Question

Illustrate the formation of single, double, and triple covalent bonds using Lewis structure

Short answer

To illustrate the formation of single, double, and triple covalent bonds using Lewis structures, we can use the example molecules H2, O2, and N2 respectively. For H2, each hydrogen atom shares its single valence electron, forming a single covalent bond: H:\( \bullet \bullet \):H. For O2, each oxygen atom shares two of its valence electrons, forming a double covalent bond: O::\( \bullet \bullet \)::\( \bullet \bullet \)::O. For N2, each nitrogen atom shares three of its valence electrons, forming a triple covalent bond: N:::\( \bullet \bullet \bullet \):::\( \bullet \bullet \bullet \):::N.

Step-by-step solution

Choose an example molecule for each bond type

To explain the formation of single, double, and triple covalent bonds, we need one example molecule for each type of bond. - Single covalent bond: H2 - Double covalent bond: O2 - Triple covalent bond: N2

Identify the number of valence electrons for each atom in the molecules

Since Lewis structures are about valence electrons, we need to know the number of valence electrons for each atom to draw the structures. - Hydrogen (H) has 1 valence electron. - Oxygen (O) has 6 valence electrons. - Nitrogen (N) has 5 valence electrons.

Draw the Lewis structure for H2

To form a single covalent bond, each hydrogen atom will share its single valence electron with the other hydrogen atom. 1. Write the symbols of the two atoms: H - H 2. Place a pair of dots (representing the shared valence electrons) between the atoms: H:\( \bullet \bullet \):H Thus, we have illustrated a single covalent bond for H2 using the Lewis structure.

Draw the Lewis structure for O2

To form a double covalent bond, each oxygen atom will share two of its six valence electrons with the other oxygen atom. 1. Write the symbols of the two atoms: O = O 2. Place two pairs of dots (representing the shared valence electrons) between the atoms: O::\( \bullet \bullet \)::\( \bullet \bullet \)::O 3. Distribute the remaining valence electrons around each atom to complete their octets (except for the shared pairs): O::\( \bullet \bullet \)::\( \bullet \bullet \)::O Thus, we have illustrated a double covalent bond for O2 using the Lewis structure.

Draw the Lewis structure for N2

To form a triple covalent bond, each nitrogen atom will share three of its five valence electrons with the other nitrogen atom. 1. Write the symbols of the two atoms: N ≡ N 2. Place three pairs of dots (representing the shared valence electrons) between the atoms: N:::\( \bullet \bullet \bullet \):::\( \bullet \bullet \bullet \):::N 3. Distribute the remaining valence electrons around each atom to complete their octets (except for the shared pairs): N:::\( \bullet \bullet \bullet \):::\( \bullet \bullet \bullet \):::N Thus, we have illustrated a triple covalent bond for N2 using the Lewis structure.

Key concepts

Lewis Structure

The Lewis structure is a visual representation of how atoms bond and form molecules. It’s a way to depict the arrangement of valence electrons, which are responsible for chemical bonds, around atoms. In a Lewis structure, each dot represents a valence electron, and a line or pair of dots between atoms represents a shared pair of electrons—a covalent bond.

When drawing a Lewis structure, you start by determining the valence electrons for each atom, then arrange those electrons to satisfy the octet rule for each atom, meaning that each atom should have eight electrons in its outer shell (with the exception of hydrogen, which aims for two). The structure visually demonstrates how atoms share electrons to reach a stable electronic configuration. If properly drawn, it helps predict the physical and chemical properties of the molecule, such as shape and reactivity.

For example, to draw the Lewis structure of a water molecule (H₂O), we start with the oxygen atom in the center and place the hydrogen atoms on either side. We then draw dots to represent oxygen's six valence electrons and lines or pairs of dots for hydrogen's single electron, so they share electrons to form covalent bonds.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in chemical bonding because they can be lost, gained, or shared with other atoms. The number of valence electrons an atom has influences how it interacts with other atoms and thus determines the types of chemical bonds it can form.

Atoms with a full outer shell, like noble gases, are typically inert because they don’t need to lose, gain, or share electrons. However, other elements, such as sodium (with one valence electron) and chlorine (with seven valence electrons), readily form bonds to achieve a stable octet.

• Sodium tends to lose its one valence electron to become positively charged.
• Chlorine tends to gain an electron to fill its valence shell, becoming negatively charged.
• When these two elements react, they form the ionic compound sodium chloride (NaCl).

Understanding valence electrons is key to determining how atoms will bond and is the first step in drawing a Lewis structure for a molecule.

Chemical Bonding

Chemical bonding refers to the force that holds atoms together in a compound. There are several types of chemical bonds, the most common being ionic, covalent, and metallic bonds.

• Ionic bonds occur when one atom donates an electron to another atom, resulting in positively and negatively charged ions that attract each other.
• Covalent bonds happen when two atoms share one or more pairs of electrons to achieve full outer shells and are commonly found in organic compounds.
• Metallic bonds involve a 'sea of electrons' that flow freely around a lattice of metal cations, giving metals their characteristic properties, such as conductivity and malleability.

Covalent bonding, the focus of our exercise, involves sharing of electrons between atoms. Based on the number of shared electron pairs, covalent bonds can be single (one shared pair), double (two shared pairs), or triple (three shared pairs). These shared electrons allow each atom in the molecule to reach a stable electronic configuration, akin to that of noble gases.