Question

Write the electron dot formula and draw the structural formula for each of the following molecules: (a) \(\mathrm{CH}_{4}\) (b) \(\mathrm{OF}_{2}\) (c) \(\mathrm{H}_{2} \mathrm{O}_{2}\) (d) \(\mathrm{NF}_{3}\)

Short answer

\(\mathrm{CH}_4\): Tetrahedral, \(\mathrm{OF}_2\): Bent, \(\mathrm{H}_2 \mathrm{O}_2\): Skewed, \(\mathrm{NF}_3\): Trigonal pyramidal.

Step-by-step solution

Understanding Electron Dot Formula

The electron dot formula, also known as the Lewis dot structure, represents the valence electrons of atoms within a molecule. Start by identifying the number of valence electrons for each atom in the molecule.

Electron Dot Formula for Methane, \(\mathrm{CH}_4\)

Carbon has 4 valence electrons, and each hydrogen has 1 valence electron. Carbon will share one electron with each hydrogen atom, forming four single bonds. Each hydrogen completes its valence shell with 2 electrons (Helium configuration), and carbon completes its with 8 electrons.

Drawing the Electron Dot Formula for \(\mathrm{CH}_4\)

Write the central carbon atom with one hydrogen atom connected at each side. Around carbon, mark one dot to represent each shared electron from hydrogen. This completes the electron dot formula for methane.

Electron Dot Formula for Oxygen Difluoride, \(\mathrm{OF}_2\)

Oxygen has 6 valence electrons, and each fluorine has 7 valence electrons. Oxygen will form single bonds with each fluorine, sharing one pair of electrons with each atom. Each fluorine also needs three lone pairs to complete its valence shell.

Drawing the Electron Dot Formula for \(\mathrm{OF}_2\)

Place the oxygen atom in the center with a single bond drawn to each fluorine. Position three lone pairs around each fluorine atom, and one pair around the oxygen to reflect its share of electrons.

Electron Dot Formula for Hydrogen Peroxide, \(\mathrm{H}_2 \mathrm{O}_2\)

Each oxygen atom has 6 valence electrons and each hydrogen has 1 valence electron. The structure will show each oxygen forming a single bond with another oxygen and a single bond with hydrogen. Oxygen atoms will possess two lone pairs each to complete the electron octets.

Drawing the Electron Dot Formula for \(\mathrm{H}_2 \mathrm{O}_2\)

Align the two oxygen atoms with a single bond between them, each attached to a hydrogen atom. Place two lone pairs on each oxygen atom, confirming the octet rule is satisfied.

Electron Dot Formula for Nitrogen Trifluoride, \(\mathrm{NF}_3\)

Nitrogen has 5 valence electrons, and each fluorine has 7 valence electrons. Nitrogen forms three single bonds, one with each fluorine, sharing a pair of electrons with each. Nitrogen also retains a lone pair, while each fluorine is surrounded with three lone pairs.

Drawing the Electron Dot Formula for \(\mathrm{NF}_3\)

Position nitrogen in the center, with one bond extending to each fluorine atom. Surround each fluorine with three lone pairs, and nitrogen with one to complete the structure.

Key concepts

Valence Electrons

Valence electrons are electrons found in the outermost shell of an atom and play a crucial role in chemical bonding. While core electrons reside in inner shells and are not typically involved in bonding, valence electrons participate directly in the formation of chemical bonds.
They determine how an atom will interact or bond with other atoms.

• These electrons help predict an element's reactivity and behavior in a molecule.
• The number of valence electrons is often equal to the atom's group number in the periodic table.
• Atoms strive to achieve a stable electronic configuration, usually similar to the nearest noble gas, by either gaining, losing, or sharing valence electrons.

Considering methane (\(\mathrm{CH}_4\)) as an example, carbon, with four valence electrons, shares these electrons with four hydrogen atoms, which each have one valence electron, to complete their valence shells. Understanding valence electrons sets the stage for comprehending molecular bonding and electron pairs.

Molecular Bonding

Molecular bonding, primarily covalent bonding, occurs when atoms share electrons to attain stability. This shared pair of electrons allows atoms to achieve the octet rule, a stable arrangement of eight electrons in their outer shell, much like noble gases.

• Covalent bonds can involve sharing one, two, or three pairs of electrons, known as single, double, or triple bonds, respectively.
• Each pair of shared electrons counts towards the octet of both bonded atoms.

In the molecule oxygen difluoride (\(\mathrm{OF}_2\)), the oxygen atom shares one pair of electrons (a single covalent bond) with each of the two fluorine atoms.
This sharing enables each fluorine atom to complete its octet, showcasing the role of covalent bonds in molecular stability. By understanding molecular bonding, one sees the basis for how molecules like these are held together.

Electron Pairs

Electron pairs in Lewis structures can be divided into two categories: bonding pairs and lone pairs.

• Bonding pairs represent shared electrons in a covalent bond between atoms.
• Lone pairs (or non-bonding pairs) are electrons that remain on a single atom and do not participate in bonding.

For instance, in hydrogen peroxide (\(\mathrm{H}_2\mathrm{O}_2\)), each oxygen atom has two lone pairs alongside the bonding pairs. These lone pairs are crucial for determining the molecule's geometry and reactivity.
By distinguishing between bonding pairs and lone pairs, we can fully comprehend the 3D structure and chemical properties of molecules like nitrogen trifluoride (\(\mathrm{NF}_3\)), where nitrogen has one lone pair that affects its resultant shape and angles. Recognizing these pairs aids in predicting molecule interactions and stability.