reaction with NH3: Agi does not dissolve in NH3. Observation: A yellow precipitate of Agi remains undissolved after adding ammonia. Table 13.6 Action of Ag+ followed by ammonia on halides Halide ion Action of aq. Ag+ ion Fluoride ion (F) No precipitate Chloride ion (CI) White precipitate (AgCI) Bromide ion (Br) Cream colour precipitate (AgBr) Iodide ion (1) Pale yellow precipitate (Agl) Summary of solubility in ammonia Halide Solubility in Dilute NH: AgX Precipitate F No precipitate Yes CI White (AgCI) No Br Cream (AgBr) No Yellow (Agl) REACTIONS OF HALIDES (X-) WITH CONCENTRATED SULFURIC 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 wit NaF. HF is a weak reducing agent, therefore, it does not react with H2S04 further NaF s) + H,SOA(e) - • Hr gas evolved • No further redox (HF 1s a poor reducing agent) Reaction with chloride When concentrated sulfuric acid reacts with sodium chloride, fumes of hydrogen chloride gas are produced. NaCls) + H,504(e) Action of aqueous Ammonia No reaction with aq. NH3 Soluble in dil. Aq. NH3 Soluble in conc. Aq. NH3 - Insoluble in aq. NH3 Solubility in Concentrated NH: Yes Yes No → NaHSO4(0) + HF (g) → NaHSO4(e) + HC (g) Reaction with bromide In case of (Br) ion, steamy fumes of hydrogen bromide (HBr) gas and brown fumes of bromine (Br) 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., 2HBrs) + H,SO 4(0) → Br (g) + 2502(g) + 2H, 0(6) Reaction with iodide When concentrated sulfuric acid reacts with sodium iodide, fumes of hydrogen iodide gas, purple fumes of solid iodine Iz and smell of H2S gas produced. HI acts as a strong reducing agent. It reduces H2SO4 to H2S and is itself oxidized to 12. Nal (ag) + H,SO 4(e) NaHSO 4(e) + H11 (g) 4(0) 412(s) + H, (8) + 41,0(+) 8HI (s) + H,S0 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 (12) vapors • Hydrogen sulfide (HS) gas (rotten egg smell) • HI is a strong reducing agent Reducing power of hydrogen halide, increasing order of reducing ability down the group. REACTIONS OF CHLORINE WITH COLD AND HOT AQUEOUS SODIUM HYDROXIDE

Reactions of chlorine with cold and hot aqueous sodium hydroxide Disproportionation 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). Cl 2(g) +2 NaOH (a) → NaCl (ag) + NaClO (ag) + H, O(e) Oxidation states of chlorine in the above reaction are given below: = 0 Chlorine in Cl = -1 Chlorine in NaCl +1 Chlorine in NaClO = • Chlorine in Cl2 is oxidized from 0 to +1 in NaOCl. Keep in mind! MCQ: Which of the following was the previous name of sodium chlorate? (a) Sodium hypochlorite (b) Sodium hypochlorite (c) Sodium chloride (d) Sodium hypochlorate • Chlorine in Clz 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 NaCIO) 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. 3C/2(g) + 6NaOH (ag) → 5 NaC(ag) + NaCO3(aq) + 3H, (,) = Chlorine in Cl = Chlorine in NaCl = Chlorine in NaCIO3 • Chlorine is oxidized from 0 to +5 in NaCIO. Chlorine is reduced from 0 to -1 in NaCl. This disproportionation reaction shows how the temperature of the reaction influences the products formed, demonstrating the versatility of chlorine in undergoing redox reactions. • NaCIO 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 NaC1O3) Quick Check 13.5 (a) How would KI react with conc. HSO‹? What does this reaction indicate about the reducing power of iodide? Ans. KI reacts with conc. H2SO4 to produce I2, SOz, HaS, and HI. This shows that I is a strong reducing agent and can reduce H2SO4 to SO2, S, or HaS. Reaction: 8KI + 5H2SO4 → 412 + H2S + 4H20 + 4KzSO4 (b) Show that the reaction of Cl with cold and hot aqueous KOH is a disproportionation reaction. In cold KOH: Ans. Cl + 2KOH → KC1 + KCO + HỌ (Chlorine is both oxidized to CIO and reduced to CI) In hot KOH: 3Cl2 + 6KOH → 5KCl + KCIO3 + 3H20 (Chlorine is oxidized to ClOs and reduced to Cl Disproportionation occurs when the same element is both oxidized and reduced in a reaction. (c) HI acts as strong reducing agent. Explain it with chemical reactions. Ans. HI is a strong reducing agent due to the weak H-I bond and the ability of I to lose electrons easily. 0 - 1 +5 Examples: 1. With H2SO4 2HI + H2SO4 → 12 + SO2 + 2H20 2. With sulfur 2HI + S → H2S + 12 These reactions show that HI readily donates electrons, reducing other substances. USE OF CHLORINE INWATER PURIFICATION

Use of chlorine in water purification • Chlorine gas (Clz) 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 (HOCI) and the chlorate (I) or hypochlorite ions (OCE). • 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 (HCIO). C/(g) + H2O,) → HC (ag) + HCIO (ag) Chloric (I) acid (HOCI) is a weak acid and partially dissociate in water to form hydrogen ions (Ht) and chlorate (I) ion (OCl). • HOCI = hypochlorous acid (also called chloric(I) acid) • HOCI is a weak acid, and it partially dissociates: HOCI H+ + OCT Disinfection Activity • HOCI and OCI are active disinfectants. • HOC! 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 OCT 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 OCT, which disrupt the cell function leading to cell death. • HOCI and OCI can oxidize and inactivate enzymes that are crucial for survival and replication of bacteria. Interesting Information! S.Q. Give antiseptic role of hypochlorite. Ans. Nucleic acids (DNA and RNA) can be oxidized by HOÇI and OCl and thus preventing bacteria from replicating and vital cellular functions. Factors affecting disinfection 1. pH • At pH around 6-7.5, HOCI 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 HOCI and OCI 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. Quick Check 13.6 (a) Why HOCl is more effective disinfectant than OCt to kill bacteria in water? Ans. HOCl is more effective because it is neutral and can penetrate microbial cell walls much more easily than the negativity charged hypochlorite ion (OCE). Once inside, HOCl oxidizes essential cellular components disrupting cell function and causing death. (b) What are the factors that affect disinfection of bacteria in water? Ans. pH: At pH = 6 - 7.5 HOCl is more effective. Chlorine dose: High chlorine dose kill microbes effectively. Contact time: The disinfectant must be in contact with water long enough to pénetrate and inactivate organisms. (c) What are the primary active species in the chlorination of water? Give equation that shows their formation. Ans. Primary active species: Hypochlorous acid (HOCl) and hypochlorite ion (OCl). Chemical equation of formation of primary active species: → HCl (ag) + HClO (aq) Cl 2(g) + H,0(0) -