Question

Draw the Lewis structure for the chlorofluorocarbon CFC-11, CFCly. What chemical characteristics of this substance allow it to effectively deplete stratospheric ozone?

Short answer

The Lewis structure of CFC-11 (CCl₃F) has a central carbon atom surrounded by three chlorine atoms and one fluorine atom, with each connected to the carbon atom by a single bond. CFC-11 is effective at depleting stratospheric ozone due to its stability, which allows it to reach the stratosphere, and its ability to release chlorine atoms upon photodissociation. The chlorine atoms then act as catalysts for the breakdown of ozone into oxygen, leading to significant depletion of the stratospheric ozone layer.

Step-by-step solution

Determine the valence electrons of each atom

To draw the Lewis structure, we need to know the valence electrons of carbon (C), chlorine (Cl), and fluorine (F). The valence electrons are as follows: - Carbon (C) has 4 valence electrons (group 4) - Fluorine (F) has 7 valence electrons (group 7) - Chlorine (Cl) has 7 valence electrons (group 7) CFC-11 (CFCly) has the chemical formula CCl₃F, meaning it has one carbon atom, three chlorine atoms, and one fluorine atom.

Draw the Lewis structure of CFC-11 (CFCly)

We know that carbon atom has four valence electrons available for bonding. The central carbon atom will be surrounded by three chlorine atoms and one fluorine atom. Each chlorine atom has 7 valence electrons, and they need one more to achieve a stable octet. Similarly, the fluorine atom also requires one more electron to achieve the octet rule. Hence, the central carbon atom will form a single bond with each of the surrounding atoms (three Cl and one F). The Lewis structure for CFC-11 (CFCly) will look like this: Cl | Cl - C - F | Cl All atoms now have an octet of electrons surrounding them, and the Lewis structure of CFC-11 (CFCly) is complete.

Chemical characteristics that enable CFC-11 to deplete stratospheric ozone

CFC-11 has a few key characteristics that make it effective at depleting ozone: 1. Stability: The C-Cl and C-F bonds in CFC-11 are relatively strong, making the molecule stable and unreactive in the troposphere. This stability allows CFC-11 to reach the stratosphere, where it becomes reactive due to high-energy UV radiation. 2. Photodissociation: When CFC-11 reaches the stratosphere, the high-energy UV radiation can break one of the C-Cl bonds, releasing a Cl atom. The chemical reaction is: \[ CCl_{3}F \xrightarrow[\text{UV radiation}]{} Cl + CCl_{2}F\] 3. Ozone depletion: The Cl atom released in the stratosphere acts as a catalyst for the breakdown of ozone (O₃) into oxygen (O₂). One chlorine atom can destroy thousands of ozone molecules through a continuous reaction cycle: \(Cl + O_3 \rightarrow ClO + O_2\\ ClO + O \rightarrow Cl + O_2\) The regenerated Cl atom then proceeds to destroy more ozone molecules, leading to significant depletion of the stratospheric ozone layer. In conclusion, the Lewis structure of CFC-11 (CFCly) is determined by the valence electrons of carbon, chlorine, and fluorine atoms. The stability of CFC-11, coupled with its ability to release chlorine atoms upon photodissociation in the stratosphere, enables it to deplete the ozone layer effectively.

Key concepts

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom that are available for bonding with other atoms. Each element has a characteristic number of valence electrons that determines how it will bond with other atoms to form molecules. For instance, in our example of CFC-11 (CFCly), carbon (C) has 4 valence electrons, while both chlorine (Cl) and fluorine (F) have 7 each.

To create stable molecules, these atoms share their valence electrons through covalent bonds. The Lewis structure represents these valence electrons and the bonds formed between them. It's essential for understanding how a molecule like CFC-11 can interact with other substances, such as ozone in the stratosphere.

Octet Rule

The octet rule is a chemical rule of thumb that atoms tend to combine in such a way that they each have eight electrons in their valence shell, giving them the same electronic configuration as a noble gas. The rule applies to main-group elements, and it's instrumental in determining the shape and stability of molecules.

For the molecule CFC-11, each chlorine and fluorine atom requires one additional electron to complete its octet, while the carbon atom needs four to achieve the same stability. By sharing electrons through covalent bonds, each atom in the molecule attains a stable octet, minimizing the molecule's energy and making it chemically stable under normal conditions.

Stratospheric Ozone Depletion

Stratospheric ozone depletion refers to the thinning of the ozone layer in the Earth's stratosphere caused by various chemical compounds, among which chlorofluorocarbons (CFCs) like CFC-11 are significant contributors. The ozone layer is crucial as it protects the Earth from harmful ultraviolet (UV) radiation from the sun.

Ozone molecules absorb UV radiation and undergo a constant cycle of destruction and reformation. However, stable compounds like CFCs that reach the stratosphere can release chlorine atoms through photodissociation, a process where strong UV light causes the bonds within the molecule to break. These chlorine atoms act as catalysts, breaking down ozone at a much higher rate than it can be naturally replenished, leading to an 'ozone hole'.

Photodissociation

Photodissociation is the chemical process by which a chemical bond is broken or a molecule is cleaved into two or more smaller parts by the absorption of light. In the upper layers of the atmosphere, the intense UV radiation can cause photodissociation of stable molecules like CFC-11.

When CFC-11 absorbs UV radiation, a chlorine atom breaks away from the rest of the molecule. This reaction is crucial in understanding how CFCs lead to ozone depletion because the chlorine atoms set off a chain reaction that destroys thousands of ozone molecules. Thus, photodissociation of molecules like CFC-11 has long-term implications for the health of the stratospheric ozone layer and, consequently, for life on Earth due to increased exposure to UV radiation.