Question

(This problem is in the nature of review of elementary inorganic chemistry and may require reference to a general chemistry book.) Write Lewis structures for each of the following compounds. Use distinct, correctly placed dots for the electrons. Mark all atoms that are not neutral with charges of the proper sign. a. ammonia, \(\mathrm{NH}_{3}\) b. ammonium bromide, \(\mathrm{NH}_{4} \mathrm{Br}\) c. hydrogen cyanide, \(\mathrm{HCN}\) d. ozone \(\left(\angle \mathrm{O}-\mathrm{O}-\mathrm{O}=120^{\circ}\right)\) e. carbon dioxide, \(\mathrm{CO}_{2}\) f. hydrogen peroxide, \(\mathrm{HOOH}\) g. hydroxylamine, \(\mathrm{HONH}_{2}\) h. nitric acid, \(\mathrm{HNO}_{3}\) i. hydrogen sulfide, \(\mathrm{H}_{2} \mathrm{~S}\) j. boron trifluoride, \(\mathrm{BF}_{3}\)

Short answer

Lewis structures drawn for each compound, showing bonding and any charges.

Step-by-step solution

Ammonia (NH3) Structure

First, count the total number of valence electrons: Nitrogen has 5 and each hydrogen has 1, totaling 8 electrons. Place nitrogen in the center, draw three single bonds to three hydrogen atoms, each sharing a pair of electrons, and place a lone pair on nitrogen.

Ammonium Bromide (NH4Br) Structure

Ammonium \(NH_4^+\) is formed by adding a proton to ammonia \(NH_3\). Bromide is \(Br^-\). Total valence electrons are: 5 (N) + 4 (H) - 1 (charge). Draw NH4+ like NH3 but with four N-H bonds and no lone pairs on nitrogen, add Br- nearby. Mark Br with a negative charge and NH4 with a positive charge.

Hydrogen Cyanide (HCN) Structure

Total electrons are 4 (C) + 1 (H) + 5 (N) = 10. With carbon in the center, draw a single bond to hydrogen and a triple bond to nitrogen, which satisfies the octet rule for both carbon and nitrogen.

Ozone (O3) Structure

Ozone has 18 valence electrons (3 x 6 for each oxygen). Draw the central O with double bonds to each side O, using resonance structures, where one oxygen has a formal charge of +1 and another has -1, represented by shifting the double bond position.

Carbon Dioxide (CO2) Structure

Carbon dioxide has 4 (C) + 2 x 6 (O) = 16 valence electrons. Draw carbon in the center with double bonds to each oxygen, ensuring that each atom satisfies the octet rule.

Hydrogen Peroxide (HOOH) Structure

Total electrons are 2 (H) + 2 x 6 (O) = 14. Connect two O's with a single bond, each with one H attached by a single bond. Complete with lone pairs on oxygens to satisfy the octet.

Hydroxylamine (HONH2) Structure

Count the total valence electrons: 1 (H) + 6 (O) + 5 (N) + 2 x 1 (H) = 14. Place oxygen bonded to nitrogen, which is then bonded to two hydrogens. Distribute electrons to complete octets, with lone pairs on oxygen.

Nitric Acid (HNO3) Structure

With total electrons 1 (H) + 5 (N) + 3 x 6 (O) = 24, place nitrogen in the center. Attach one O with a single and another with a double bond. The single bonded O holds the H atom. The third O has a lone pair with a positive charge on N and negative on the single-bonded O.

Hydrogen Sulfide (H2S) Structure

H2S has a total of 2 (H) + 6 (S) = 8 electrons. Place sulfur in the center, bonded to two hydrogens, with lone pairs on sulfur to satisfy the octet rule.

Boron Trifluoride (BF3) Structure

Boron has 3 electrons and each F contributes 7, totaling 3 + 3x7 = 24 electrons. Place Boron at center with single bonds to each F, which achieves 6 electrons around B, below octet, but stable for BF3.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in chemical bonding. They determine how an atom interacts with others to form molecules.

• Each element in the periodic table has a characteristic number of valence electrons. For example, nitrogen has five valence electrons, while hydrogen has one.
• Valence electrons are used in forming bonds, as they are the electrons involved in either sharing or transferring between atoms.
• In Lewis structures, valence electrons are depicted as dots placed around element symbols.

Understanding the count and arrangement of valence electrons is the first step in drawing accurate Lewis structures as demonstrated with compounds such as ammonia (\(\mathrm{NH}_3\)) and hydrogen sulfide (\(\mathrm{H}_2\mathrm{S}\)). Counting valence electrons helps to ensure each atom's contributions are accounted for in the molecule.

Octet Rule

The octet rule is a chemical rule of thumb that suggests atoms are most stable when they have eight electrons in their valence shell, similar to a noble gas configuration.

• This rule helps predict how atoms combine and form stable molecules or ions.
• Atoms will share, donate, or receive electrons to achieve a full octet, influencing the structure of molecules like carbon dioxide (\(\mathrm{CO}_2\)) and nitric acid (\(\mathrm{HNO}_3\)).
• Some elements follow the rule strictly, while others, like boron in boron trifluoride (\(\mathrm{BF}_3\)), can have slightly different arrangements.

The rule serves as a guideline, although exceptions exist. For instance, in molecules like ozone (\(\mathrm{O}_3\)), formal charges balance the octet, while in boron trifluoride, boron is content with less than a complete octet.

Chemical Bonds

Chemical bonds are the connections between atoms that result from the sharing or transferring of valence electrons. These bonds determine the structural integrity and function of molecules.

• Covalent bonds are formed when atoms share electrons, as seen in \(\mathrm{NH}_3\) and \(\mathrm{HCN}\), where electrons are shared between hydrogen, nitrogen, and carbon.
• Ionic bonds involve transferring electrons, like in ammonium bromide (\(\mathrm{NH}_4\mathrm{Br}\)), where there is a positive ammonium ion and a negative bromide ion.
• Double and triple bonds occur when atoms share two or three pairs of electrons, enhancing the stability and reducing the energy state of the molecule, as observed in \(\mathrm{CO}_2\) and \(\mathrm{HCN}\).

Understanding the type of bond and its formation is essential to predicting and rationalizing molecular properties, shapes, and reactivity.

Formal Charge

Formal charge is a tool used to determine the most feasible Lewis structure for a molecule by examining the distribution of electrons.

• It is calculated as: \[\text{{Formal charge}} = \text{{valence electrons}} - \text{{non-bonding electrons}} - \frac{{1}}{{2}} \text{{bonding electrons}}\]
• In ozone (\(\mathrm{O}_3\)), resonance structures can be drawn by allocating formal charges to optimize stability with +1 and -1 charges on different oxygen atoms.
• A formal charge closer to zero often contributes to a more stable structure, as in the arrangement of atoms in nitric acid (\(\mathrm{HNO}_3\)).

The goal in checking formal charges is to achieve the most balanced structure, adhering to the octet rule when possible and considering the overall charge of ions or neutral molecules.