6.1 Shapes of Molecules and Ions & 6.2 Electronegativity and Polarity

Question 1

How does the electron-pair repulsion theory work?

Answer 1

• Electrons all have negative charge, so electron pairs repel one another
• This causes them to be arranged as far as possible from each other to minimise repulsion
• The bonded atoms are therefore held in a definite shape (so electron pairs surrounding the central atom determine the shape of the molecule/ ion)

Question 2

How do you represent molecules in 3 dimensions?

Answer 2

• Imagine the 3D molecule in front of you, and a piece of paper (that’s surface is facing you) going through it vertically
• Change the orientation of the molecule so the piece of paper “cuts” through as many of the atoms as possible
• Draw the atoms in the plane of the paper (the ones the paper cuts through) using regular solid lines in their 3D positions
• Draw atoms coming out (in front) of the plane of paper with with a plain wedge
• Draw atoms going into (behind) the plane of paper with a dashed wedge
• You may have to rotate the molecule so that all of the atoms can be drawn, even if this means more will be drawn with wedges (e.g. sulfur hexafluoride)

Question 3

What are the 2 types of electron pairs? What does each mean?

Answer 3

• Bonded pairs of electrons are involved in a covalent bond
• Lone pairs of electrons are not in a covalent bond

Question 4

How do lone pairs differ to bonded pairs in terms of repulsion?

Answer 4

• Lone pairs repel more strongly than a bonded pair

Question 5

Molecules with 4 electron pairs are based off of what shape? How do lone pairs affect this?

Answer 5

• Tetrahedral
• Lone pairs repel more strongly than bonded pairs
• The bonded pairs are repelled by the lone pairs which push them closer together, which reduces their bond angle; the angle between a bonded pairs of electrons
• Each additional lone pair reduces the bond angle by 2.5°

Question 6

List the number of electron, lone and bonding pairs in methane, and therefore the name of its shape and the bond angle.

Answer 6

• 4 electron pairs
• 4 bonding pairs, no lone pairs
• Tetrahedral
• 109.5°

Question 7

List the number of electron, lone and bonding pairs in ammonia, and therefore the name of its shape and the bond angle.

Answer 7

• 4 electron pairs
• 3 bonding pairs, 1 lone pair
• Pyramidal
• 107°

Question 8

List the number of electron, lone and bonding pairs in water, and therefore the name of its shape and the bond angle.

Answer 8

• 4 electron pairs
• 2 bonding pairs, 2 lone pairs
• Non-linear
• 104.5°

Question 9

When drawing molecular shapes, how are multiple bonds treated? Give an example.

Answer 9

• Multiple (e.g. double/ triple) bonds are treated as a single bonding region
• Carbon dioxide has 2 double bonds, so this is counted as 2 bonding regions

Question 10

List the number of bonding regions in carbon dioxide, and therefore the name of its shape and the bond angle.

Answer 10

• 2 bonding regions, 0 lone pairs
• Linear
• 180°

Question 11

List the number of lone and bonding pairs in BF3, and therefore the name of its shape and the bond angle.

Answer 11

• 3 bonding pairs, 0 lone pairs
• Trigonal planar
• 120°

Question 12

What does it mean if a molecule is planar?

Answer 12

• All of its atoms are in the same plane

Question 13

List the number of lone and bonding pairs in SF6, and therefore the name of its shape and the bond angle.

Answer 13

• 6 bonded pairs, no lone pairs
• Octahedral
• 90°

Question 14

How is a NH4+ ion bonded together? How is this shown in a displayed formula?

Answer 14

• A molecule of ammonia and a positive hydrogen ion bond together as the hydrogen ion shares what was nitrogen’s lone pair with the atom of nitrogen, through a dative covalent bond (which behaves as a regular bonding pair)
• As the hydrogen ion was missing an electron, the overall charge of the ion is 1+

Question 15

List the number of lone and bonding pairs in NH4, and therefore the name of its shape and the bond angle.

Answer 15

• 4 bonding pairs, no lone pairs
• Tetrahedral
• 109.5°

Question 16

What is covalent bonding? Required.

Answer 16

• The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

Question 17

What is electronegativity? Required.

Answer 17

• The ability of a bonded atom to attract the bonding electrons in a covalent bond

Question 18

What is a Pauling electronegativity value?

Answer 18

• A value assigned as a measure of the relative attraction of a bonded atom for the pair of electrons in a covalent bond

Question 19

Where are electronegativity values the highest in the periodic table?

Answer 19

• Up a group
• Across (to the right of) a period
• (With the exception of group 0)

Question 20

Why do electronegativity values increase up a group?

Answer 20

• The atomic radius is smaller
• The bonded pairs of electrons are attracted more strongly to the nucleus of the atom

Question 21

Why do electronegativity values increase across a period?

Answer 21

• The nuclear charge increases
• This attracts the bonding pairs of electrons more strongly

Question 22

Which elements are the most electronegative?

Answer 22

• Fluorine
• Oxygen
• Nitrogen

Question 23

What is a non-polar bond, and what is it also known as?

Answer 23

• A bond where the electrons are shared equally between the bonded atoms
• A pure covalent bond

Question 24

What causes a bond to be non-polar? Give examples.

Answer 24

• The atoms have the same/ similar electronegativities, like carbon and hydrogen
• The atoms are the same, such as in diatomic molecules

Question 25

What is a polar bond? Give an example.

Answer 25

• A bond where the electron pair is shared unequally
• Hydrogen chloride

Question 26

What happens to atoms as a result of the existence of a polar bond?

Answer 26

• A permanent dipole is formed; there is a separation of opposite partial charges
• The pair of electrons sit closer to the atom with the higher electronegativity value, which causes the bond to be polarised as it has a small partial negative charge (δ-) on one side, and a small partial positive charge (δ+) on the other

Question 27

Why is the delta sign used to show dipoles?

Answer 27

• It shows that the charges are small

Question 28

How can electronegativity values be used to guess the type of bonding between atoms?

Answer 28

• If there is no difference in the electronegativity values of 2 bonded atoms, they will form a pure covalent bond
• A small difference creates a polar covalent bond
• If there is a large difference between the electronegativity values, one of the atoms will have a much greater attraction for the bonded pair of electrons, and would even take control of the electrons and form an ionic bond

Question 29

What makes molecules polar?

Answer 29

• If it is not symmetrical, the molecule may have an overall dipole as the dipoles wouldn’t cancel out

Question 30

How can you tell if a molecule is polar?

Answer 30

• If the dipoles act in opposing directions (think of them pointing from positive to negative), they cancel out and the molecule is non-polar
  (- If you were to pull at the dipoles around the central atom in its 3D structure, would it move? If the central atom would move, the molecule is polar.)

Question 31

Polar molecules only dissolve in what type of solvents? What about non-polar molecules?

Answer 31

• Polar solvents
• Non-polar molecules only dissolve in non-polar solvents

Question 32

How do ionic compounds (e.g. NaCl) dissolve in water?

Answer 32

• Water molecules attract and surround the Na+ and Cl- ions
• The ionic lattice breaks down as it dissolves
• Na+ ions are attracted to the oxygen of the water molecules (δ-)
• Cl- ions are attracted to the hydrogen of the water molecules (δ+)