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) \(\mathrm{NH}_{2} \mathrm{Cl}\).

Short answer

The Lewis structures for the given molecules are as follows: (a) H2CO: H-C=O-H (b) H2O2: H-O-O-H (c) C2F6: F-C(1/2)-C-F (with all other fluorine atoms bonded to carbon atoms) (d) AsO3^3-: O=As-O(-)(1/2)-O= (with 2 other oxygen atoms with a single bond and a -1 charge) (e) H2SO3: O=S-O-H | O-H (f) NH2Cl: N-H | H-Cl

Step-by-step solution

Count the valence electrons

In the H2CO molecule, we have 2 hydrogen atoms, 1 carbon atom, and 1 oxygen atom. Hydrogen has 1 valence electron, carbon has 4, and oxygen has 6. Thus, we have 2(1) + 4 + 6 = 12 valence electrons.

Arrange the atoms

Based on the problem statement, both hydrogen atoms are bonded to the carbon atom, which is bonded to the oxygen atom.

Form bonds and distribute remaining electrons

We use the valence electrons to form bonds and lone pairs. First, we use 2 electrons to form a C-O bond, while the C-H bonds require 2 electrons each. This leaves us with 12 - 6 = 6 electrons. We place 4 of those as lone pairs on the oxygen, and then add the remaining 2 electrons as a double bond between carbon and oxygen. Lewis structure for H2CO: H-C=O-H (b) H2O2

Count the valence electrons

There are 2 hydrogen atoms and 2 oxygen atoms in the H2O2 molecule. Thus, we have 2(1) + 2(6) = 2 + 12 = 14 valence electrons.

Arrange the atoms

The oxygen atoms are in the center, each bonded to one hydrogen atom.

Form bonds and distribute remaining electrons

First form single bonds between O-H and O-O, using 4 electrons. This leaves us with 14 - 4 = 10 electrons. We then place 3 lone pairs on each oxygen atom, which leaves 10-6=4 electrons. We place the remaining electrons as a single bond between the two oxygen atoms. Lewis structure for H2O2: H-O-O-H (c) C2F6 (contains a C-C bond)

Count the valence electrons

In the C2F6 molecule, we have 2 carbon atoms and 6 fluorine atoms. This gives us 2(4) + 6(7) = 8 + 42 = 50 valence electrons.

Arrange the atoms

We have 2 carbon atoms bonded together in the center, and 6 fluorine atoms distributed around the carbon atoms.

Form bonds and distribute remaining electrons

We start by forming the C-C bond, which uses 2 electrons, leaving us with 50 - 2 = 48 electrons. We then form the C-F bonds using 2 electrons each. There are 6 C-F bonds, which use 6*2=12 electrons, leaving us with 48 - 12 = 36 electrons. We distribute the remaining electrons as lone pairs on each fluorine atom, using 3 lone pairs for each F (3*2*6=36 electrons). Lewis structure for C2F6: F-C(1/2)-C-F with all other fluorine atoms bonded to carbon atoms. (d) AsO3^3-

Count the valence electrons

In the AsO3^3- ion, we have 1 arsenic atom, 3 oxygen atoms, and a -3 charge. This gives us 5 + 3(6) + 3 = 5 + 18 +3= 26 valence electrons.

Arrange the atoms

We place the arsenic atom in the center, with the oxygen atoms surrounding it.

Form bonds and distribute remaining electrons

We start by forming single bonds between the As-O, which uses 3*2=6 electrons. This leaves us with 26 - 6 = 20 electrons. We then distribute the remaining electrons as lone pairs on each oxygen atom, using 2 lone pairs for each oxygen (2*2*3=12). Finally, we will place the remaining 20 - 12 = 8 electrons as two double bonds between the arsenic and two oxygen atoms, completing the octets. Lewis structure for AsO3^3-: O=As-O(-)(1/2)-O= with 2 other oxygen atoms with a single bond and a -1 charge. (e) H2SO3 (H is bonded to O)

Count the valence electrons

In the H2SO3 molecule, we have 2 hydrogen atoms, 1 sulfur atom, and 3 oxygen atoms. This gives us 2(1) + 6 + 3(6) = 2 + 6 + 18 = 26 valence electrons.

Arrange the atoms

Sulfur is surrounded by three oxygen atoms, which are bonded to hydrogen atoms.

Form bonds and distribute remaining electrons

First form single bonds between S-O and O-H, which uses 2 electrons each, and a total of 4*2=8 electrons. This leaves us with 26 - 8 = 18 electrons. We then place two lone pairs on two oxygen atoms and one lone pair on the third oxygen atom. The sulfur atom has two double bonds with the two oxygen atoms that have two lone pairs each. Lewis structure for H2SO3: O=S-O-H | O-H (f) NH2Cl

Count the valence electrons

In the NH2Cl molecule, we have 1 nitrogen atom, 2 hydrogen atoms, and 1 chlorine atom. This gives us 5 + 2(1) + 7 = 5 + 2 + 7 = 14 valence electrons.

Arrange the atoms

The nitrogen atom is in the center, with the two hydrogen atoms and chlorine atom surrounding it.

Form bonds and distribute remaining electrons

First form single bonds between N-H and N-Cl, which uses 2 electrons each and a total of 3*2=6 electrons. This leaves us with 14 - 6 = 8 electrons. We then place these remaining electrons as lone pairs - pairs on the nitrogen atom and 3 pairs on the chlorine atom. Lewis structure for NH2Cl: N-H | H-Cl

Key concepts

Valence Electrons

Valence electrons are the outermost electrons in an atom and are crucial in determining how atoms bond with each other. These electrons are located in the highest energy level of an atom and are responsible for the chemical properties of the element.
The number of valence electrons can usually be determined from an element's group number in the periodic table.

• For example, hydrogen belongs to Group 1, so it has 1 valence electron.
• Carbon is in Group 4 and has 4 valence electrons, while oxygen in Group 6 has 6 valence electrons.

Understanding the number of valence electrons is essential for predicting how atoms will interact and bond to form molecules. For instance, in Lewis structures, we represent valence electrons as dots around the symbol of an element. This visualization helps to identify how atoms will be shared or transferred in chemical bonds.

Chemical Bonds

Chemical bonds are the glue that holds atoms together in a molecule. They involve the interaction between valence electrons of the bonding atoms. There are several types of chemical bonds, but the most common ones include ionic bonds and covalent bonds.
Covalent bonds occur when atoms share pairs of electrons to attain a stable electronic configuration similar to those of noble gases. Ionic bonds are formed when one atom donates an electron to another atom, resulting in positively and negatively charged ions that attract each other.In the context of our problem:

• In \( \mathrm{H}_{2} \mathrm{CO} \), carbon forms covalent bonds with two hydrogen atoms and double bonds with oxygen using its valence electrons.
• Oxygen in \( \mathrm{H}_{2} \mathrm{O}_{2} \) shares electrons with hydrogen and itself to form stable single bonds.
• In \( \mathrm{C}_{2} \mathrm{F}_{6} \), each carbon atom shares electrons with fluorine atoms, forming covalent bonds despite the electronegative nature of fluorine.

Recognizing the type of bond and how electrons are shared is fundamental in drawing complete and accurate Lewis structures.

Covalent Bonding

Covalent bonding is a specific type of chemical bond where electrons are shared between atoms. This sharing allows each atom to achieve a noble gas electron configuration, thereby becoming more stable. Covalent bonds can be single, double, or triple bonds, depending on how many pairs of electrons are shared.

• Single bonds involve one pair of shared electrons, like the connections between hydrogen and oxygen in \( \mathrm{H}_{2} \mathrm{O}_{2} \).
• Double bonds involve two pairs of shared electrons, seen in the carbon-oxygen bond in \( \mathrm{H}_{2} \mathrm{CO} \).
• Triple bonds require three shared pairs of electrons and are less common in typical classroom problems.

Covalent bonds are crucial for the structural integrity of molecules and play a vital role in defining the properties of substances. For example, as in our problem examples, a covalent network can give rise to unique compounds with varying physical states and reactivities.