Shapes of Molecules/Intermolecular Forces

Question 1

State the shape of this molecule, give the bond angles and give one example of a molecule with this structure.

Answer 1

1. It is ‘linear’.
2. The bonds are 180° apart.
3. A molecule with this structure is CO₂.

Question 2

State the shape of this molecule, give the bond angles and give one example of a molecule with this structure.

Answer 2

1. It is ‘trigonal planar’.
2. The bonds are 120° apart.
3. A molecule with this structure is BF₃.

Question 3

State the shape of this molecule, give the bond angles and give one example of a molecule with this structure.

Answer 3

1. It is ‘tetrahedral’.
2. The bonds are 109.5° apart.
3. A molecule with this structure is methane (CH₄).

Question 4

State the shape of this molecule, give the bond angles and give one example of a molecule with this structure.

Answer 4

1. It is ‘v-shaped’.
2. The bonds are 104° apart.
3. A molecule with this structure is water (H₂O).

Question 5

State the shape of this molecule, give the bond angles and give one example of a molecule with this structure.

Answer 5

1. It is ‘trigonal pyramidal’.
2. The bonds are 107° apart.
3. A molecule with this structure is ammonia (NH₃).

Question 6

What shapes could molecues with three atoms possibly be?

Answer 6

1. V-shaped.
2. Linear.

Question 7

What shapes could molecules with four atoms possibly be?

Answer 7

1. Trigonal planar.
2. (Trigonal) pyramidal.

Question 8

What factors affect the bond angles in a molecule?

Answer 8

1. The type of electron pairs (i.e. bonding or lone pairs)
2. The number of electron pairs around the central atom.

Question 9

Explain van der Waal’s forces?

Answer 9

Van der Waal’s foces are caused by the movement of electrons with in a molecule. Van der Waals forces occur in non-polar covalent molecules.

Electrons move randomly within a bond. At times, the electons of a bond can be closer to one atom than the other. As a result a temopary polarity (dipole) is formed.

This temporary dipole can be attracted to other temporary dipoles or can induce other moloecules into having temporary dipoles, which it is then attracted to.

Van der Waal’s forces are weak compared to convalent bonds and can vary in strength. The more elctrons in a molecule the greater number of possible temporary dipoles and the stronger the force.

Question 10

What evidnece was there for van der Waal’s forces?

Answer 10

The fact that non-polar gases such as hydrogen and an oxygen can be liquidfied suggests that some attractive focre(van der Waal’s) exists between the molecules.

Question 11

Explain dipole-dipole interactoins between polar molecules.

Answer 11

This is a force that exists between polar substances. In polar substances, there is a permanent positive and negative dipole. These dipoles are attracted to oppositely charged dipoles of other molecules.

Because the dipoles are permanent, this force is stronger than that of van der Waal’s forces.

Question 12

What evidnece is ther for dipole-dipole interactions?

Answer 12

Ethene(M_r = 28) amd Metanal(M_r = 30) would be expected to have similar boiling points due to their similar M_r . However, ethene boils at 169K and methanal boils at 252K.

Methanal contains a C=O bond. This bond is very polar(electronegaitve value of 1.0). A permenmant dipole exist with O being negative and C being positive.

This creates dipole-dipole interaction forces which are stronger than van der Waals forces than the van der Waals forces in Ethene

Question 13

What is hydrogen bonding?

Answer 13

Hydrogen bonding is a special type of dipole-dipole interaction between small, highly electronegative atoms such as O, N or F.

Question 14

Explain the lower boiling point of H₂S compared to H₂O

Answer 14

• The H-S bond is less polar than the O-H bond
• The partrial negative is more diffused in the larger Sulfur atom than the smaller Oxygen atom and thus is less effective
• There is no hydrogen bonding between the H₂S molecules.

Question 15

Compare water and the hydrides elements in IV and VI

Answer 15

• The boiling point of the hydrides in IV show the normal trends for boiling points with increasing with M_r
• The hydrides in the group VI show a similar trend except for H₂O.
• H2O showes an exceptionally high boiling point due to the hydorgen bonding between molecules.

Question 16

Compare the boiling points of ammonia, hydrogen fluoride and water.

Answer 16

• The boiling point of water is much higher than that of ammonia. While ammonia can form one more hydrogen bond that water, oxygen has a much higher electronegative value than nitrogen, leading to stronger hydrogen bonds.
• The boiling point of water is higher than hydrogen fluoride. The greater number of hydrogen bond each water molecule can form is of a greater significance than the stronger hydrigen bonds in hydrogen fluoride due to fluorine’s higer electronegativity.

Question 17

Why is there no evidence for hydorgen bonding in hydrogen chloride?

Answer 17

Despite the sizable electronegative diffierence in hydrogen chloride, the chlorine atom is too large for hydrogen bonding and consquently diffuses the charge on the atom

Question 18

Why can only small electronegative atons form hydrigen bonds?

Answer 18

When hydrogen is bonded to a very electronegative small atom it’s share of the bonding electrons is very small, so that it is essentiallly a bare proton.

This creates a very positive diople which readily attracts to other eletronegative atoms. If this atom has a lone pair of electrons,in a small orbit, it can share this pair with the hydrogen atom.

The hydrogen atom is now bonded to two small atoms, one convalently and the other through hydrogen bonding.