Question

(a) Draw a dot diagram for \(\mathrm{NH}_{3}\) and one for \(\mathrm{PH}_{3}\). (b) Is either molecule polar? (The electronegativities of \(\mathrm{N}, \mathrm{P}\), and \(\mathrm{H}\) are \(3.0,2.1\), and 2.1, respectively.) (c) In which substance are the London forces stronger? Explain. (d) Basing your answer solely on the London forces in the two substances, which substance would you expect to have the higher boiling point? Explain.

Short answer

(a) For NH3, the central Nitrogen atom is surrounded by three Hydrogen atoms, with two lone pair electrons on Nitrogen. For PH3, the central Phosphorus atom is surrounded by three Hydrogen atoms but no lone pairs. (b) NH3 is polar (\(|3.0-2.1|=0.9)\), while PH3 is nonpolar (\(|2.1-2.1|=0\)). (c) London forces are stronger in PH3 (14 electrons) than in NH3 (10 electrons). (d) PH3 is expected to have a higher boiling point due to stronger London forces.

Step-by-step solution

Drawing Dot Diagrams for NH3 and PH3

To draw the dot diagrams, we need to consider the valence electrons of the atoms involved in the molecules (Nitrogen, Phosphorus, and Hydrogen). For NH3: 1. Nitrogen (N) has 5 valence electrons. 2. Hydrogen (H) has 1 valence electron. 3. The molecule has three Hydrogen atoms and one Nitrogen atom, so the total count of valence electrons becomes 5 + 3*(1) = 8. Place the Nitrogen atom in the center and surround it with three Hydrogen atoms. Then, distribute the valence electrons in such a way that hydrogen completes 1 valence electron and nitrogen completes 8, which is the number of electrons required for a stable shell structure. For PH3: 1. Phosphorus (P) has 5 valence electrons. 2. Hydrogen (H) has 1 valence electron. 3. Repeat the process stated for NH3 but replace Nitrogen with Phosphorus as the central atom.

Checking Polarity of the Molecules

Polarity depends on the difference in electronegativities between bonded atoms. Calculate the electronegativity difference between Nitrogen and Hydrogen and between Phosphorus and Hydrogen, using the electronegativity values provided (N: 3.0, P: 2.1, H: 2.1). For NH3: Electronegativity difference between N and H: \(|3.0 - 2.1| = 0.9\). For PH3: Electronegativity difference between P and H: \(|2.1 - 2.1| = 0\). In general, a difference greater than 0.5 is considered sufficient for a bond to be considered polar. Thus, NH3 is a polar molecule, whereas PH3 is nonpolar.

Comparing London Forces

London forces are temporary attractive forces that result from the random distribution of electrons in molecules. Larger molecules with more electrons usually experience stronger London forces. NH3 has 10 electrons: 7 from Nitrogen and 3 from Hydrogens. PH3 has 14 electrons: 5 from Phosphorus and 3 from Hydrogens. Since PH3 has more electrons, the London forces are stronger in PH3 than in NH3.

Predicting Boiling Points from London Forces

As mentioned earlier, London forces are generally stronger in larger molecules with more electrons. Substances with stronger London forces usually have higher boiling points because more energy is required to overcome these forces. Since PH3 has stronger London forces than NH3, we would expect PH3 to have a higher boiling point due to London forces alone.

Key concepts

Dot Diagrams

In chemical bonding, dot diagrams help us visualize the valence electrons in a molecule. They are also known as Lewis structures. These diagrams show how electrons are shared or transferred between atoms, allowing us to predict molecular geometry and bonding characteristics.

To create a dot diagram, you need to know the number of valence electrons for each atom. For \(NH_3\), nitrogen has 5 valence electrons, while each hydrogen atom has 1. This totals to 8 valence electrons that need to be arranged so that nitrogen is in the center, forming bonds with three hydrogens.

Similarly, in \(PH_3\), phosphorus has 5 valence electrons. Again, with three hydrogens contributing 1 electron each, a total of 8 electrons are distributed with phosphorus at the center.

Polar Molecules

A molecule's polarity is determined by the distribution of electrical charge over its atoms. When a molecule is polar, it has a slightly positive and slightly negative end due to an uneven distribution of electrons.

The key to determining polarity is the electronegativity difference between the atoms involved. For \(NH_3\), the difference between nitrogen (3.0) and hydrogen (2.1) is 0.9, which signifies a polar molecule.

• A polar molecule has an asymmetric shape and clear positive and negative poles.
• NH₃’s trigonal pyramidal structure contributes to its polarity.

On the other hand, \(PH_3\) has an electronegativity difference of 0, making it nonpolar. Without a difference, charges are evenly distributed, eliminating poles.

London Dispersion Forces

These are the weakest intermolecular forces but are present in all molecules. London dispersion forces arise from temporary fluctuations in electron density, creating temporary dipoles.

The strength of London forces increases with the size of the electron cloud. This means that molecules with more electrons will have stronger dispersion forces. In the case of \(PH_3\) versus \(NH_3\), PH₃, with more electrons, shows stronger London forces.

• Larger electron clouds mean increased temporary dipoles.
• Even though they are weaker than other forces, their impact grows with molecular size.

These forces are particularly notable in heavier, nonpolar molecules.

Electronegativity Differences

Electronegativity measures an atom's ability to attract electrons within a bond. Differences in electronegativity between atoms determine bond polarity and, consequently, molecule polarity.

If the electronegativity difference is greater than 0.5, the bond is likely polar. For \(NH_3\), the nitrogen-hydrogen bond has a difference of 0.9, indicating polarity. In contrast, the phosphorus-hydrogen bond in \(PH_3\) has no difference, resulting in a nonpolar bond.

• Electronegativity differences can shape molecular geometry.
• They help predict molecular interactions, like solubility and boiling/melting points.

Understanding these differences is crucial for predicting how compounds behave in different environments.