Chapter 6: Shapes of Molecules and Intermolecular Forces

Question 1

Explain electron-pair repulsion theory.

Answer 1

In molecules and polyatomic ions, electron pairs repel one another, hence, they are arranged as far apart as possible. The arrangement of electron pairs minimises repulsion and thus holds the bonded atoms in a definite shape. Different numbers of electron pairs result in different shapes.

Question 2

Show how chemists draw three-dimensional molecular shapes.

Answer 2

Chemists use wedges to help visual structures in three dimensions.

Question 3

Compare bonded and lone-pair electron repulsions.

Answer 3

A lone pair is slightly closer to the central atom, and occupies more space, than a bonded pair. This results in a lone pair repelling more strongly than a bonding pair.

Question 4

Explain the structure and shape of a molecule of methane, CH₄.

Answer 4

A molecule of methane, CH₄, is symmetrical with four C–H covalent bonds. Four bonded pairs surround the central carbon atom. The four electron pairs repel one another as far apart as possible in three-dimensional space.

The result is a tetrahedral shape with four equal H–C–H bond angles of 109.5°.

Question 5

Methane (CH₄), ammonia (NH₃) and water (H₂O) all have four electron pairs surrounding the central atom, but in ammonia and water, the electron pairs are a mixture of bonded pairs and lone pairs.

Explain how the bond angle and shape of an ammonia and a water molecule is determined.

Answer 5

• The four electron pairs around the central atom repel one another as far apart as possible into a tetrahedral arrangement.
• Lone pairs repel more strongly than bonded pairs.
• Therefore, lone pairs repel bonded pairs slightly closer together, decreasing the bond angle — the angle between the bonded pairs of electrons.
• The bond angle is reduced about 2.5° for each lone pair.

Question 6

In determining molecular shape, how are molecules containing multiple bonds treated?

Answer 6

In molecules containing multiple bonds, each multiple bond is treated as a bonding region.

Question 7

Describe and explain the shape and bond angles of boron trifluoride.

Answer 7

Boron trifluoride, BF₃, has only three bonded pairs around the cetral boron atom. Electron-pair repulsion gives a trigonal planar shape with equal bond angles of 120°.

Question 8

Describe and explain the shape and bond angles of sulfur hexafluoride.

Answer 8

Sulfur hexafluoride, SF6, has six bonded pairs of electrons around the central sulfur atom. Electron-pair repulsion gives an octahedral shape with equal bond angles of 90°.

Question 9

Describe and explain the shape and bond angles for an ammonium ion.

Answer 9

The ammonium ion, NH₄⁺, has four bonded pairs surrounding the central nitrogen atom. This is the same number of bonded electron pairs as a methane molecule, therefore, NH₄⁺ has the same tetrahedral arrangement and bond angles (109.5°) as a methane molecule.

Question 10

Describe and explain the shapes and bond angles of:

1. the carbonate ion, CO₃^2–
2. the nitrate ion, NO₃^–
3. the sulfate ion, SO₄^2–.

Answer 10

1. CO₃^2– has three regions of electron density surrounding the central atom. So it has the same shape as the BF₃ molecule: trigonal planar shape with bond angles of 120°.
2. NO₃^– has three regions of electron density surrounding the central atom. So it has the same shape as the BF₃ molecule: trigonal planar shape with bond angles of 120°.
3. SO₄^2– has four regions of electron density surrounding the central atom. So it has the same shape as a methane molecule: tetrahedral shape with bond angles of 109.5°.

Question 11

1. What is electronegativity? Explain.
2. How is electronegativity measured?
3. What is the periodic trend in electronegativity. Explain.
4. What elements are most electronegative?
5. What elements are least electronegative?

Answer 11

1. Shared pairs of electrons in a covalent bond often experience more attraction from one of the bonded atoms than the other. Electronegativity is a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond.
2. The Pauling scale is used to compre the electronegativity of the atoms of different elements.
3. Across the periodic table, the nuclear charge increases while the atomic radius decreases. Hence, electronegativity increases across the periodic table.
   Up the periodic table, the energy level of the outer shell decreases so the distance between the outer shell and the nucleus decreases. Hence, electronegativity increases up the periodic table.
4. The non-metals nitrogen, oxygen, fluorine and chlorine have the most electronegative atoms.
5. The Group 1 metals, including lithium, sodium and potassium have the least electronegative atoms.

Question 12

1. Explain the effects of their being a large difference in electronegativity between two atoms.
2. In what regions of electronegativity difference are bonds covalent, polar covalent and ionic.

Answer 12

1. If the electronegativity difference is large, one bonded atom will have a much greater attraction for the shared pair than the other bonded atom. The more electronegative atom will have gained control of the electrons and the bond will now be ionic rather than covalent.
2. When the electronegativity difference is 0, the bond is covalent. When the electronegativity difference is between 0 and 1.8, the bond is polar covalent. When the electronegativity difference is higher than 1.8, the bond is ionic.

Question 13

1. What is a non-polar bond? Under what circumstances is a bond non-polar?
2. What is a polar covalent bond? Under what circumstances is a bond polar?
3. What is a permanent dipole?
4. How does a permanent dipole distinguish from an induced dipole?

Answer 13

1. In a non-polar bond, the bonded electron pair is shared equally between the bonded atoms. A bond will be non-polar when the bonded atoms have the same electronegativity.
2. In a polar covalent bond, the bonded electron pair is shared unequally between the bonded atoms. A bond will be polar when the bonded atoms are different and have a different electronegativity.
3. A permanent dipole refers to the separation of opposite charges in a polar covalent bond (delta-positive on one atom and delta-negative on the other atom).
4. A permanent dipole does not change.

Question 14

What happens when there are two or more polar bonds in a molecule?

Answer 14

Depending on the shape of the molecule, the dipoles may reinforce one another to produce a larger dipole over the whole molecule, or cancel out if the dipoles act in opposite directions.

Question 15

Describe the polarity of water, H₂O.

Answer 15

A water molecule is polar. The two O–H bonds have a permanent dipole. The two dipoles act in different directions but they do not exactly oppose each other. Overall the oxygen end of the molecule has a delta-negative charge and the hydrogen end of the molecule has a delta-positive charge.

Question 16

Describe the polarity of carbon dioxide, CO₂.

Answer 16

A carbon dioxide molecule is non-polar. The two C=O bonds have a permanent dipole. The two dipoles act in opposite directions and exactly oppose one another. Therefore, over the whole molecule, the dipoles cancel and the overall dipole is zero.

Question 17

Describe and explain the process of solid sodium chloride dissolving in water to form Na⁺ (aq) and Cl^– (aq).

Answer 17

• Water molecules attract Na⁺ and Cl^– ions.
  
  • Na⁺ ions are attracted towards the oxygen end of the molecule (delta-negative).
  • Cl^– ions are attracted towards the hydrogen end of the molecule (delta-positive).
• The ionic lattice breaks down as it dissolves.
• In the resulting solution, water molecules surround the Na⁺ and Cl^– ions.

Question 18

1. Define intermolecular forces.
2. What types of intermolecular forces are there?
3. What properties do intermolecular forces factor?

Answer 18

1. Intermolecular forces are weak interactions between dipoles of different molecules.
2. Intermolecular forces fall into three main categories: induced dipole-dipole interactions (London forces); permanent dipole dipole interactions; and hydrogen bonding.
3. Intermolecular forces are largely responsible for physical properties such as melting and boiling points.

Question 19

Give the typical ranges of bond enthalpies for the three types on intermolecular forces. Compare this to the bond enthalpy of a single covalent bond.

Answer 19

Question 20

1. What are induced dipole-dipole interactions (London forces)?
2. Specifically, where do London forces act?
3. Describe the temporary nature of induced dipoles.

Answer 20

1. London forces are weak intermolecular forces that exist between all molecules, whether polar or non-polar.
2. They act between induced dipoles in different molecules.
3. Induced dipoles are temporary. In the next instant of time, the induced dipoles may dissapear, only for the whole process to take place amongst other molecules.

Question 21

Describe and explain the origin of induced dipoles.

Answer 21

The movement of electrons produces a changing dipole in a molecule. At any instant, an instantaneous dipole will exist, but its position is constantly shifting.

The instantaneous dipole induces a dipole on a neighbouring molecule.

The induced dipole induces further dipoles on neighbouring molecules, which then attract one another.

Question 22

1. Explain the factors that affect the strength of induced dipole-dipole interactions.
2. How does this affect boiling points?

Answer 22

1. Induced dipoles result from interactions of electrons between molecules. The more electrons in each molecule, the larger the instantaneous and induced dipole, and the greater the induced dipole-dipole interactions, and the stronger the attractive forces between molecules.
2. Larger numbers of electrons mean larger induced dipoles. So more energy is required to overcome the intermolecular forces, increasing the boiling point.

Question 23

Compare the intermolecular forces in hydrogen chloride with the intermolecular forces in fluorine.

Where do these forces act?

Why is there a difference in boiling point?

Answer 23

• Fluorine molecules are non-polar and only have London forces between molecules. London forces act between induced dipoles in different molecules, regardless of polarity.
• Hydrogen chloride molecules are polar and have London forces and permanent dipole-dipole interactions between molecules. Permanent dipole-dipole interactions act between the permanent dipoles in different polar molecules.
• Extra energy is needed to break the additional permanent dipole-dipole interactions between hydrogen chloride molecules.
• The boiling point of hydrogen chloride is therefore higher than fluorine.

Question 24

1. What is a simple molecular substance?
2. Describe the structure of simple molecules in the solid state.

Answer 24

1. A simple molecular substance is made up of simple molecules — small units containing a definite number of atoms with a definite molecular formula, such as neon (Ne) or water (H₂O).
2. In the solid state, simple molecules form a regular structure called a simple molecular lattice where the molecules are held in place by weak intermolecular forces. The atoms within each molecule are bonded together strongly by covalent bonds.

Question 25

Describe and explain the physical state of simple molecules. Refer to melting and boiling points.

Answer 25

At room temperature, simple molecular substances may exist as solids, liquids or gases. Simple molecular substances have low melting and boiling points as the weak intermolecular forces can be broken with relatively little energy. The covalent bonds are strong and do not break.

Question 26

Describe and explain the solubility of non-polar simple molecular substances in non-polar solvents.

Answer 26

• When a simple molecular compound is added to a non-polar solvent (e.g. hexane) intermolecular forces can form between the molecules and the solvent.
• The interactions weaken the intermolecular forces in the simple molecular lattice. The intermolecular forces break and the compound dissolves.

Therefore, non-polar simple molecular substances tend to be soluble in non-polar solvents.

Question 27

Describe and explain the solubility of non-polar simple molecular substances in polar solvents.

Answer 27

• When a non-polar simple molecular substance is added to a polar solvent, there is little interaction between the molecules in the lattice and the solvent molecules.
• The intermolecular bonding within the polar solvent is too strong to be broken.

Therefore non-polar simple molecular substances tend to be insoluble in polar solvents.

Question 28

Describe and explain the solubility of polar simple molecular substances in polar solvents. What is it dependant on?

Answer 28

Polar simple molecular substances may dissolve in polar solvents as the polar solute molecules and the polar solvent molecules can attract each other. The solubility depends on the strength of the dipole and can be hard to predict.

Question 29

1. Explain why sugar dissolves in water.
2. Show how this solubility extends to liquids and gases by using an example.

Answer 29

1. Sugar is a polar covalent compound with many polar O–H bonds, which attract and bond with polar water molecules.
2. This solubility can extend to liquids and gases. For example, hydrogen chloride is a gas with a polar H–Cl bond that is extremely soluble in water, forming hydrochloric acid

Question 30

Explain the electrical conductivity of simple molecular structures.

Answer 30

• There are no mobile charged particles in simple molecular structures.
• With no charged particles that can move, there is nothing to complete an electrical circuit.

Therefore simple molecular structures are non-conductors of electricity.

Question 31

What is a hydrogen bond?

Answer 31

A hydrogen bond is a special type of permanent dipole-dipole interaction found between molecules containing:

• an electronegative atom with a lone pair of electrons (e.g. nitrogen, oxygen, or fluorine)
• a hydrogen atom attached to an electronegative atom (e.g. H–O, H–N, or H–F).

Question 32

1. Where does a hydrogen bond act?
2. How relatively strong are hydrogen bonds?
3. How is a hydrogen bond drawn?

Answer 32

1. The hydrogen bond acts between a lone pair of electrons on an electronegative atom in one molecule and a hydrogen atom in a different molecule.
2. Hydrogen bonds are the strongest type of intermolecular interactions.
3. A hydrogen bond is shown by a dashed line.

Question 33

Show the hydrogen bonding between molecules of water (H₂O) and ammonia (NH₃).

Answer 33

Question 34

Explain the unique properties that hydrogen bonding gives water that supports the existence of life on Earth.

Answer 34

• Hydrogen bonds hold water molecules apart in an open lattice structure. There are two lone pairs on the oxygen atom and two hydrogen atoms, so each water molecule can form four hydrogen bonds.
• In solid ice, the hydrogen bonds extend outwards, holding water molecules slightly apart and form an open tetrahedral lattice full of holes. The bond angle about the hydrogen atom involved in the hydrogen bond is close to 180°.
• The holes in the open lattice structure decrease the density of water on freezing. When ice melts, the ice lattice collapses and the molecules move closer together. So liquid water is denser than solid ice.
• Solid ice is less dense than liquid water and floats.

Question 35

Explain why water has a relatively high melting and boiling point.

Answer 35

As with all molecules, water has London forces between molecules.

• Hydrogen bonds are extra forces, on top of the London forces.
• A relatively large quantity of energy is needed to break the hydrogen bonds in water, so water has much higher melting and boiling points than would be expected from just London forces.
• When the ice lattice breaks, the rigid arrangement of hydrogen bonds in ice is broken. When water boils, the hydrogen bonds break completely.