Question

Draw Lewis structures for the following: (a) \(\mathrm{H}_{2} \mathrm{~S}\) (b) \(\mathrm{CS}_{2}\) (c) \(\mathrm{NH}_{3}\) (d) \(\mathrm{NH}_{4} \mathrm{Cl}\)

Short answer

Lewis structures: (a) H2S: S central atom with two single bonds to H and two lone pairs. (b) CS2: C central atom with double bonds to two S atoms with two lone pairs on each S. (c) NH3: N central atom with three H atoms bonded and one lone pair. (d) NH4Cl: NH4+ ion with N center, four H atoms bonded, and Cl with an ionic bond to NH4+ and three lone pairs.

Step-by-step solution

Draw Lewis Structure for H2S

Start by arranging the atoms with sulfur (S) in the center and the two hydrogen (H) atoms around it. Sulfur has 6 valence electrons, and each hydrogen atom has 1 valence electron. Connect each hydrogen to sulfur with a single bond, which accounts for two electrons. Distribute the remaining electrons around sulfur to complete its octet.

Draw Lewis Structure for CS2

Place the carbon (C) atom in the center with the two sulfur (S) atoms on either side. Carbon has 4 valence electrons, while each sulfur atom has 6 valence electrons. Form double bonds between carbon and each sulfur atom to give carbon a full octet and each sulfur 8 valence electrons.

Draw Lewis Structure for NH3

Nitrogen (N) is the central atom with three hydrogen (H) atoms around it. Nitrogen has 5 valence electrons and each hydrogen atom has 1 valence electron. Form single bonds between nitrogen and each hydrogen atom, using up 6 valence electrons. Place the remaining pair of electrons (lone pair) on nitrogen.

Draw Lewis Structure for NH4Cl

For ammonium chloride (NH4Cl), start with the ammonium ion (NH4+). Nitrogen is the central atom surrounded by four hydrogen atoms. Each bond represents one electron from hydrogen and one from nitrogen. Since the ammonium ion has a +1 charge, remove one electron from the total count for nitrogen. Next, chloride (Cl) has 7 valence electrons, connect it to the ammonium ion with an ionic bond signifying the transfer of the electron, fulfilling Cl's octet.

Key concepts

Chemical Bonding

Understanding chemical bonding is fundamental to mastering chemistry. It is the process that holds atoms together within molecules or compounds. Bonds are formed through the interactions of atomic electrons, where atoms can either share or transfer valence electrons to achieve stable electron configurations. There are three primary types of chemical bonds: ionic, covalent, and metallic.

When elements with very different electronegativities (like metals and nonmetals) interact, they tend to form ionic bonds. One atom donates its valence electrons to another, resulting in positively and negatively charged ions that attract each other. In contrast, covalent bonding involves the sharing of electrons between nonmetals. It can vary from a single pair of electrons being shared (single bond) to multiple electron pairs (double or triple bonds). The metallic bond is distinctive to metals, where electrons flow freely between atoms. The type of bond dictates the properties of the substance, such as melting point, hardness, and conductivity.

Valence Electrons

Valence electrons play a pivotal role in chemical bonding. These are the electrons found in the outermost shell of an atom and are involved in the formation of chemical bonds. The number of valence electrons that an element possesses can often be determined by its group number on the periodic table. For instance, the elements in group 1 have 1 valence electron, while those in group 17 have 7.

Knowledge of valence electrons allows us to predict how atoms will bond. Atoms with few valence electrons, such as alkali metals, tend to lose electrons easily to achieve stability. Conversely, atoms with nearly full valence shells, such as halogens, look to gain electrons. Elements in group 18, known as noble gases, typically do not participate in bonding due to their full valence shells, which makes them unusually stable.

Octet Rule

The octet rule is a crucial concept in chemistry that revolves around the stability associated with having eight valence electrons. Most elements strive to achieve an electron configuration similar to that of the noble gases which is perceived as chemically stable. It is this desire for stability that drives the formation of bonds.

In the case of the exercise examples, the octet rule is applied to atoms like sulfur, carbon, nitrogen, and chlorine to obtain full valence shells. The only exception is hydrogen, which seeks to attain a duplet (two electrons), similar to the noble gas helium. The rule is a guideline more than a strict law; in some cases, molecules with an odd number of electrons or those that can have expanded octets (elements found in periods 3 and below of the periodic table) do not follow the octet rule strictly.

Electron Dot Structures

Electron dot structures, also known as Lewis dot structures, are a visual representation of the valence electrons in an atom or molecule. These diagrams help predict the bonding behavior of atoms, the arrangement of atoms within a molecule, and the presence of lone pairs of electrons. Lewis structures consist of element symbols surrounded by dots that represent valence electrons.

In the exercise examples, we visualize the shared pairs of electrons as lines (or dashes) and non-bonding electrons as dots around the atoms. For instance, in the Lewis structure of ammonia (\( NH_3 \)), the three hydrogen atoms are bonded to nitrogen with single covalent bonds while a pair of non-bonding electrons remain on the nitrogen. These structures not only illustrate the connectivity of atoms but also provide insight into the molecular geometry and reactivity of substances.