Group 17 - Inorganic

General Introduction to Group 17 Elements

• Classification: Group 17 elements are known as the halogens. They are p-block elements.

• Electronic Configuration: They possess a characteristic outer-shell electron configuration of $ns^2np^5$.

• Valence Electrons and Bonding:

~ Each halogen has 7 valence electrons, which is one electron short of the stable octet found in noble gases.

~ Because of this, they are highly reactive non-metals.

~ They form ionic compounds when reacting with metals.

~ They form covalent compounds when reacting with non-metals.

• Group Trends: The elements are very similar to one another, displaying a gradual change in properties as the atomic number increases down the group.

Characteristic Physical Properties of Halogens

• Color and State at Room Temperature ($25\,^{\circ}\text{C}$):

  ~ All halogens are colored. The depth of the color increases as the atomic number increases.

~ Chlorine ($Cl_2$): A yellow-green gas.

~ Bromine ($Br_2$): A brown liquid.

~ Iodine ($I_2$): A black solid that produces a purple vapor.

• Boiling Points: [$Cl_2$ : $-35\,^{\circ}\text{C}$] [$Br_2$: $58\,^{\circ}\text{C}$]  [ $I_2$ : $183\,^{\circ}\text{C}$]

• Solubility in Water:

~ Chlorine ($Cl_2$): Moderately soluble in water. Aqueous chlorine turns litmus paper red (due to its acidic nature) and then subsequently bleaches it.

The reaction is: $Cl_2(g) + H_2O(l) \rightarrow Cl^-(aq) + ClO^-(aq) + 2H^+(aq)$.

~ Bromine ($Br_2$): Slightly soluble in water.

~ Iodine ($I_2$): Insoluble in water. However, it dissolves in aqueous potassium iodide ($KI$) solution because of the formation of triiodide ($I_3^-(aq)$) ions. Aqueous iodine is a brown solution.

The reaction is: $I_2(s) + I^-(aq) \rightleftharpoons I_3^-(aq)$.

• Solubility in Organic Solvents: In organic solvents, halogens ($X_2$) exist as relatively free molecules, similar to their behavior in the gas phase.

~ Chlorine ($Cl_2$): Forms a yellow solution.

~ Bromine ($Br_2$): Forms a brown solution.

~ Iodine ($I_2$): Forms a violet or purple solution.

• Observations with Chloroform: Upon adding chloroform to the solutions, the layers separate showing distinct colors: Chlorine appears yellow, Bromine appears brown, and Iodine appears purple.

Volatility and Intermolecular Forces

• Volatility Trend: Volatility decreases down the group. Volatility refers to how easily a substance evaporates.

• Structure: Halogens have a simple molecular structure composed of diatomic covalent molecules ($X_2$).

• Intermolecular Forces: These molecules are held together by instantaneous dipole-induced dipole ($id-id$) forces (also known as London dispersion forces).

• Reasoning for the Trend:

~ As one moves down the group, the molecules become larger.

~ The total number of electrons in the molecule increases.

~ Consequently, the intermolecular $id-id$ forces increase in strength.

~ Increasing strength of intermolecular forces leads to higher boiling points.

~ This explains the physical transition down the group from gaseous $Cl_2$ to liquid $Br_2$ to solid $I_2$.

Chemical Reactivity and Oxidising Ability

• Reactivity Trend: The chemical reactivity of the halogens decreases down the group.

• Reaction with Iron ($Fe$):

~ Chlorine: Reacts vigorously with iron. Heating is only necessary to initiate the reaction.

The equation is: $2Fe + 3Cl_2 \rightarrow 2FeCl_3$.

~ Bromine: Shows a steady reaction. Heating is required continuously throughout the process.

The equation is: $2Fe + 3Br_2 \rightarrow 2FeBr_3$.

~ Iodine: Shows a very slow reaction, even when heated strongly. The equation is: .

• Oxidising Power: All halogens are oxidising agents. The oxidising power decreases down the group (Cl_2 > Br_2 > I_2) because the elements become less reactive.

• Displacement Reactions: Each halide ion can be oxidised by the halogen located above it in the group. A more reactive halogen will displace a less reactive one from its compounds.     *

Example:

$2I^-(aq) + Br_2(aq) \rightarrow I_2(aq) + 2Br^-(aq)$ .

~ In this reaction, bromide displaces iodide; $I^-$ is oxidised to $I_2$ by $Br_2$.

• Specific Displacement Scenarios:

~ Chlorine ($Cl_2$): Displaces Bromine from bromides

($Cl_2 + 2NaBr \rightarrow Br_2 + 2NaCl$)

and displaces Iodine from iodides

($Cl_2 + 2NaI \rightarrow I_2 + 2NaCl$).

~ Bromine ($Br_2$): Displaces Iodine from iodides ($Br_2 + 2NaI \rightarrow I_2 + 2NaBr$)

but has no reaction with chlorides.

~ Iodine ($I_2$): Cannot displace chlorine or bromine; it has no reaction with chlorides or bromides.

Reaction of Halogens with Hydrogen

• General Reaction: All halogens react with hydrogen to form covalent hydrides ($HX$).

~ The general equation is: $H_2 + X_2 \rightarrow 2HX$.

• Reactivity Variations:

~ $H_2 + Cl_2 \rightarrow 2HCl$ : A rapid reaction that explodes in the presence of light.    

~ $H_2 + Br_2 \rightarrow 2HBr$: A slow reaction requiring a temperature of $200\,^{\circ}\text{C}$ and the presence of a Platinum ($Pt$) catalyst.

~ $H_2 + I_2 \rightleftharpoons 2HI$: No reaction occurs unless strongly heated. The reaction is incomplete, resulting in an equilibrium mixture.

Thermal Stability of Hydrides

• Trend: Thermal stability of the hydrides decreases down the group (HCl > HBr > HI).

• Scientific Explanation:

~ As the size of the halogen increases down the group, the $H-X$ bond length increases.

~ Longer bonds are weaker bonds.

~ Weaker bonds require less energy to break, making the hydride easier to decompose thermally.

• Bond Energy Data: [$HCl$ : $431\,kJ\,mol^{-1}$]    [$HBr$: $366\,kJ\,mol^{-1}$] [$HI$: $299\,kJ\,mol^{-1}$]

• Decomposition Observations

~ $HI$ : Has the lowest bond energy and is the most easily decomposed.

Heating results in the observation of the purple vapor of $I_2(g)$.

Equation: $2HI(g) \rightarrow H_2(g) + I_2(g)$.

~ $HBr$ : Decomposes only slightly. Slight browning or orange-brown vapor ($Br_2$) is observed upon heating.

~ $HCl$: Does not decompose; it is stable to heat.

Acid Strength and Ease of Oxidation of Hydrides

• Acid Strength Trend: Acid strength increases down the group (HI > HBr > HCl).

~ Hydrides react with water to form strongly acidic solutions: $HX + H_2O \rightarrow H_3O^+ + X^-$

~ The decrease in $H-X$ bond strength down the group means the bond is more easily broken, allowing $H_3O^+$ and $X^-$ to form more readily.

• Ease of Oxidation Trend: Ease of oxidation increases down the group from $HCl$ to $HI$

~ $HCl$ : Can be oxidised only by strong agents such as acidified $KMnO_4$, concentrated $H_2SO_4$, or $MnO_2$.     

~ $HBr$: Oxidised fairly easily by less powerful agents like concentrated $H_2SO_4$ or $H_2O_2$.

~ $HI$: Very unstable and readily oxidised, even by atmospheric oxygen, to produce $I_2$.

Equation: $4HI + O_2 \rightarrow 2H_2O + 2I_2$. Consequently, $HI$ is a very strong reducing agent.

Identification of Halide Ions ($Cl^-$, $Br^-$, $I^-$)

• Reaction with Silver Ions ($Ag^+$): Addition of aqueous silver nitrate ($AgNO_3$) precipitates silver halides ($AgX$). Equation: $Ag^+(aq) + X^-(aq) \rightarrow AgX(s)$.

• Distinguishing the Silver Halides:

1. Chloride ($Cl^-$): Forms a white precipitate ($AgCl$). It dissolves readily in dilute aqueous ammonia and also in concentrated ammonia to form a colorless solution ($[Ag(NH_3)_2]^+(aq)$).     

2. Bromide ($Br^-$): Forms a cream precipitate ($AgBr$). It is insoluble in dilute ammonia but dissolves in concentrated ammonia to form a colourless solution ($[Ag(NH_3)_2]^+(aq)$).     

3. Iodide ($I^-$): Forms a yellow precipitate ($AgI$). It is insoluble in both dilute and concentrated aqueous ammonia.

Reactions with Concentrated Sulfuric Acid ($H_2SO_4$)

• Chloride ($Cl^-$):

Reaction: $NaCl + H_2SO_4 \rightarrow HCl + NaHSO_4$ .

Observation: White fumes of $HCl(g)$.

Type: Acid-base reaction only (no oxidation occurs).

• Bromide ($Br^-$):

Primary Reaction: $NaBr + H_2SO_4 \rightarrow HBr + NaHSO_4$ (Acid-base).

Secondary Reaction: $2HBr + H_2SO_4 \rightarrow Br_2 + SO_2 + 2H_2O$ (Oxidation).

Observation: Brown fumes ($Br_2$) mixed with white fumes ($HBr$).

• Iodide ($I^-$):

Primary Reaction: $NaI + H_2SO_4 \rightarrow HI + NaHSO_4$ (Acid-base).

Secondary Reaction: $8HI + H_2SO_4 \rightarrow 4I_2 + H_2S + 4H_2O$ (Oxidation).

Observation: Purple vapor ($I_2$), black solid ($I_2$), white fumes ($HI$), and a stinking gas ($H_2S$).

Preparation Note: Because $HI$ is so readily oxidised by $H_2SO_4$, it cannot be prepared this way. Phosphoric acid ($H_3PO_4$), a non-oxidising acid, is used instead: $2NaI + H_3PO_4 \rightarrow 2HI + Na_2HPO_4$.

• Summary of Reducing Power: The reducing power of halides increases down the group (I^- > Br^- > Cl^-).

Reactions of Chlorine with NaOH and Water

• Disproportionation: A reaction where an element is simultaneously oxidised and reduced.

• Chlorine with Cold Aqueous NaOH ($15\,^{\circ}\text{C}$):

~ Product: Sodium chlorate(I), $NaClO$

Equation: $Cl_2(g) + 2NaOH(aq) \rightarrow NaCl(aq) + NaClO(aq) + H_2O(l)$

Oxidation Change: $Cl_2$ ($0$) to $Cl^-$ ($-1$) [reduction] and $ClO^-$ ($+1$) [oxidation]

Uses: Bleach and disinfectant.

• Chlorine with Hot Aqueous NaOH ($70\,^{\circ}\text{C}$):

~ Product: Sodium chlorate(V), $NaClO_3$.

Equation: $3Cl_2(g) + 6NaOH(aq) \rightarrow 5NaCl(aq) + NaClO_3(aq) + 3H_2O(l)$.

Oxidation Change: $Cl_2$ ($0$) to $Cl^-$ ($-1$) [reduction] and $ClO_3^-$ ($+5$) [oxidation].

Uses: Weed killer.

• Chlorination of Water:

Reaction: $Cl_2(aq) + H_2O(l) \rightarrow HCl(aq) + HClO(aq)$

$HClO$ is chloric(I) acid. It kills bacteria to sterilize water.

$HClO$ dissociates to provide $ClO^-(aq)$, which also acting as a sterilizing agent. This makes water safe to drink.