Question

Using the table of electronegativity values (Table \(11.5\) ), indicate which element is more positive and which is more negative in these compounds: (a) \(\mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{RbCl}\) (c) \(\mathrm{NH}_{3}\) (d) \(\mathrm{PbS}\) (e) \(\mathrm{PF}_{3}\) (f) \(\mathrm{CH}_{4}\)

Short answer

The more positive elements are H in H2O, Rb in RbCl, H in NH3, Pb in PbS, P in PF3, and C in CH4. The more negative elements are O in H2O, Cl in RbCl, N in NH3, S in PbS, F in PF3, and H in CH4 (though in CH4, it's a special nonpolar case).

Step-by-step solution

Identify Electronegativity Values for Each Atom

Refer to Table 11.5 to find the electronegativity values for hydrogen (H), oxygen (O), rubidium (Rb), chlorine (Cl), nitrogen (N), phosphorus (Pb and P), sulfur (S), fluorine (F), carbon (C), and lead (Pb). Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

Compare Electronegativity Values in H2O

Oxygen has a higher electronegativity than hydrogen. Therefore, in a water molecule \(\mathrm{H}_{2}O\), oxygen is the more negative atom, while hydrogen is the more positive atom.

Compare Electronegativity Values in RbCl

Chlorine has a higher electronegativity than rubidium. In rubidium chloride \(\mathrm{RbCl}\), chlorine is the more negative atom, and rubidium is the more positive atom.

Compare Electronegativity Values in NH3

Nitrogen has a higher electronegativity than hydrogen. In ammonia \(\mathrm{NH}_{3}\), nitrogen is the more negative atom, while hydrogen is the more positive atom.

Compare Electronegativity Values in PbS

Sulfur has a higher electronegativity than lead. In lead(II) sulfide \(\mathrm{PbS}\), sulfur is the more negative atom, and lead is the more positive atom.

Compare Electronegativity Values in PF3

Fluorine has a higher electronegativity than phosphorus. In phosphorus trifluoride \(\mathrm{PF}_{3}\), fluorine is the more negative atom, and phosphorus is the more positive atom.

Compare Electronegativity Values in CH4

Carbon has a lower electronegativity than hydrogen, however, the difference is very minimal. Methane \(\mathrm{CH}_{4}\) tends to be nonpolar. Nonetheless, if we must choose, then carbon would be the more positive atom, and hydrogen would be the more negative atom, but this is a special case due to the small difference in electronegativity and the symmetrical shape of the molecule, leading to a cancellation of dipole moments.

Key concepts

Chemical Bonding

Chemical bonding is the force that holds atoms together in compounds. These bonds form as a result of the interaction between the valence electrons of different atoms. Atoms can achieve a more stable electron configuration by sharing, donating, or receiving electrons, which leads to the formation of different types of bonds such as ionic, covalent, and metallic bonds.

For instance, in an ionic bond, one atom will donate an electron to another, creating ions with opposite charges that are held together by electrostatic forces. For example, in the case of \(\mathrm{RbCl}\), rubidium donates an electron to chlorine, resulting in positively charged rubidium ions (Rb+) and negatively charged chloride ions (Cl-).

In contrast, covalent bonds involve the sharing of electrons between atoms, as seen in \(\mathrm{H}_2\mathrm{O}\), where hydrogen shares its electrons with oxygen to form a stable electron arrangement for both participants. Understanding the nature of these bonds helps us to anticipate the properties of the compounds formed.

Molecular Polarity

Molecular polarity refers to the distribution of electrical charge over the atoms joined by the chemical bond. If the charge is equally distributed, the molecule is nonpolar, meaning there are no positive or negative poles. When electrons are shared unevenly between atoms due to differences in electronegativity, it results in a dipole with a positive end and a negative end, making the molecule polar.

Water (\(\mathrm{H}_2\mathrm{O}\)) is a classic example of a polar molecule due to the high electronegativity of oxygen compared to hydrogen. This causes a partial negative charge on the oxygen atom, while the hydrogens bear partial positive charges. The bent shape of water amplifies its polarity, as the dipoles don't cancel out. In contrast, methane (\(\mathrm{CH}_4\)) is nonpolar because of the symmetrical arrangement of hydrogen atoms around the central carbon atom, which leads to an even charge distribution despite the small difference in electronegativity.

Compound Electronegativity Differences

The concept of electronegativity differences between atoms in a compound is essential to determining the type of bond and polarity of a molecule. Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond, and these differences can predict bond characteristics.

Substantial electronegativity differences, as observed between metals and nonmetals like in \(\mathrm{RbCl}\) and \(\mathrm{PbS}\), often lead to the formation of ionic compounds. In contrast, smaller differences, usually found between nonmetals such as in \(\mathrm{NH}_3\) and \(\mathrm{PF}_3\), result in polar covalent bonds with uneven electron sharing and charged poles on the molecules. When electronegativity differences are insignificant or nonexistent, as with carbon and hydrogen in \(\mathrm{CH}_4\), the compound typically forms nonpolar covalent bonds. Recognizing these differences is crucial for predicting a compound's structure and its physical and chemical properties.