6. Draw The Lewis Structure For C2h4.

Ethylene, also known as ethene, with the chemical formula C2H4, is a fundamental building block in the petrochemical industry and a vital plant hormone. Understanding its molecular structure, particularly through the Lewis structure, is crucial for comprehending its chemical properties and reactivity. Let's delve into a step-by-step guide on drawing the Lewis structure for C2H4, along with an explanation of the underlying principles and related concepts.

Drawing the Lewis Structure for C2H4: A Step-by-Step Guide

Constructing a Lewis structure involves representing the valence electrons of atoms within a molecule to illustrate chemical bonds. Here’s how to do it for ethylene:

Step 1: Determine the Total Number of Valence Electrons

Carbon (C) is in group 14 (or IVA) of the periodic table, so each carbon atom has 4 valence electrons.
Hydrogen (H) is in group 1, so each hydrogen atom has 1 valence electron.

In C2H4, there are two carbon atoms and four hydrogen atoms. Therefore, the total number of valence electrons is:

(2 carbons * 4 valence electrons/carbon) + (4 hydrogens * 1 valence electron/hydrogen) = 8 + 4 = 12 valence electrons

Step 2: Draw the Skeletal Structure

The skeletal structure shows how the atoms are connected. Carbon atoms tend to bond with each other, so in C2H4, the two carbon atoms are bonded together, with each carbon atom also bonded to two hydrogen atoms. The initial structure looks like this:

H H | | C--C | | H H

Step 3: Distribute the Valence Electrons as Single Bonds

Each single bond represents two shared electrons. In the skeletal structure above, there are five single bonds (one C-C bond and four C-H bonds). Distributing the valence electrons as single bonds:

One C-C bond uses 2 valence electrons.
Four C-H bonds use 4 * 2 = 8 valence electrons.

Total electrons used so far: 2 + 8 = 10 valence electrons.

Step 4: Check Remaining Valence Electrons

We started with 12 valence electrons and have used 10, leaving 12 - 10 = 2 valence electrons.

Step 5: Place Remaining Electrons as Lone Pairs

If there were remaining electrons, they would typically be placed as lone pairs on the atoms, starting with the most electronegative atoms (excluding hydrogen). However, in this case, we only have 2 electrons left, and neither carbon atom has a complete octet.

Step 6: Form Multiple Bonds to Satisfy the Octet Rule

Each carbon atom currently has only 6 electrons around it (2 from the C-C bond and 2 from each of the two C-H bonds). To satisfy the octet rule, we need to form a double bond between the two carbon atoms. This means sharing two pairs of electrons between the carbon atoms instead of just one.

Replace the single bond between the carbon atoms with a double bond:

H H | | C=C | | H H

Now, each carbon atom has 8 electrons around it (4 from the C=C double bond and 2 from each of the two C-H bonds). All 12 valence electrons have been used, and each carbon atom has a complete octet.

Step 7: Final Lewis Structure

The final Lewis structure for C2H4 (ethylene) is:

H H | | C=C | | H H

This representation accurately depicts the arrangement of atoms and the distribution of valence electrons in the ethylene molecule.

Understanding the Significance of the Lewis Structure

The Lewis structure provides valuable insights into the properties and behavior of ethylene:

Double Bond: The presence of a double bond between the two carbon atoms is crucial. This double bond consists of one sigma (σ) bond and one pi (π) bond. The pi bond is weaker than the sigma bond, making the double bond more reactive than a single bond.
Planar Geometry: The double bond restricts rotation around the C=C axis, causing the ethylene molecule to be planar. All six atoms (two carbons and four hydrogens) lie in the same plane.
Reactivity: The double bond in ethylene is a site of high electron density, making it susceptible to electrophilic attack. This explains why ethylene undergoes addition reactions readily, such as hydrogenation, halogenation, and hydration.

The Science Behind the Lewis Structure

The Lewis structure is based on the octet rule, which states that atoms tend to gain, lose, or share electrons in order to achieve a full outer shell of 8 electrons (except for hydrogen, which seeks 2 electrons). The formation of chemical bonds allows atoms to achieve this stable electron configuration.

In ethylene, carbon atoms share electrons to form covalent bonds. Covalent bonds occur when atoms share electrons rather than transferring them (as in ionic bonds). The sharing of electrons allows both carbon atoms to achieve a stable octet configuration, resulting in a stable molecule.

Key Properties and Reactivity of Ethylene (C2H4)

Ethylene is one of the most important organic chemicals produced worldwide. Its properties and reactivity make it a versatile building block for a wide range of industrial processes:

Colorless Gas: Ethylene is a colorless, flammable gas at room temperature and pressure.
Slightly Sweet Odor: It has a faint, sweet odor.
Low Solubility in Water: Ethylene is not very soluble in water but is soluble in organic solvents.
Addition Reactions: Ethylene readily undergoes addition reactions, where atoms or groups of atoms add across the double bond. Common addition reactions include:
- Hydrogenation: Addition of hydrogen (H2) to form ethane (C2H6).
- Halogenation: Addition of halogens (e.g., Cl2, Br2) to form dihaloalkanes (e.g., 1,2-dichloroethane).
- Hydration: Addition of water (H2O) to form ethanol (C2H5OH).
- Polymerization: Polymerization of ethylene to form polyethylene, a widely used plastic.
Oxidation Reactions: Ethylene can be oxidized to form ethylene oxide, which is then used to produce ethylene glycol (antifreeze) and other chemicals.
Flammability: Ethylene is highly flammable and can form explosive mixtures with air.

Industrial Uses of Ethylene

Ethylene is a crucial feedstock in the petrochemical industry, serving as a precursor for numerous products:

Polyethylene Production: The largest use of ethylene is in the production of polyethylene (PE), a versatile plastic used in packaging, films, containers, and many other applications. There are several types of polyethylene, including low-density polyethylene (LDPE), high-density polyethylene (HDPE), and linear low-density polyethylene (LLDPE).
Ethylene Oxide Production: Ethylene oxide is an important intermediate used in the production of ethylene glycol, surfactants, and other chemicals. Ethylene glycol is primarily used as antifreeze in automotive cooling systems and as a precursor for polyester fibers and resins.
Ethylene Dichloride Production: Ethylene dichloride (EDC) is used to produce vinyl chloride, which is then polymerized to form polyvinyl chloride (PVC), a widely used plastic in construction, pipes, and other applications.
Ethylbenzene Production: Ethylbenzene is used to produce styrene, which is then polymerized to form polystyrene, a versatile plastic used in packaging, insulation, and other products.
Ethanol Production: Ethylene can be hydrated to produce ethanol, which is used as a solvent, fuel additive, and chemical intermediate.
Other Uses: Ethylene is also used in the production of various other chemicals, including acetaldehyde, acetic acid, and alpha-olefins.

Biological Role of Ethylene

In addition to its industrial importance, ethylene also plays a significant role in plant biology:

Plant Hormone: Ethylene is a plant hormone that regulates various aspects of plant growth, development, and senescence.
Fruit Ripening: Ethylene promotes fruit ripening, causing changes in color, texture, and flavor. It is often used commercially to ripen fruits such as bananas, tomatoes, and apples.
Flower Senescence: Ethylene promotes flower senescence, causing wilting and abscission (shedding) of petals.
Leaf Abscission: Ethylene promotes leaf abscission, causing leaves to fall off trees in the autumn.
Stress Response: Ethylene is produced by plants in response to various stresses, such as wounding, disease, and flooding. It can help plants cope with these stresses by triggering defense mechanisms.

Common Mistakes When Drawing Lewis Structures

Drawing Lewis structures can sometimes be challenging, and certain common mistakes are often made:

Incorrectly Counting Valence Electrons: Always double-check the number of valence electrons for each atom. Refer to the periodic table to ensure accuracy.
Violating the Octet Rule: Ensure that each atom (except hydrogen) has 8 electrons around it. If an atom has fewer than 8 electrons, consider forming multiple bonds (double or triple bonds) to satisfy the octet rule.
Incorrectly Placing Lone Pairs: Place lone pairs on atoms to satisfy the octet rule, starting with the most electronegative atoms (excluding hydrogen).
Forgetting to Minimize Formal Charges: While not always necessary for basic Lewis structures, minimizing formal charges can help determine the most stable arrangement of atoms and electrons. Formal charge is calculated as:

Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (½ Bonding Electrons)
Drawing Incorrect Skeletal Structures: Make sure the skeletal structure accurately represents how the atoms are connected. Carbon atoms typically form chains or rings, and hydrogen atoms are usually terminal (bonded to only one other atom).

Advanced Concepts Related to Ethylene

For a deeper understanding of ethylene, consider exploring these advanced concepts:

Molecular Orbital Theory: Molecular orbital (MO) theory provides a more sophisticated description of chemical bonding than Lewis structures. MO theory explains how atomic orbitals combine to form bonding and antibonding molecular orbitals, which determine the electronic structure and properties of molecules.
Hybridization: Hybridization is the mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies. In ethylene, each carbon atom is sp2 hybridized, meaning that one s orbital and two p orbitals mix to form three sp2 hybrid orbitals. The remaining p orbital forms the pi bond of the double bond.
Resonance Structures: While ethylene itself does not exhibit resonance, the concept of resonance is important for understanding other molecules with delocalized electrons. Resonance occurs when a molecule can be represented by two or more Lewis structures that differ only in the arrangement of electrons.
Spectroscopy: Spectroscopic techniques, such as infrared (IR) spectroscopy and nuclear magnetic resonance (NMR) spectroscopy, can be used to study the structure and properties of ethylene and other molecules. IR spectroscopy measures the vibrations of molecules, while NMR spectroscopy measures the magnetic properties of atomic nuclei.

Conclusion

Drawing the Lewis structure for ethylene (C2H4) is a fundamental exercise in understanding chemical bonding and molecular structure. By following the step-by-step guide outlined above, you can accurately represent the arrangement of atoms and the distribution of valence electrons in this important molecule. The Lewis structure reveals the presence of a double bond between the carbon atoms, which is responsible for ethylene's planar geometry and high reactivity. Ethylene's properties make it a crucial feedstock in the petrochemical industry and a vital plant hormone. Understanding ethylene's structure and properties is essential for anyone studying chemistry, biology, or related fields.