2. Bonding

Question 1

How are ionic compounds formed?

Answer 1

Transfer of electrons from metal atoms to non metal atoms

Question 2

Define an ion

Answer 2

Atoms which have lost or gained electrons to gain an overall charge

Question 3

How do metals form ions

Answer 3

Lose electrons to become positively charged

Question 4

How do non-metals form ions

Answer 4

Gain electrons to become negatively charged

Question 5

Define a cation

Answer 5

Positive ions

Question 6

Define an anion

Answer 6

Negative ions

Question 7

Define an ionic bond

Answer 7

The electrostatic attractive between oppositely charged ions

Question 8

Why do atoms lose/gain electrons?

Answer 8

To get a noble gas configuration

Question 9

Define ionic lattice

Answer 9

The structure where ions are held in a 3D framework

Question 10

Define lattice

Answer 10

Regular repeated 3D arrangement of ions, atoms or molecules in a solid

Question 11

What causes the crystalline nature of ionic compounds?

Answer 11

The regular pattern of the ions

Question 12

Define decrepitation

Answer 12

The cracking noise that ionic compounds make when they’re heated caused by the crystalline structure

Question 13

Why do ionic compounds have high melting/boiling points?

Answer 13

Because lots of energy is required to overcome the strong electrostatic forces of attraction between oppositely charged ions

Question 14

What factors affect the strength of the ionic bond?

Answer 14

The size of the ion and the size of the charge

Question 15

Why are +ve generally smaller?

Answer 15

Bc they’ve lost ions so fewer shells and increased effective nuclear charge so shells drawn in closer

Question 16

Why are -ve ions generally larger?

Answer 16

Gain electrons so more repulsion of electrons pushes them further apart
Smaller effective nuclear charge bc more electrons with same no of protons

Question 17

Are ionic compounds soluble?

Answer 17

Most of the time, depending on the strength of the ionic bonds

Question 18

Why can ionic compounds only conduct electricity in the molten/aqueous state?

Answer 18

Because the ions are free to move and carry charge in these states

Question 19

Define a covalent bond

Answer 19

One or more shared pairs of electrons between two non-metal atoms

Question 20

Define a single covalent bond

Answer 20

One shared pair of electrons

Question 21

Define a double covalent bond

Answer 21

Two pairs of shared electrons

Question 22

Define a triple covalent bond

Answer 22

Three pairs of shared electrons

Question 23

What is a bonding pair of electrons?

Answer 23

Pair of electrons that are shared between two atoms

Question 24

What is a lone pair of electrons?

Answer 24

An unshared pair of electrons that are not involved in bonding

Question 25

Define a coordinate bond

Answer 25

A bond which has a shared pair of electrons where both electrons are supplied by one atoms

Question 26

How are ions in metals arranged?

Answer 26

In layers with the outer shell electrons not bound to an atom

Question 27

Define delocalised electrons

Answer 27

Electrons that are not bound to one particular atom

Question 28

What is a metallic bond?

Answer 28

The electrostatic attraction between delocalised electrons and positive ions in the lattice

Question 29

Why can metals conduct electricity?

Answer 29

Delocalised electrons are free to move throughout the structure and can carry charge

Question 30

Why can metals conduct heat?

Answer 30

Delocalised electrons can move and carry heat energy throughout the metal

Question 31

Why are metals malleable and ductile/

Answer 31

Because the layers can slide over each other without disrupting the bonding

Question 32

Why do metals have high densities?

Answer 32

Because the positive ions are packed tightly together

Question 33

Why do metals have high melting points?

Answer 33

Hella energy required to overcome the strong forces of attraction between the positive ions and delocalised electrons

Question 34

Why do some metals have a higher melting point than others?

Answer 34

Because they have more delocalised electrons and they may have smaller ions so they're packed more tightly together

Question 35

What are molecular covalent substances?

Answer 35

Substances which exist as single molecules

Question 36

What are molecular covalent crystals?

Answer 36

Molecular covalent substances which form crystalline structures

Question 37

Why are molecular covalent crystals brittle?

Answer 37

Because they dont have strong bonds holding them together

Question 38

Why do molecular covalent crystals not conduct electricity?

Answer 38

Bc they don't have any charged particles to carry charge

Question 39

Define an allotrope

Answer 39

Different forms of the same element in the same physical state

Question 40

What are the two allotropes of carbon?

Answer 40

Diamond and graphite

Question 41

Why is diamond hard?

Answer 41

Many strong covalent bonds which require hella energy to overcome

Question 42

Why does diamond have a high melting point?

Answer 42

Strong covalent bonds that require hella energy to overcome

Question 43

In diamond how many bonds does each carbon form?

Answer 43

4

Question 44

Why does diamond not conduct electricity?

Answer 44

No charged particles which can move and carry charge

Question 45

In graphite how many bonds does each carbon form?

Answer 45

Three

Question 46

Why can graphite conduct electricity?

Answer 46

Delocalised electrons which can move and carry charge

Question 47

Why does graphite have high melting point

Answer 47

Strong covalent bonds which require a lot of energy to break

Question 48

Why is graphite a good lubricant?

Answer 48

Layers can slide over each other because of weak forces in between layers

Question 49

What are the types of crystalline structures?

Answer 49

Ionic, metallic, macromolecular and molecular

Question 50

Why do lone pairs have a greater repulsive force?

Answer 50

Because they're closer to the central atom

Question 51

Linear: bonding pairs and angles

Answer 51

2 bonding pairs with 180 degrees between them

Question 52

Trigonal planar: bonding pairs and angles

Answer 52

3 bonding pairs with 120

Question 53

Tetrahedral: bonding pairs and angles

Answer 53

4 pairs with 109.5

Question 54

Trigonal bipyramidal: bonding pairs and angles

Answer 54

5 pairs with 90 and 120

Question 55

Octahedral: bonding pairs and angles

Answer 55

6 bonding pairs all 90

Question 56

Pyramidal: bonding pairs and angles

Answer 56

3 bonding pairs, one lone pair - 107

Question 57

Bent: bonding pairs and angles

Answer 57

2 bonding pairs, 2 lone pairs - 104.5

Question 58

Define electronegativity

Answer 58

Power of an atom to attract the pair of electrons in a covalent bond

Question 59

What factors affect electronegativity?

Answer 59

``` Atomic radius (distance between pair of electrons and nucleus) Nuclear charge Shielding ```

Question 60

Why does electronegativity increase across a period?

Answer 60

Decrease in atomic radius so stronger attraction between nucleus and bonding pair Increase in nuclear charge so more attraction

Question 61

Why does electronegativity decrease down a group?

Answer 61

Increased atomic radius and weaker nuclear charge b more shielding so weaker attraction between nucleus and bonding pairs

Question 62

How are induced dipoles formed?

Answer 62

Electrons are in constant motion so they may be distributed more on one side of the molecule creating a temp dipole

Question 63

What is the relationship between van der Waals forces and the size of the molecule?

Answer 63

Larger molecules have more electrons so greater induced dipoles so greater van der Waals forces between molecules

Question 64

How are hydrogen bonds formed?

Answer 64

Occur between delta positive charge of a H in a bond and the lone pair of electrons of an O, N or F of another molecule or atom