Question

Write Lewis structures for the following: (a) \(\mathrm{H}_{2} \mathrm{CO}\) (both \(\mathrm{H}\) atoms are bonded to \(\mathrm{C}\) ), (b) \(\mathrm{H}_{2} \mathrm{O}_{2}\), (c) \(\mathrm{C}_{2} \mathrm{~F}_{6}\) (contains a \(\mathrm{C}-\mathrm{C}\) bond \(),\) (d) \(\mathrm{AsO}_{3}{ }^{3-}\), (e) \(\mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}\) is bonded to \(\mathrm{O})\), (f) \(\mathrm{C}_{2} \mathrm{H}_{2}\).

Short answer

The Lewis structures for the given molecules/ions are as follows: a) H2CO: O :: C - H : H b) H2O2: O - H :: : O - H c) C2F6: :F - C - C - F: : : : F F F d) AsO3 ^{3-}: [::O - As = O]^{3-} : O e) H2SO3 (H is bonded to O): O - S - O :: : :: H H f) C2H2: H - C ≡ C - H

Step-by-step solution

a) \(\mathrm{H}_{2} \mathrm{CO}\)

1. Count the total number of valence electrons: C has 4 valence electrons, O has 6, and two hydrogens have 1 each. Altogether, we have \(4 + 6 + 1 + 1 = 12\) electrons. 2. Connect the atoms with a single bond: As mentioned in the problem, both H atoms are bonded to C, so we connect them accordingly. 3. Distribute the remaining electrons: There are 8 remaining electrons to satisfy the octet rule. The next step is to place the remaining electrons, giving priority to the more electronegative element first (in this case, oxygen). The Lewis structure for \(\mathrm{H}_{2}\mathrm{CO}\) is: O :: C - H : H

b) \(\mathrm{H}_{2} \mathrm{O}_{2}\)

1. Count the total number of valence electrons: 2 oxygens have 6 electrons each, and two hydrogens have 1 each. We have a total of \(6 + 6 + 1 + 1 = 14\) electrons. 2. Connect the atoms with a single bond: In hydrogen peroxide, the hydrogen atoms are bonded to different oxygen atoms. 3. Distribute the remaining electrons: There are 10 remaining electrons to satisfy the octet rule. Distribute the remaining electrons to fulfill the octet rule for each of the oxygen atoms. The Lewis structure for \(\mathrm{H}_{2} \mathrm{O}_{2}\) is: O - H :: : O - H

c) \(\mathrm{C}_{2} \mathrm{F}_{6}\)

1. Count the total number of valence electrons: 2 carbons have 4 valence electrons each, and 6 fluorine atoms have 7 each. We have a total of \(4 + 4 + 42 = 50\) electrons. 2. Connect the central atoms with a single bond: It is given that there is a C-C bond. 3. Connect peripheral atoms with single bonds: Connect 3 F atoms to each of the C atoms. 4. Distribute the remaining electrons: There are 32 remaining electrons to satisfy the octet rule. Assign the remaining electrons to the F atoms. The Lewis structure for \(\mathrm{C}_{2}\mathrm{F}_{6}\) is: :F - C - C - F: : : : F F F

d) \(\mathrm{AsO}_{3}{ }^{3-}\)

1. Count the total number of valence electrons: As has 5 valence electrons, 3 oxygens have 6 electrons each, and 3 extra electrons due to the 3- charge. We have a total of \(5 + 18 + 3 = 26\) electrons. 2. Connect the central atom with peripheral atoms: As is the central atom, with 3 bonds to each oxygen, forming a trigonal planar geometry. 3. Distribute the remaining electrons: There are 14 remaining electrons to satisfy the octet rule. Assign the remaining electrons to the oxygen atoms. The Lewis structure for \(\mathrm{AsO}_{3}{ }^{3-}\) is: [::O - As = O]^{3-} : O

e) \(\mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}\) is bonded to \(\mathrm{O})\)

1. Count the total number of valence electrons: S has 6 valence electrons, 3 oxygens have 6 electrons each, and 2 hydrogens have 1 electron each. We have a total of \(6 + 18 + 2 = 26\) electrons. 2. Connect the central atom with peripheral atoms: S is the central atom, with bond to all three oxygen atoms. 3. Connect hydrogen to oxygen atoms: It is given that both H atoms are bonded to O atoms. 4. Distribute the remaining electrons: There are 14 remaining electrons to satisfy the octet rule. Assign these electrons to the oxygen atoms. The Lewis structure for \(\mathrm{H}_{2} \mathrm{SO}_{3}\) is: O - S - O :: : :: H H

f) \(\mathrm{C}_{2} \mathrm{H}_{2}\)

1. Count the total number of valence electrons: 2 carbons have 4 valence electrons each, and 2 hydrogens have 1 electron each. We have a total of \(4 + 4 + 1 + 1 = 10\) electrons. 2. Connect the central atoms with a bond: It is given that there is a C-C bond. 3. Connect peripheral atoms with a single bond: Each C atom has one hydrogen atom. 4. Distribute the remaining electrons: There are 6 remaining electrons to satisfy the octet rule. Assign these electrons to the carbon atoms. The Lewis structure for \(\mathrm{C}_{2} \mathrm{H}_{2}\) is: H - C ≡ C - H

Key concepts

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. They are crucial because they determine how an atom will interact or bond with other atoms. When drawing Lewis structures, the amount of valence electrons guides us on how many bonds an atom can form. For example:

• Carbon (C) has 4 valence electrons, allowing it to form up to 4 bonds.
• Oxygen (O) has 6 valence electrons, typically forming 2 bonds to achieve a stable structure.
• Hydrogen (H), with only 1 valence electron, can form a single bond.

By knowing the number of valence electrons, we can accurately predict and draw the arrangement of atoms in molecules, leading to the correct Lewis structure.

Octet Rule

The Octet Rule is a key principle in chemistry stating that atoms are most stable when they have eight electrons in their valence shell. This rule helps explain the chemical bonding of many elements. For example, most metals lose electrons to achieve an octet and become positively charged ions, while nonmetals gain electrons to complete their octet.
When applying the octet rule in Lewis structures:

• Each atom, typically, strives to have 8 electrons around it through bonds or lone pairs.
• There are exceptions, such as hydrogen, which only needs 2 electrons for a full shell.
• To fulfill the octet rule, atoms share electrons, leading to the formation of covalent bonds.

Understanding the octet rule allows us to better predict the stability and reactivity of molecules.

Chemical Bonding

Chemical bonding refers to the attraction between atoms that enables the formation of chemical substances. Bonds arise due to the sharing or transfer of valence electrons, aiming to achieve a full outer shell according to the octet rule. There are several types of bonds, including:

• Covalent Bonds: Form when two atoms share one or more pairs of valence electrons. For instance, in the molecule \(\mathrm{H}_2\mathrm{CO}\), carbon shares electrons with hydrogen and oxygen.
• Ionic Bonds: Occur when electrons are transferred between atoms, forming charged ions which attract each other. This is common in salts.
• Double and Triple Bonds: Double bonds involve the sharing of two pairs of electrons, like in \(\mathrm{O}_2\), while triple bonds involve three pairs, as found in \(\mathrm{C}_2\mathrm{H}_2\).

The type of chemical bond affects the molecule's properties, such as its strength and melting point.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. It is largely determined by the Lewis structure and the repulsion between electron pairs, following the VSEPR (Valence Shell Electron Pair Repulsion) theory.
Different geometries include:

• Linear: A straight-line shape with a 180-degree bond angle, as seen in molecules like \(\mathrm{C}_2\mathrm{H}_2\).
• Trigonal Planar: A flat, triangular shape with 120-degree angles, typical for molecules with three bonding pairs around a central atom, like \(\mathrm{AsO}_3^{3-}\).
• Tetrahedral: A three-dimensional shape where four bonds are equally spaced around a central atom, seen in \(\mathrm{C}_2\mathrm{F}_6\).

Understanding molecular geometry helps predict the physical and chemical properties of substances, including their reactivity, polarity, and interactions with other molecules.