• Melting point & boiling point
• Color intensity
• Atomic and ionic radius
• Density

• Ionization energy
• Electronegativity,  F<Cl<Br<I<At
• Oxidizing power

• Electrolysis of HF

• Electrolysis of NaCl (Down Cell)
• Electrolysis of brine in Kellner-Solvay method or Chlor-alkali cell

• Passing Cl₂ gas into concentrated sea water

• Ability of halogen as oxidizing agent decrease down the group.

• X₂ + H₂S --> 2HX + S   (X=Cl, Br, I)
• 
  

• Except F, halogen react with water in disproportionation reaction (undergo both oxidation and reduction).

• General disproportionation reaction for Cl,Br and I. Hypohalite (XO^-) tend to disproportionate further.

• Halogen form covalent bond in compound with non-metallic element.
• Aqueous solution of non-metallic halide undergo hydrolysis reaction.
• CCl₄ does not hydrolyze because C has no d orbital to receive a pair of electron from water molecule.

• Halide of s-block metals are ionic
• MX₂ have some covalent character
• Halide of d-block and p-block metals in low oxidation state are ionic and in high oxidation state show some covalent character

• H₂ + X₂ —> 2HX

• 2NaX+H₂SO₄ —> Na₂SO₄ + 2HX (X=Cl, Br, I)
• 2HX+ H₂SO₄ —> 2H₂O + SO₂ + X₂        (X= I, Br)
• 3KX+ H₃PO₄ —>K₃PO₄+ 3HX     (X=Cl, Br, I)

• Thermal stability: HF>HCl>HBr>HI
• Acid strength: HF<<HCl<HBr<HI            (The more easier the H+ can be released ,the higher the acidity )
• HF is a weak acid due to the small size and high electronegativity of F, the bond between F and hydrogen is very strong. They have a very high electronegativity difference. That is why F doesn't  let go hydrogen easily, more energy needed to break the bond.

• React with metals above hydrogen in electrochemical series
• React with metal oxides, hydroxides and carbonates
• HX as reducing agent
• Strength as reducing agent              HCl< HBr < HI

• +7 HOXO₃ Perhalic acid -highest oxidation number
• +5 HOXO₂ Halic acid
• +3 HOXO Halous acid
• +1 HXO Hypohalous acid -lowest oxidation number 

• Disproportionation reaction

• Ba(OH)₂ + H₂O₂ + 2ClO₂ -->Ba(ClO₂)₂ + 2H₂O + O₂
• Ba(ClO₂)₂+H₂SO₄ --> BaSO₄ + 2HClO₂

• I₂ + 10HNO₃ --> 2HOIO₂ + 10NO₂ + 4H₂O

• Hypobromite salt unstable
• Hypoiodite ion doesn't exist
• Hypochlorite salt used as bleaching agent,oxidizing agent and disinfectant

• Electrolysis of cold brine
• CaO/CA(OH)₂ with chlorine

• Only NaClO₂ industrially important
• Salt used as bleaching and oxidizing agent
• Salts produced by reaction of ClO₂  with NaOH and H₂O₂

• NaClO₃ + H₂O --> NaClO₄ + H₂
• Potassium and ammonium salt produced by: NaClO₄ + MCl --> MClO₄ + NaCl     (M= K,NH₄⁺)

• Red molecule
• Preparation; anhydride chloric + perchloric acid, ClO₂⁺ & ClO₄^-

• Dehydrating concentrated perchloric acid with P2O5 and distilling (-35C ,1 mm Hg pressure)

• White solid 
• Decompose at temperature > 300 C
• Preparation: dehydrating ionic acid at 200C in a stream of dry air     I₂O₅+5CO —> I₂ + 5CO₂
  
  
  

• melting & boiling point
• colour intensity
• atomic & ionic radius
• density

• IE
• electronegativity
• oxidising power

• oxidising agent (ability decrease from F2, Cl2, Br2 to I2 )
• react with H2S
• react with water ( disproportionation except F2)
• react with alkali solution
• react with non-metallic element (form covalent bonds)
• reaction with metals

• "Hals"-salt,"gennan"-generate. 
• found as salts of the halide ion-high reactivity.
• has 7 valence electrons (ns²np⁵)
• form singly bonded diatomic molecules,X₂ - only need one more electron to achieve stable octet electron arrangement.

1. melting point and boiling point-increase in Van der Waals forces.
2. colour intensity
3. atomic radius and ionic radius
4. density of the elements.

1. ionization energy
2. electronegativity
3. oxidizing power (E°SRP)

• only has 1 oxidation state.

• electrolysis of HF dissolved in molten KF(as electrolyte bcs pure HF is non-conducting molecular compound)
• Anode : 2F^- → F₂ +2e^-/ Cathode : 2H⁺ +2e^-  →  H₂

1. Electrolysis of molten NaCl in Downs Cell. (produce sodium metal in cathode,need extremely high temperature)
2. Electrolysis of brine in Kellner-Solvay method/Chlor-alkali cell(produce chlorine,H₂ and sodium hydroxide,need room temperature).

• Cl₂ + 2Br^-  →  2Cl^- + Br₂

• SO₃^2- + Br₂  →   SO₄^2- + 2Br^-

1. Reacting a solution of the halide ions with any substance that is stronger oxidizing agent.
2. Reaction of metal halide with strong oxidizing agent such as MnO₂ in the presence of acid.
3. Oxidation of conc.HCl with MnO₂ or KMnO₄. (Chlorine only)
4. Reacting bleaching powder,CaCl₂^.Ca(ClO)^₂ or bleaching solution,NaOCl with HCl produced a solution of Cl₂ in water.

• oxidizing agents(high electronegativity, tend to gain electrons)
• in displacement reaction. The ability of oxidized decrease from F₂ to l₂. But F is rarely used as it is difficult to handle)

• Halogen reacts with water in a disproportionation reaction.
• reactants serve as both oxidizing and reducing agents.
• But only F₂ and Cl₂ can oxidize water.

• X₂ + 2OH^-  → X^- + XO^- +H₂O
• 3XO^-  →  2X^- + XO₃^-

• form covalent bonds
• Non-metal halides(aqueous solution) undergo hydrolysis. But CCl4 does not hydrolyze(don't have d orbital to receive a pair of e^- from H₂O)

• halides of s-block metals - ionic.
• halides of d-block and p-block metals (low in oxidation state) - ionic, (high oxidation state) - covalent.

• Cl^- - release HCl gas
• Br^- - HBr,SO₂ and Br₂ liberated.
• l^- - SO₂,H₂S and l₂ liberated.

• Cl^- - produced white ppt. dissolved in NH3.
• Br^- - yellowish-white ppt. dissolved in conc.NH₃.
• l^- - Yellow ppt. sparingly dissolved in conc.NH₃.

• Cl^- - No reaction
• Br^- - Br₂ released and formed brown CCl₄ layer.
• l^- - l₂ released and formed pink CCl₄ layer.

• manufactured industrially by the reaction of fluospar,CaF₂ with sulphuric acid at 200-350℃
• used in manufacture of aluminium fluoride,boron trifluoride,uranium tetrafluoride and ammonium hydrogen fluoride, in manufacture of fluorohalogen hydrocarbons and in pickling stainless steel.

• manufactured industrially by direct reaction of hydrogen and halogen. H₂ + X₂  →  2HX
• HCl - used in pickling of metals,acid treatment of oil and gas wells,neutralization in inorganic and organic chemistry,hydrolysis of protein and carbohydrates,manufacture of ClO2.
• In lab, prepared by reaction chlorine salts with conc. H₂SO₄. (not suitable for HBr and Hl,the hydrides will be oxidize by the acid).
• Can prepared by reacting the phophoric acid and halide salts(Cl,Br,I) in the presence of P₂O₅.

• thermal stability : HF > HCl > HBr > Hl
• Acid strength :  HF << HCl < HBr < Hl 
  • High electronegativity difference, bond between F and H is very strong.Hence F does not let got H easily, more energy needed to break the bond.

• react with metals above H in the Electrochemical series.
• react with metal oxides, hydroxides and carbonates.
• HF reacts with glass.
• Hx as reducing agent.

• Fluorine - only formed hypofluorous acid,HOF.
• Cl,Br,l - form four types of oxoacid~Perhalic acid(+7),halic acid(+5),halous acid(+3),hypohalous acid(+1)

• by Hydrolysis of halogen(disproportionation reaction)
• HOX are weak acid, HOCl>HOBr>HOl
• oxidizing agents.
• Hydrolysis of dichlorine oxide,Cl₂O (industrial preparation)

• Chlorous acid,HOClO more stable,prepared by reacting the chlorine dioxide with hydrogen peroxide in the presence of BaOH.Ba(ClO₂)₂ precipitated and reacted with H₂SO₄ to release chlorous acid.

• only HOlO₂ can be easily prepared by reaction of iodine with nitric acid.

• HOClO₃ and HOlO₃ can be prepared from the catalytic oxidation of halate,XO₃^- followed by the addition of a strong acid.
• HOXO₃ are strong acid and powerful oxidizing agent. It oxidizes organic materials explosively when heated.
• Evaporation of aqueous HOlO₃ gives crystals of orthoperiodic acid H₅lO₆.

• Acid strength of oxoacid oc chlorine :
  •  HOCl < HOClO < HOClO₂ < HOClO₃ 

• prepared by electrolysis of cold brine (15℃, Chlor-alkali industry)
  1. Anode : 2Cl^-  →  Cl₂ + 2e^-
  2. Cathode : 2H₂O + 2e^-  →  H₂ + 2OH^-
• Chlorine gas at anode mixed with hydroxide ion at cathode.
  1. Cl₂ + 2OH^-  →  ClO^- + Cl^- +H₂O

• If the solution is hot, reaction produce mixture of chloride and chlorate.
• Reaction of CaO or Ca(OH)₂ with chlorine

• hypochlorite salt used as bleaching agent, oxidizing agent and disinfectants.

• undergo disproportionate reaction.
  1. 3XO^-  →  2X^- + XO₃^-

• used as bleaching and oxidizing agent,deodorizing agent of odorous industrial emission.
• produced by the reaction of ClO₂ with NaOH and H₂O₂.
  1. 2ClO₂ + 2NaOH + 2H₂O₂  →  2NaClO₂ + 2H₂O + O₂
• if carefully controlled, a mixture of chlorate (+5) and hypochlorite (+1) solution ions can be converted into chlorite (+3) ion solution.
  1. ClO₃^- + ClO^-  →  ClO₂^-

• prepared by electrolysis of hot brine (70℃)
  1. 3X₂ + 6OH^-  →  XO₃^- + 5X^- +3H₂O
• prepared by oxidation of halide ion (-1) by hypochlorite ion (+1) [applicable for preparation of bromate and iodate salt]
• NaClO₃-defoliant and herbicide.
• KClO₃- primary oxidant in fireworks and matches,source of obtaining oxygen gas in lab experiment(with MnO₂)
  1. 2KClO₃  →  2KCl + 3O₂
• lO₃^-,iodates - source for production of iodine.

• mainly utilized in fireworks.

• manufactured from sodium chlorate(+5) by electrochemical synthesis.
  1. NaClO₃ + H₂O  →  NaClO₄ + H₂

• produced by metathesis with sodium perchlorate.
  1. NaClO₄ + MCl  →  MClO₄ + NaCl (M=K,NH₄⁺)
• Ammonium perchlorate - used as oxidation agent in rocket fuel.

• mostly unstable.

• solid phase, mixture of anhydride ClO₂⁺ and ClO₄^-

• melting & boiling point
• colour intensity
• atomic & ionic radius
• density of the elements

• Ionization energy
• electronegativity
• oxidizing power

• melting point & boiling point
• colour intensity
• atomic radius & ionic radius
• density of the elements

• ionization energy
• electronegativity
• oxidizing power

• +7 HOXO₃ (Perhalic acid) 
• +5 HOXO₂ (Halic acid)
• +3 HOXO (Halous acid)
• +1 HXO (Hypohalous acid)

• Ba(OH)₂ + H₂O₂ + 2ClO₂ -> Ba(ClO₂)₂ + 2H₂O + O₂
• Ba(ClO₂)₂ + H₂SO₄ -> BaSO₄ + 2HOClO₂

• I₂ + 10HNO₃ -> 2HOIO₂ + 10NO₂ + 4H₂O

• +1 XO (hypo - ite ion)
• +3 XO₂ ( - ite ion)
• +5 XO₃ ( - ate ion)
• +7 XO₄ (per - ate ion)

• Red molecule
• Preparation - mixture of anhydride chloric and perchloric acid

• Dehydrating concentrated perchloric acid with P₂O₅ and distilling product at -35^oC, 1mm Hg pressure

• White solid 
• Decompose at temperature higher than 300^oC 
• Preparation - dehydrating iodic acid at 200^oC in a stream of dry air

•  ^{HXO (Hypohalous acid) +1}
•  ^{HOXO (Halous acid) +3}
•  ^{HOXO2 (Halic acid) +5} 
•  ^{HOXO3 (Perhalic acid) +7} 

•  ^{-ite (+1,+3)}
•  ^{-ate (+5,+7)}
•  ^{hypo- lowest oxidation no.}
•  ^{per- highest oxidation no.}

•   Hydrogen halides

• Oxoacid of Halogen

• Red molecule
• Preparation - mixture of anhydride chloric and perchloric acid

• Dehydrating concentrated perchloric acid with P₂O₅ and distilling product at -35^oC, 1mm Hg pressure

• White solid 
• Decompose at temperature higher than 300^oC 

• HOX(+1)
• HOXO(+3)
• HOXO2(+5)
• HOXO3(+7)

•  HXO (Hypohalous acid) +1
•  HOXO (Halous acid) +3
•  HOXO2 (Halic acid) +5 
•  HOXO3 (Perhalic acid) +7 

•  -ite (+1,+3)
•  -ate (+5,+7)
•  hypo- lowest oxidation no.
•  per- highest oxidation no.

• Down the group :

• Industrial Preparation

• Electrolysis of molten NaCl in Downs Cell
• Electrolysis of brine in Kellner-Solvay method or using Chlor-alkali cell

• Extracted by passing chlorine gas into concentrated sea water

• Laboratory Preparation

1. Reacting a solution of the halide ions with any substance that is stronger oxidizing agent
2. Reaction of metal halide with strong

• Application of Elemental halogen

• Chemical properties

• Industrial preparation

• Laboratory preparation of HX (X = Cl, Br, I)

• PHYSICAL PROPERTIES

• Oxoacid of Halogen

• OXYANION

• Preparation

• NaCIO3 + H20 -> NaCIO4 + H2
• Potassium and ammonium salt

• Oxides

• Down Group - increases in:

• Down Group - decreases in:

• Chemical Properties

• Compounds of Halogen

• diatomic molecules
• has 7 valance electron

• melting and boiling point
• color intensity
• atomic and ionic radius
• density

• IE
• electronegativity
• oxidizing power

• electrolysis of molten NaCl in Down Cell

• electrolysis of HF dissolved in molten KF

• by passing chlorine gas into conc. sea water

• reacting a solution of halide ion with strong oxidizing agent
• reaction of metal halide with strong oxidizing agent with presence of acid

• Fluorine - as fluorinating agent
• Chlorine- in manufacture of organic and inorganic chemical
• Bromine- in manufacture of organic molecule

• HF, HCl, HBr,HI
• Thermal stability : HF>HCl>HBr>HI
• Acid strength : HI>HBr>HCl>HF

• HXO, Hypohalous acid, +1
• HOXO, Halous acid, +3
• HOXO2 , Halic acid, +5
• HOXO3, Perhalic acid, +7
• The acid strength : HOX<HOXO<HOXO2<HOXO3

• HClO, Hypochlorous acid, +1
• HOClO, Chlorous acid, +3
• HOClO2, Chloric acid, +5
• HOClO3, Perchoric acid, +7

1. ionization energy
2. electronegativity

so physical state of halogen at r.t is
F - colorless gas
Cl - yellowish-green gas
Br - reddish-brown liquid
I - violet solid
At - solid

anode: 2F^- -> F₂ + 2e
cathode: 2H⁺ + 2e -> H₂

graphite anode: 2Cl^- -> Cl₂ + 2e
cylindrical cathode: 2Na⁺ + 2e -> 2NaCl

Cl₂(g) + 2Br^-(aq) -> 2Cl^-(aq) + Br₂(g)

Br₂ + SO₃^2- -> 2Br^- + SO₄^2-

CaCl₂∙Ca(ClO₂)(s) + 4H⁺(aq) + 2Cl^-(aq) -> Ca²⁺(aq) + H2O(l) + Cl₂(aq)

2HCl(aq) + NaOCl(aq) -> NaCl(aq) + H₂O(l) + Cl₂(aq)

Hence, only halogen at higher position can displace halogen at lower position in periodic table.

X₂(g) + H₂S -> 2HX(aq) + S(s) (X=Cl, Br, I)

• disproportionation reaction (except F₂)

• only F₂ and O₂ can oxidize water

• In presence of light, HOCl will decompose

X₂(g) + 2OH^-(aq) -> X^-(aq) + XO^-(aq) + H₂O(l)

3XO^-(aq)(+1) -> 2X^-(aq)(-1) + XO₃^-(aq)(+5)

• halogens form covalent bonds in compounds
• the compounds formed is interhalogen compounds.
• aq solution of non-metal halides undergo hydrolysis reactions

• halides of s-block elements are ionic

• neutral in aq solutions

• MX₂ (M= Be & Mg, X= Cl, Br, I) have some covalent character.
• halides of p-block and d-block metals in low/high oxidation stable are ionic/show some covalent.

1. heat with conc. H₂SO₄
2. AgNO₃ + NH₃
3. Cl₂/acidic NaClO(aq) + CCl₄(l)

• Fluospar (acidspar), CaF2, with H2SO4 at 200°C - 350°C

• Halide salts + conc H₂SO₄

• not suitable as some of the hydride will be oxidize by acid.

• H₃PO₄ + Halide salts (presence of P₂O₅)

Thermal stability decrease
Acid strength increase
- due to high electronegativity difference in H-F bonds.

• Reacts with metals above hydrogen in electrochemical series.
• Reacts with metal oxides, hydroxides & carbonates.
• 2HF + SiO₂ -> H₂SiF₆ + 2H₂O
• HX as reducing agent (HCl < HBr < HI)

• Cl, Br, I form types of oxoacid
• F form only HOF due to its extreme difficulty to oxidize it.

• hydrolysis of halogen: disproportionation reaction
• weak acid. acid strength: HOCl > HOBr > HOI
• Oxidizing agents
• Industrial preparation of HOCl

• by hydrolysis of dichlorine oxide, Cl2O

• Only chlorous acid, HOClO stable
• HOClO prepared by 

• Cl₂ with H₂O₂ in the presence of Ba(OH)₂. Ba(ClO₂)₂ ppted is then react with H₂SO₄ to release HOClO

• prepared by isolated from its aq solution

• reaction of I₂ with HNO₃

• prepared from catalytic oxidation of halate, XO₃^- followed by addition of strong acid.(X= Cl, I)

• Hypohalite ions in aq solution will undergo disproportionate reaction
• rate of reaction: ClO- < BrO- < IO-

Hence, HBrO and HIO doesn't exist as they are very unstable or evenly do not exist.

• Electrolysis of cold (15°C) brine

• Cl₂ produced at anode is allowed to mix with OH- ions generated at cathode

• Reaction of CaO/Ca(OH)₂ with Cl₂

• produced by reaction of ClO2 with NaOH and H2O2
• possible to convert mixture of chlorate (+1) & hypochlorite (+5) into a solution of chlorite (+3) ion.

• Electrolysis of hot (70°C) brine
• Oxidation of X^- ion (-1) by XO₃^- ion (+1) (bromate & iodate salts)

• Na, K, NH₄⁺ perchlorates ✔
• NaClO₄ is manufactured from NaClO₃ by electrochemical synthesis.
• K⁺ & NH₄⁺ salts are produced by metathesis with NaClO₄.

• Cl₂O₆ is highly unstable red molecule
• Cl₂O₇ may be prepared by dehydrating conc. perchloric acid with P₂O₅ and distilling the product at -35°C under 1mm Hg pressure
• I₂O₅ is a white solid compound and decomposes only at T > 300°C. Produced by dehydrating iodic acid at 200°C in a stream of dry air.

• a powerful oxidizing agent

• can estimate I₂ liberated by Na₂S₂O₃ to allow quantitative determination of CO

• m.p & b.p (Van der Waals forces increase)
• colour intensity
• atomic & ionic radius
• density

• I.E.
• Electronegativity
• Oxidizing power

• The number of lone oxygen atom attached to central atom affects the acidity of the oxoacids. Number of lone oxygens ↑, oxoacid becomes a better proton donor,

•  HOCl < HOClO < HOClO2 < HOClO3

• Electrolysis of cold (15°C) brine
• Cl2 produced at anode is allowed to mix with OH- ions generated at cathode
• Reaction of CaO/Ca(OH)2 with Cl2

1. Na, K, NH4+ perchlorates 
2. NaClO4 is manufactured from NaClO3 by electrochemical synthesis.
3. K+ & NH4+ salts are produced by metathesis with NaClO4

• 7 valence electrons (ns²np⁵)
• Diatomic molecules, X₂
• Increases going down the group:

• Decreases down the group

• F only have oxidation state of -1 due to it has the highest electronegativity.

• Industrial Preparation

1. Fluorine : Electrolysis of HF dissolved in molten
2. Chlorine : Electrolysis in

1. Bromine :  Extracted by passing chlorine gas into conc. sea  water. 

• Laboratory Preparation

• Application of Elemental Halogen

1. Fluorine : Mainly used as fluorinating agent for organic compounds.
2. Chlorine : 

• Chemical Properties

• Tests to identify halide ions

1. Add conc.H2SO4 and heat
2. Add AgNO3(aq), followed by NH3(aq)
3. Add Cl2(aq) or acidic NaClO(aq) and then CCl4(l) 

• Compounds of Halogen

• Hypochlorite ion (+1) and its salt

• Halite ion (+3) and its salts

• Halate ion (+5) and its salts 

1. Electrolysis of hot (70 °C) brine. 
2. Oxidation of halide ion (-1) by hypochlorite ion (+1).  This method is applicable for preparation of bromate  and iodate salt.

•  Oxides of Halogen

1.  Dichlorine hexoxide, Cl2O6

• melting point
• colour intensity
• atomic & ionic radius
• density

• ionization energy
• electronegativity
• oxidizing power

1. reacting a solution of halide ions with any substance that is strong oxidizing agent.
2. Reaction of metal halide with strong oxidizing agent in presence of acid.

• The Group VIIA elements are called halogens.
• Halogens are highly reactive nonmetallic elements in the periodic table.
• Astatine is radioactive and rapidly decays to other, least reactive compared to other Group VIIA elements.
• have seven valence electrons in their highest-energy orbitals (ns²np⁵).  Receive an electron to achieve a full octet of eight electrons, so these elements tend to form anions having -1 charges, known as halides. 
• In their elemental form, the halogens form diatomic molecules, X₂, connected by single bonds.  Since all of the halogens have one unpaired electron in their atomic forms, it is easy for them to "pair up" to form diatomic molecules.  The X₂ molecules are nonpolar, so the only interactions between them are fairly weak London forces, but as the size of the atoms increases, the London forces become stronger, increasing their melting and boiling points

1.  melting point & boiling point - caused by the increase in van der Waals forces (London dispersion force) 
2. colour intensity ( become darker from colourless to yellowish green to reddish brown then purple)
3. atomic radius & ionic radius 
4. density of the elements

1. Ionization energy
2. Electronegativity
3. Oxidising power (E°SRP) 
   
   
   

• a colourless gas

1. by the electrolysis of a solution of potassium fluoride in anhydrous hydrofluoric acid

• Uses

1. used as a fluorinating agent for organic compounds 

• Yellowish green gas

1. By electrolysis molten sodium chloride in Down cell which also produces metallic sodium; chlorine is evolved at the anode
2.  Electrolysis of brine in Kellner-Solvay  method or using Chlor-alkali cell (Hydrogen is generated at the cathode and chlorine at the anode)

• Uses

1. Water treatment
2. Production in the paper product
3. Food seasoning
4. Household Disinfectant (household chlorine bleach) 
5. Manufacture of organic and inorganic chemicals 

• Reddish-brown liquid

1.  Reacting the bromide ion with chlorine
2. Directly oxidising bromide with manganese dioxide

• Uses

1. Mainly utilized in the manufacture of organic chemical

1. Add concentrated H₂SO₄ and heat
2. Add AgNO₃(aq), followed by NH₃(aq)
3. Add Cl₂(aq) or acidic NaClO(aq) and then CCl₄ (l) 

 Hydrogen Halides (HX) 

HF

• Uses

1. used for etching glass and metal
2. used to make refrigerants
3. Pickling stainless steel 

HCl, HBr, HI

1.  direct reaction of hydrogen and halogen
2. reaction of phosphoric acid (a poor oxidizing agent) with halide salts in the presence of P₂O₅
3.  Reaction of halide salt with concentrated sulfuric acid (* Not suitable for HBr and HI)

• Uses

1. HCl: cleaning, pickling, electroplating metals, tanning leather, and refining and producing a wide variety of products
2. HBr:  to make chemicals and drugs, as a solvent and as a veterinary drug 
3. HI:  as a reducing agent and an analytical reagent. It is also used in the manufacturing of pharmaceuticals, disinfectants 

 Oxoacid of Halogen 

• Halogens generally form four series of oxoacids:

1. Hypohalous acids, HOX (+1 oxidation state)
2. Halous acids, HOXO (+3 oxidation state), 
3. Halic acids, HOXO₂ (+5 oxidation state)
4. Perhalic acids, HOXO₃ (+7 oxidation state)

Hypohalous acids: 

Halous acids: 

Halic acids: 

Perhalic acids: 

 Oxyanion and salts of halogen 

1. Hypochlorous acid, ClO^- hypochlorite ion (+1 oxidation state)
2. Chlorous acid, ClO₂^- chlorite ion (+3oxidation state) 
3. Chloric acid, ClO₃^- chlorate ion (+5oxidation state)
4. Perchloric acid, ClO₄^- perchlorate ion (+7oxidation state)

1. Electrolysis of cold (15 °C) brine (Chlor-alkali industry) 
2. Reaction of CaO or Ca(OH)₂ with chlorine 

1.  reaction of ClO₂ with NaOH and H₂O₂ 

 Oxides of Halogen 

1. Dichlorine hexoxide, Cl₂O₆ 

• highly unstable molecule

• white solid compound 
• decomposes only at temperature higher than 300 °C
• As a powerful oxidizing agent, it will quantitatively convert CO to CO₂

• 7 valence electrons
• ns²np⁵
• Diatomic molecules, X₂

• Atomic & Ionic radius
• Melting & boiling point
• Colour intensity
• Density

• IE
• Electronegativity
• Oxidizing power

• Electrolysis of HF dissolved in molten KF

• Electrolysis of molten NaCl in Down cell
• Electrolysis of brine in Kellner-Solvay method or using Chlor-alkali cell

• By passing chlorine gas into concentrated seawater

•  By reacting a solution of the halide ions with any substance that is a stronger oxidizing agent 
•  (Chlorine) by oxidation of concentrated HCl with MnO₂ or KMnO₄ 
•  A solution of Cl₂ in water is readily obtained by reacting bleaching powder, CaCl₂.Ca(ClO)₂ or bleaching solution, NaOCl with HCl 

•  Adding concentrated H₂SO₄ and heat
•  Adding AgNO₃ (aq), followed by NH₃ (aq) 
•  Adding Cl₂ (aq) or acidic NaClO(aq) and then CCl₄ (l)

• HF, HCl, HBr, HI
• Thermal stability: HF>HCl>HBr>HI
• Acid strength : HI>HBr>HCl>HF

• Hypohalous acid, HXO, +1
• Halous acid, HOXO,  +3
• Halic acid, HOXO₂,+5
• Perhalic acid, HOXO₃, +7
• Acid strength: HOX<HOXO<HOXO₂<HOXO₃

• HClO, Hypochlorous acid, +1
• HOClO, Chlorous acid, +3
• HOClO₂, Chloric acid, +5
• HOClO₃, Perchloric acid, +7

• ‌Fluorinating agent for organic compound
• ‌refrigerant fluids
• aerosol propellants
  
  

• ‌in water treatment
• in cleaning and sanitation‌
• in manufacture of organic and inorganic chemicals‌
• in paper and pulp industry
  
  

• ‌manufacture of organic chemicals
  
  

• ‌reacts with metal above hydrogen in the Electrochemical Series
• reacts with metal oxides, hydroxides, carbonates
• HF reacts with glass
• HX as reducing agent

• Dichlorine  hexoxide, Cl₂O₆

• Dichlorine heptoxide

• Diiodine pentoxide, I₂O₅