C6- shapes of molecules and intermolecular forces

Question 1

What is the arrangement of electron pairs and why?

Answer 1

Electron pairs repel each other so that they are arranged as far apart as possible.
The arrangement of electron pairs minimises repulsion and holds bonded atoms in a definite shapes.

Question 2

What is the electron pair repulsion theory?

Answer 2

A model used for explaining and predicting the shapes of molecules and polyatomic ions.
Different number of electron pairs result in different shapes.

Question 3

How is each bond in a molecule represented in the 3D representation?

Answer 3

Solid line : bond in the plane of the paper

Solid wedge: bond comes out of the plane of the paper (front)

Dashed wedge: bond goes into the plane of the paper (back)

Question 4

What is a lone pair of electrons and how do they affect the bond angle?

Answer 4

a pair of valence electrons that are not shared with another atom

A lone pair of electrons is closer to the central atom and occupies more space than a bonded pair.
Therefore a lone pair repels more strongly than a bonded pair- decreasing the bond angle.
The bond angle is reduced by 2.5* for each lone pair

Question 5

What is meant by bond angle?

Answer 5

The angle between the bonded pairs of electrons.

Question 6

What is the order of pair repulsion from strongest to weakest?

Answer 6

lone pair/ lone pair

Bonded pair/ lone pair

Bonded pair/ bonded pair

Question 7

Linear

Answer 7

2 bonded pairs
0 lone pairs
180*

Question 8

Trigonal planar

Answer 8

3 bonded pairs
0 loan pairs
120*

Question 9

Tetrahedral

Answer 9

4 bonded pairs
0 loan pairs
109.5*

Question 10

Octahedral

Answer 10

6 bonded pairs
0 loan pairs
90*

Question 11

Trigonal pyramidal

Answer 11

3 bonded pairs
1 lone pair
107* (109.5 - 2.5 )

Question 12

Non- linear

Answer 12

2 bonded pairs
2 loan pairs
104.5* (109.5 - (2.5x2) )

Question 13

Bent

Answer 13

2 bonded pairs
1 loan pair
120*

Question 14

What is the bond angle for NH3, CH4, H2O ?

Answer 14

NH3: 107
CH4: 109.5
H2O: 104.5

Question 15

What is electronegativity?

Answer 15

The attraction of a bonded atom for the pair of electrons in a covalent bond.

Atom with larger electronegativity value has delta- charge.
Atom with smaller electronegativity value has delta+ charge /

Question 16

In a covalent bond, what happens with the bonded atoms and the electrons?

Answer 16

The nuclei of the bonded atoms attract the shared pair of electrons.
In molecules of elements (eg. H2), the atoms are the same element and so the bonded electron pair is shared evenly.

Question 17

What are the changes in a covalent bond when the bonded atoms are different elements?

Answer 17

• nuclear charges are different.
• atoms may be different sizes.
• shared pair of electrons may be closer to one nucleus and may now experience more attraction from one of the bonded atoms than the other.

Question 18

What is the Pauling scale and what does it show?

Answer 18

The Pauling used to compare electronegativity:

-As you go across the periodic table electronegativity increases.
(Across the table nuclear charge increases & atomic radius decreases)

-As you go up the periodic table electronegativity also increases.

Question 19

How can you predict the bond type using electronegativity?

Answer 19

The difference in electronegativity tells you the bond type:

Covalent: 0
Polar covalent: 0-1.8
Ionic: greater than 1.8

Question 20

Explain a non-polar bond

Answer 20

The bonded electron pair is shared equally between the bonded atoms.
A bond will be non-polar when:
- bonded atoms are the same
- bonded atoms have the same/similar electronegativity

Question 21

What is a pure covalent bond?

Answer 21

In molecules of elements, the bonded atoms come from the same element and the electron pair is shared equally so therefore it is a pure covalent bond.

Question 22

Explain polar bonds

Answer 22

The bonded electron pair is shared unequally between bonded atoms. Happens when:

• bonded atoms are different
• bonded atoms have different electronegativities.

= polar covalent bond:
Contains a permanent dipole as the dipole does not change.

Question 23

What is a dipole?

Answer 23

The separation of opposite charges.

Represented by an arrow in the direction of the dipole movement.

Question 24

What are intermolecular forces and the 3 main categories?

Answer 24

Weak interactions between dipoles of different molecules.

> Vanderwaal’s forces:

• induced dipole-dipole interactions ( london forces)
• permanent dipole-dipole interactions.

-hydrogen bonding.

Question 25

What are intermolecular forces and covalent bonds each responsible for?

Answer 25

Intermolecular forces are responsible for physical properties such as melting and boiling points.
Covalent bonds determine the identity and chemical reactions of molecules.

Question 26

Explain how induced dipole-dipole interactions (london forces) are created.

Answer 26

They exist between all molecules but are TEMPORARY.

1. Electron movement produces a changing dipole in a molecules.
2. At any instant, and instantaneous dipole will exist (very weak as it only lats for an instant) but its position is constantly shifting.
3. The instantaneous dipole induces a dipole on a neighbouring particle.
4. Further dipoles are induced which then attract to each other.

They exist over short distances so molecules need to be close together to induce dipoles.

Question 27

How does the number of electrons in a molecule affect an induced dipole?

Answer 27

The more electrons in each molecules:

• the larger the instantaneous and induced dipoles
• the greater the induced dipole-dipole interactions.
• the stronger the attractive forces between molecules. (This increases the boiling point as more energy is needed to overcome the IMF)

Question 28

Explain permanent dipole-dipole interactions.

Answer 28

Act between the permanent dipoles in different polar molecules.

If a molecule has London forces and permanent dipole- dipole interactions between molecules, extra energy is needed to break the additional permanent dipole- dipole interactions so therefore boiling point is often high.

Question 29

Explain a simple molecular substance and its structure.

Answer 29

A simple molecular substance is made up of simple molecules- small units containing a definite number of atoms with a definite molecular formula.

In the solid state, they can form a simple molecular lattice:

• the molecules are held in place by weak intermolecular forces.
• the atoms within each molecule are bonded together strongly by covalent bonds.

Question 30

Explain the melting and boiling point of simple molecular substances.

Answer 30

They have a low melting and boiling point:

All simple molecular substances can be solidified into simple molecular lattices by reducing the temperature.

The weak intermolecular forces in the lattice can be broken even by the energy present at low temperatures.
During melting the covalent bonds do not break.

Question 31

Explain the solubility (non-polar) of simple molecular substances.

Answer 31

Non-polar simple molecular substances tend to be soluble in non-polar solvents:
> simple molecular compound+ non-polar solvent = intermolecular forces form, therefore, weaken the lattice and the compound dissolves.

They tend to be insoluble in polar solvents:
>little interaction between the lattice and solvent so intermolecular bonding within polar solvent is too strong to break.

Question 32

Explain the solubility (polar) of simple molecular substances.

Answer 32

• Polar covalent substances may dissolve in polar solvents as the solute molecules and solvent molecules attract each other.
• solubility depends on the strength of the dipole (so it’s hard to predict)
• compounds that contain both polar and non-polar parts can dissolve in either solvents.
• some molecules = hydrophobic (non-polar) and hydrophilic (polar so interacts with water)

Question 33

Explain the electrical conductivity in simple molecular substances.

Answer 33

There are no mobile charges particles so therefore they are non-conductors of electricity.

Question 34

Explain a hydrogen bond.

Answer 34

A special type of permanent dipole- dipole interaction found between molecules containing:

• an electronegative atom with a lone pair of electrons. (Eg. Oxygen, fluorine,nitrogen).
• a hydrogen atom attached to an electronegative atom.

Hydrogen bonds are the strongest type of intermolecular attraction.

The shape around the hydrogen atom involved in the hydrogen bond is linear.

Question 35

Explain the first anomalous property of water (ice is less dense than liquid water).

Answer 35

With 2 lone pairs on the oxygen atom and two hydrogen atoms, each water molecule can form 4 hydrogen bonds.
The hydrogen bond extend outwards, holding water molecules slightly apart and forming an open tetrahedral lattice full of holes.
(The bond angle about the hydrogen atom involved in the hydrogen bond is 180*)

The holes in the open lattice structure decrease the density of water on freezing.
When ice melts, the lattice collapses and the molecules move closer together making liquid more dense.

Question 36

Explain the second and third anomalous properties of water (high MP & BP , high surface tension& viscosity).

Answer 36

Water has a relatively high melting and boiling point:
water has London forces between molecules.
- hydrogen bonds are extra forces, above London forces so a large quantity of energy is needed to break the hydrogen bonds in water.

• when the ice lattice breaks, the rigid arrangement of hydrogen bonds in ice is broken. When water boils, the hydrogen bonds break completely.

Without hydrogen bonds water would have a boiling point of about -75*C and woukd exist as a gas at room temperature and pressure.

> Water also has a relatively high surface tension and viscosity

Question 37

How does ionic bonding hold an ionic compound together?

Answer 37

Electrostatic attraction between positively charged ions

Question 38

Explain how you would predict the bond angle in a molecules.

Answer 38

1. Draw dot and cross diagram.
2. Show loan pairs clearly.
3. State the number of bonded pairs and loan pairs.
4. Reduce 2.5* for each loan pair if necessary.
5. State the bond angle.