Question

For each of the following bonds, draw a figure indicating the direction of the bond dipole, including which end of the bond is positive and which is negative. For each of the following bonds, draw a figure indicating the direction of the bond dipole, including which end of the bond is positive and which is negative. a. \(\mathrm{P}-\mathrm{F}\) b. \(\mathrm{P}-\mathrm{O}\) c. \(\mathrm{P}-\mathrm{C}\) d. \(\mathrm{P}-\mathrm{H}\)

Short answer

The bond dipoles for the given bonds are as follows: a. P-F bond: \( \overset{ - }{F} \leftarrow \overset{ + }{P} \) b. P-O bond: \( \overset{ - }{O} \leftarrow \overset{ + }{P}\) c. P-C bond: \( \overset{ - }{C} \leftarrow \overset{ + }{P}\) d. P-H bond: \( \overset{ - }{P} \leftarrow \overset{ + }{H}\)

Step-by-step solution

Analyzing electronegativity values

Before drawing the dipoles, we need to know the electronegativity values of the elements involved in each bond. Electronegativity values have a tendency to decrease down a group and increase across a period in the periodic table. Using the periodic table, we can approximate electronegativity values for the relevant elements involved in each bond to determine the direction of the dipole.

Direction of bond dipoles for each bond

a. P-F bond: Since fluorine is the most electronegative element in the periodic table, the bond dipole will point towards the F atom. P is less electronegative than F; thus, it will have a partial positive charge, and F will have a partial negative charge. We can represent the bond dipole as \( \overset{ - }{F} \leftarrow \overset{ + }{P} \). b. P-O bond: Oxygen is quite electronegative. However, P is less electronegative than O. Thus, the bond dipole will point towards the O atom. We can represent the bond dipole as \( \overset{ - }{O} \leftarrow \overset{ + }{P}\). c. P-C bond: Carbon and phosphorus have similar electronegativity values. However, since C is slightly more electronegative than P, the bond dipole will point towards the C atom, but it will be small due to their similar electronegativity. We can represent the bond dipole as \( \overset{ - }{C} \leftarrow \overset{ + }{P}\). d. P-H bond: Hydrogen is less electronegative than phosphorus. Therefore, the bond dipole will point towards the P atom, and the P atom will have a partial negative charge. We can represent the bond dipole as \( \overset{ - }{P} \leftarrow \overset{ + }{H}\).

Drawing figures for each bond

Now that we've determined the bond dipoles' directions, we can draw figures for each bond: a. P-F bond: b. P-O bond: +(P)-->(F)- +(P)-->(O)- c. P-C bond: d. P-H bond: +(P)->(C)- -(P)<--(H)+

Key concepts

Electronegativity

Electronegativity is a fundamental concept in chemistry which describes the tendency of an atom to attract electrons towards itself when it forms a chemical bond. The more electronegative an atom is, the more strongly it will pull shared electrons in a covalent bond towards itself, resulting in a polar bond if it's bonded to a less electronegative atom.

Elements display a range of electronegativities, measured on the Pauling scale, with fluorine holding the highest value of around 4.0. This scale is useful when trying to determine the charge distribution within a molecule, and thus the direction of the bond dipoles. Understanding electronegativity is crucial when predicting whether a bond will be ionic, covalent, or polar covalent. Fluorine being the most electronegative element, will always have a partial negative charge in a bond with another nonmetal, as electrons will be closer to it.

Bond Polarity

Bond polarity is a direct consequence of the differences in electronegativity between two bonded atoms. A polar bond has a positive end and a negative end, known as dipoles, due to an unequal sharing of electrons. When two atoms have an electronegativity difference usually greater than 0.5, a bond dipole is established.

For instance, in a P-F bond, the significant electronegativity difference causes the electrons to be drawn towards the fluorine atom, leading to a negative charge on the fluorine end and a positive charge on the phosphorus end. This dipole indicates that the bond between the atoms is polar. The larger the difference in electronegativity, the more polarized the bond will be, affecting the molecule's physical properties, such as solubility and boiling points.

Periodic Table Trends

The periodic table is designed in such a way that it reveals trends in the properties of elements, including electronegativity. Electronegativity tends to increase across a period, from left to right, and decrease down a group on the periodic table. As we move from left to right, the number of protons in the nucleus increases, pulling the electrons closer and thus increasing the ability of the atoms to attract electrons.

Conversely, as we move down a group, the shells of electrons increase, the increase in distance and the shielding effect cause the electronegativity to decrease. These trends are key to predicting the polarity of bonds and understanding reactive behaviors. For example, when comparing elements such as phosphorus, oxygen, carbon, and hydrogen, we can anticipate the bond polarity by looking at their positions on the periodic table.

Chemical Bonding

Chemical bonding involves the joining of atoms to form new substances. Bonds can be ionic, covalent, or metallic. In an ionic bond, electrons are transferred from one atom to another leading to the formation of oppositely charged ions. Covalent bonds involve the sharing of electron pairs between atoms, while metallic bonds are the force of attraction between valence electrons and the metal ions.

Covalent bonds can be nonpolar, where there is an equal sharing of electrons (usually between atoms of the same element), or polar, where there is a partial transfer of electron density due to differences in electronegativity. The type and degree of bonding greatly influence the properties of compounds, such as hardness, electrical and thermal conductivity, melting and boiling points, and solubility. In the context of phosphorus compounds in the exercise, understanding the nature of the chemical bonds and how they form between diverse elements helps us predict the molecular structure and behavior of the compounds.