Bonding

Question 1

What is ionic bonding

Answer 1

The transfer of electrons between metals and non-metals

Question 2

Structure of ionic compounds

Answer 2

Ionic compounds have a giant ionic lattice structure

Question 3

Properties of ionic compounds (4)

Answer 3

-high melting / boiling points
-electrical conductors when molten or dissolved
-usually soluble in water and solid at room temperature
-brittle and shatter easily when given a sharp blow

Question 4

Why do ionic compounds have high melting points

Answer 4

The giant ionic lattices are held together by many strong electrostatic forces of attraction between oppositely charged ions which require lots of energy to overcome

Question 5

Why can ionic compounds conduct electricity when molten or dissolved

Answer 5

This is because when molten or dissolved, the ions are free to move (delocalised) and can carry a charge whereas in a solid, the ions are held in fixed positions by the strong ionic bonds

Question 6

Why are ionic compounds soluble in water

Answer 6

Water molecules are polar and can therefore pull the ions away from the lattice, causing it to dissolve

Question 7

Why are ionic compounds brittle and shatter easily when given a sharp blow

Answer 7

The blow may distort the ions resulting in like charges to coming into contact with each other, causing them to repel (so the compound shatters)

Question 8

What is covalent bonding

Answer 8

The sharing of electrons between non-metals

Question 9

Structure and bonding in simple covalent compounds (small molecules/ simple molecular compounds)

Answer 9

Atoms in the molecules are held together by strong covalent bonds but the molecules are held together by weak intermolecular forces of attraction

Question 10

Properties of simple covalent compounds (3)

Answer 10

• can’t conduct electricity
• low melting point
• soluble in water

Question 11

Why can’t simple covalent compounds conduct electricity

Answer 11

There are no delocalised electrons to carry the charge as they have full outer shells

Question 12

Why do you simple covalent compounds have a low melting point

Answer 12

The many weak intermolecular forces of attraction between the molecules requires little energy to overcome

Question 13

Examples of giant covalent structures

Answer 13

-diamond
-graphite
-silicone dioxide

Question 14

Structure and bonding in graphite

Answer 14

Graphite is a giant covalent molecule. Each carbon atom forms covalent bonds with 3 other carbon atoms leaving one delocalised electron per atom. Graphite has a hexagonal structure and has layers (one layer is called graphene)

Question 15

Properties of graphite (5) and explanations

Answer 15

-soft and slippery: graphite has layers which have weak intermolecular forces of attraction between them which are easily broken allowing the layers to slide over one another
-high melting point: it’s a giant covalent structure and has many strong covalent bonds between the atoms which require a lot of energy to overcome
-can conduct electricity: has delocalised electrons which a pre free to carry the charge
- low-density: layers are far apart compared to the length of the covalent bonds
-insoluble: covalent bonds in the layers are too difficult to break

Question 16

Structure and bonding in diamond

Answer 16

Diamond has a giant covalent structure, each carbon atom forms covalent bonds with 4 other carbon atoms leaving no delocalised electrons per atom. Diamond has a tetrahedral structure

Question 17

Properties of diamond (3) and explanations

Answer 17

Hard and strong/ rigid due to tetrahedral arrangement
Can’t conduct electricity as there’s no delocalised electrons to carry the charge
High melting point as it’s a giant covalent structure, the many strong covalent bonds require a lot of energy to overcome

Question 18

Definition of dative covalent bonding

Answer 18

When an atom with a lone pair of electrons donates the lone pair to an electron deficient atom

Question 19

Define metallic bonding

Answer 19

Metals consist of a lattice of positive ions existing in a sea of delocalised electrons

Question 20

Properties of metals (4) and explanations

Answer 20

• Good electrical and thermal conductors
• strong, the more electrons (more electrostatic attraction) and the smaller the ion (nucleus closer to outer shell electron), the stronger the metal is
• malleable and ductile as metal ion is still in the same environment after a small distortion
• high melting points due to many strong electrostatic forces of attraction between positively charged ions and delocalised electrons

Question 21

Definition of electronegativity

Answer 21

The power of an atom to attract a bonding pair or electrons in a covalent bond

Question 22

Factors affecting electronegativity and reasons (3)

Answer 22

• nuclear charge (increases across a period as number of protons increases while shielding stays the same)
• electron shielding
• distance between nucleus and outer shell electrons (the smaller the atom, the closer the nucleus is to the shared pair of electrons which increases electronegativity)

Question 23

What’s the relationship between electronegativity and bond polarity

Answer 23

As electronegativity difference between atoms increases, bond polarity also increases (if atoms have the same electronegativities, the molecule will be non-polar as there’s no overall dipole moment)

Question 24

Why aren’t symmetrical molecules polar

Answer 24

Symmetrical molecules can have polar bonds but the dipoles cancel each other out as they’re in equal and opposite directions and therefore the overall molecule is non polar as there’s no net dipole moment

Question 25

Define electron pair repulsion theory

Answer 25

Each pair of electrons around an atom will repel all other electron pairs, so pairs of electrons will take up positions as far apart as possible to minimise repulsion

Question 26

Lone pairs vs bonding pairs (repulsion) and how they affect bond angles

Answer 26

Lone pairs repel more than bonding pairs (having lone pairs on an atom decreases the bond angle by 2.5 degrees)

Question 27

Shapes of molecules- linear with bond angle

Answer 27

2 bond pairs
180 degrees

Question 28

Shapes of molecules- bent and bond angle

Answer 28

2 bond pairs
2 lone pairs
104.5 degrees

Question 29

Shapes of molecules- trigonal planar and bond angle

Answer 29

3 bond pairs
120 degrees

Question 30

Shapes of molecules- trigonal pyramidal and bond angle

Answer 30

3 bond pairs
1 lone pair
107 degrees

Question 31

Shapes of molecules- tetrahedral and bond angle

Answer 31

4 bond pairs
109.5 degrees

Question 32

Shapes of molecules- square planar and bond angle

Answer 32

4 bond pairs
2 lone pairs
90 degrees

Question 33

Shapes of molecules- trigonal bipyramidal and bond angle

Answer 33

5 bond pairs
90 and 120 degrees

Question 34

Shapes of molecules- octahedral and bond angle

Answer 34

6 bond pairs
90 degrees

Question 35

What are the 3 types of intermolecular forces

Answer 35

Van deer waals
Dipole-dipole
Hydrogen bonding

Question 36

Explain how van der waals forces arise

Answer 36

AKA London dispersion forces (weakest IMF)
Uneven distribution of electrons
One side of molecule slightly more electronegative than the other which induces a temporary dipole in the neighbouring molecule
This creates a temporary van der waals force between the electronegative side of one molecule and the electropositive side of another molecule

Question 37

Explain dipole-dipole forces

Answer 37

(Only in polar molecules)
Dipole-dipole forces form between all molecules with an electronegativity difference

Question 38

Explain hydrogen bonding

Answer 38

(Strongest IMF)
occur in molecules that have a significant electronegativity difference so can only form between a hydrogen and an oxygen/nitrogen/fluorine