Question

On the basis of the electronegativity values given in Figure \(11.3,\) indicate whether each of the following bonds would be expected to be ionic, covalent, or polar covalent. a. \(\mathrm{H}-\mathrm{O}\) b. \(\mathrm{O}-\mathrm{O}\) c. \(\mathrm{H}-\mathrm{H}\) d. \(\mathrm{H}-\mathrm{Cl}\)

Short answer

The bond types for the given bonds are as follows: a. H-O: Polar covalent bond b. O-O: Covalent bond c. H-H: Covalent bond d. H-Cl: Polar covalent bond

Step-by-step solution

Understand electronegativity and its relationship with bond types

Electronegativity is a measure of an atom's ability to attract electrons towards itself when forming a chemical bond. When the difference in electronegativity between the two bonded atoms is large, there is a greater likelihood that the bond will be ionic. When the difference is small, the bond will likely be covalent or polar covalent. As a general rule, the bond types can be classified based on electronegativity difference: - Ionic bond: Electronegativity difference \(\geq\) 1.7 - Polar covalent bond: Electronegativity difference between 0.4 and 1.7 - Covalent bond: Electronegativity difference \(<\) 0.4

Determine electronegativity values of given atoms

Identify the electronegativity values of the atoms given in the problem using Figure 11.3. For example, electronegativity values for Hydrogen (H), Oxygen (O), and Chlorine (Cl) are: - H: 2.1 - O: 3.5 - Cl: 3.0

Calculate the electronegativity differences

For each given bond, subtract the electronegativity values of the participating atoms and find the absolute difference. This will help determine the bond type. a. H-O: \(|3.5-2.1|=1.4\) b. O-O: \(|3.5-3.5|=0\) c. H-H: \(|2.1-2.1|=0\) d. H-Cl: \(|3.0-2.1|=0.9\)

Determine the bond type based on electronegativity differences

Now simply classify the bond types based on the electronegativity differences calculated in Step 3, using the general rule stated in Step 1: a. H-O: Electronegativity difference = 1.4, which is between 0.4 and 1.7, so this bond is polar covalent. b. O-O: Electronegativity difference = 0, which is less than 0.4, so this bond is covalent. c. H-H: Electronegativity difference = 0, which is less than 0.4, so this bond is covalent. d. H-Cl: Electronegativity difference = 0.9, which is between 0.4 and 1.7, so this bond is polar covalent.

Key concepts

Chemical Bonding Basics

Understanding chemical bonding is fundamental to the study of chemistry. Atoms bond together to form molecules, and the way they bond is crucial in determining the properties of the substances they create. The electrons involved in bonding belong to the outermost shell of an atom, known as the valence electrons. Bonds form to achieve a stable configuration, often resembling the nearest noble gas in terms of electron configuration.

There are three main types of bonds: ionic, covalent, and polar covalent. The type of bond that forms between two atoms depends on the relative electronegativity of the atoms. Electronegativity is a measure of how strongly an atom attracts bonding electrons to itself. When atoms bond, they either share electrons, transfer electrons, or have an uneven sharing, leading to different types of chemical bonds.

Ionic Bond

An ionic bond is a type of chemical bond where one atom donates one or more electrons to another atom, resulting in the formation of ions. These ions are oppositely charged: the electron donor becomes a positively charged cation, while the electron acceptor becomes a negatively charged anion. Ionic bonds typically form between metals and nonmetals due to their significant difference in electronegativity; typically, an electronegativity difference of 1.7 or greater indicates an ionic bond.

In an ionic compound such as sodium chloride (NaCl), the sodium (Na) atom donates its electron to the chlorine (Cl) atom, leading to a crystal lattice structure held together by the electrostatic forces between the ions. Ionic bonds create compounds with high melting and boiling points, and when dissolved in water, they can conduct electricity.

Covalent Bond

A covalent bond forms when two atoms share a pair of electrons. This type of bond often occurs between nonmetal atoms that have similar electronegativities, so there is no significant electronegativity difference causing one atom to donate electrons to the other. When the electronegativity difference is less than 0.4, atoms tend to share electrons equally, resulting in a pure covalent bond.

An example is the hydrogen molecule (H₂), where two hydrogen atoms share their single electrons, filling their valence shells and achieving a stable configuration. Covalent bonds can be single, double, or triple, depending on how many electron pairs are shared, and they give rise to a wide variety of molecular shapes and chemical properties.

Polar Covalent Bond

When two atoms with differing electronegativities share electrons, they do so unequally, resulting in a polar covalent bond. The more electronegative atom pulls the shared electrons closer to itself, leading to a partial negative charge, while the less electronegative atom acquires a partial positive charge. A polar covalent bond typically has an electronegativity difference between 0.4 and 1.7.

Water (H₂O) is a classic example of a molecule with polar covalent bonds. The oxygen atom, being more electronegative, attracts the shared electrons more than the hydrogen atoms, creating a dipole moment. Polar molecules have distinct physical properties such as a higher boiling point compared to nonpolar molecules and the ability to dissolve in water or other polar solvents.