The reactivity of halogenoalkanes is determined by competing factors: the polar nature of the carbon-halogen bond and the strength of the covalent bond. We use relative rates of hydrolysis to demonstrate these properties experimentally, and we study the environmental impact of chlorine radicals on the ozone layer.
🔑 Key Principle
The reactivity of halogenoalkanes is governed by bond enthalpy, not bond polarity. Although the C-F bond is the most polar due to fluorine's high electronegativity, it is the least reactive because the C-F bond is extremely strong (highest bond enthalpy). Conversely, the C-I bond is the weakest (lowest bond enthalpy) and hydrolyses the fastest.
Bond Enthalpy vs Bond Polarity
Two opposing factors affect the rate of reaction of a halogenoalkane:
- Bond Polarity: Electronegativity decreases down Group 7. Therefore, the C-F bond is the most polar (\(\text{C}^{\delta+}\text{-F}^{\delta-}\)) and the carbon atom is the most electron-deficient, which should make it easiest for a nucleophile to attack.
- Bond Enthalpy: Bond enthalpy decreases down Group 7. As the halogen atoms get larger, the overlap of orbitals becomes less effective, making the C-X covalent bond weaker down the group.
Experimental evidence shows that the rate of hydrolysis increases down Group 7. This proves that bond enthalpy is the dominant factor in determining halogenoalkane reactivity.
Key Definitions
The amount of energy required to break one mole of a specific covalent bond in the gaseous state. Lower bond enthalpy indicates a weaker, more easily broken bond.
The catalytic destruction of ozone (\(\text{O}_3\)) molecules in the stratosphere, primarily by chlorine free radicals derived from chlorofluorocarbons (CFCs).
Testing Reactivity: Silver Nitrate Hydrolysis
To measure the rate of hydrolysis of halogenoalkanes, they are reacted with water (acting as a nucleophile) in the presence of silver nitrate (\(\text{AgNO}_3\)) and an ethanol solvent (to allow the water and halogenoalkane to mix):
\[ \text{R-X(l)} + \text{H}_2\text{O(l)} \rightarrow \text{R-OH(aq)} + \text{H}^+\text{(aq)} + \text{X}^-\text{(aq)} \]
The halide ions (\(\text{X}^-\)) released react immediately with the silver ions (\(\text{Ag}^+\)) to form insoluble silver halide precipitates:
\[ \text{Ag}^+\text{(aq)} + \text{X}^-\text{(aq)} \rightarrow \text{AgX(s)} \]
By timing how long it takes for each precipitate to appear, we determine the relative rate of reaction:
| Halogenoalkane | Bond Enthalpy / kJ mol⁻¹ | Precipitate Identity | Precipitate Colour | Rate of Hydrolysis |
|---|---|---|---|---|
| Fluoroalkane | 467 | AgF (soluble) | No precipitate | Extremely slow / no reaction |
| Chloroalkane | 338 | AgCl | White precipitate | Slow |
| Bromoalkane | 276 | AgBr | Cream precipitate | Medium |
| Iodoalkane | 238 | AgI | Yellow precipitate | Fastest |
When asked to explain the trend in rate of hydrolysis, you must attribute it to decreasing bond enthalpy down the group. Do not write about electronegativity or polarity, as this would predict the opposite trend. You must explicitly state that the C-I bond is weaker/has a lower bond enthalpy than the C-Cl bond, meaning less energy is required to break it.
Ozone Depletion by CFCs
Chlorofluorocarbons (CFCs) were widely used in aerosols, air conditioners, and refrigerators because they are non-toxic, non-flammable, and extremely stable in the troposphere. However, their stability allows them to diffuse into the stratosphere, where they are exposed to high-energy ultraviolet (UV) radiation.
The Free Radical Mechanism
Under UV light, homolytic fission of the C-Cl bond occurs (because the C-Cl bond is weaker than the C-F bond), generating chlorine free radicals (\(\text{Cl}\bullet\)):
1. Initiation:
\[ \text{CF}_2\text{Cl}_2 \xrightarrow{uv} \text{CF}_2\text{Cl}\bullet + \text{Cl}\bullet \]
2. Propagation:
The chlorine free radical is highly reactive and attacks ozone (\(\text{O}_3\)), destroying the ozone molecule and forming a chlorine monoxide radical (\(\text{ClO}\bullet\)). The chlorine monoxide radical then reacts with another ozone molecule (or oxygen atom), regenerating the chlorine free radical:
\[ \text{Cl}\bullet + \text{O}_3 \rightarrow \text{ClO}\bullet + \text{O}_2 \]
\[ \text{ClO}\bullet + \text{O}_3 \rightarrow \text{Cl}\bullet + 2\text{O}_2 \]
3. Overall Equation:
Because the chlorine radical is used in the first propagation step and regenerated in the second, it acts as a catalyst. A single chlorine radical can destroy thousands of ozone molecules before terminating:
\[ 2\text{O}_3 \rightarrow 3\text{O}_2 \]
The Catalytic Ozone Cycle
The diagram below shows how the chlorine radical is continuously regenerated in the stratosphere, driving the decomposition of ozone:
Alternative Compounds
To protect the ozone layer, international agreements (like the Montreal Protocol) banned the use of CFCs. Chemists developed alternatives:
- Hydrochlorofluorocarbons (HCFCs): These contain hydrogen, chlorine, fluorine, and carbon. They are less stable than CFCs and break down in the lower atmosphere (troposphere) before reaching the ozone layer. However, they still possess some ozone depletion potential.
- Hydrofluorocarbons (HFCs): These contain only hydrogen, fluorine, and carbon. Because they do not contain chlorine, they do not release chlorine radicals in the stratosphere and have an ozone depletion potential of zero. Examples include \(\text{CF}_3\text{CH}_2\text{F}\).
Worked Examples
Answer:
Observation: 1-iodobutane will form a yellow precipitate of silver iodide much faster than 1-chlorobutane forms a white precipitate of silver chloride.
Explanation: Reactivity is determined by the carbon-halogen bond enthalpy. The C-I bond is longer and weaker (bond enthalpy of 238 kJ mol⁻¹) than the C-Cl bond (bond enthalpy of 338 kJ mol⁻¹). Consequently, less energy is required to break the C-I bond, allowing it to hydrolyse at a significantly faster rate.
Answer:
Initiation: UV light breaks the weakest C-Cl bond to form chlorine free radicals:
\[ \text{CFCl}_3 \xrightarrow{uv} \text{CFCl}_2\bullet + \text{Cl}\bullet \]
Propagation Step 1: The chlorine radical destroys ozone:
\[ \text{Cl}\bullet + \text{O}_3 \rightarrow \text{ClO}\bullet + \text{O}_2 \]
Propagation Step 2: The chlorine radical is regenerated:
\[ \text{ClO}\bullet + \text{O}_3 \rightarrow \text{Cl}\bullet + 2\text{O}_2 \]
The chlorine radical is regenerated at the end of propagation, proving it acts as a catalyst.
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