Topic 2 - Bonding and Structure

Question 1

What is the definition of an Orbital?

Answer 1

The area around an atom where there is a high probability of finding an electron

Question 2

What is the definition of Bonding?

Answer 2

How atoms or ions form attractive forces

Question 3

What is the definition of Structure?

Answer 3

How atoms or ions are arranged

Question 4

What is the definition of Isolelectronic?

Answer 4

When two atoms, ions or molecules have the same electronic configuration and same number of valence electrons e.g Na(+), Mg(2+) and Al(3+)

Question 5

Why do atoms form bonds?

Answer 5

In order to become more stable (they do not have to have a full outer shell of this to be the case)

Question 6

Definition of Ionic Bonding

Answer 6

The electrostatic forces of attraction between oppositely charged ions

Question 7

Definition of Covalent bonding

Answer 7

The electrostatic forces of attraction between the negative shared pair of electrons and their adjacent, positively charged nuclei

Question 8

Definition of Metallic Bonding

Answer 8

The electrostatic forces of attraction between the delocalised (sea of) electrons and the positively charged ions

Question 9

What factors affect the strength of the ionic bond

Answer 9

• The charge of the ions involved
• the radius of the ions (ionic radius)

Question 10

What is the trend in ionic radius down a group?

Answer 10

The ionic radius increases due to more filled shells of electrons meaning shielding increases so there is less attraction between the electrons and the nucleus

Question 11

Ionic Charge density formula

Answer 11

= charge/ionic radius

Question 12

Which two atoms form ions that are electron deficient?

Answer 12

Be is in group 2 and only has 4 electrons in its outer shell when it forms BeCl2
B is in group 3 and only has 6 electrons in its outer shell when it forms BF3

Question 13

Which two atoms from ions that are expanded octets and break the octet rule?

Answer 13

P has 5 electrons in its outer shell and forms PCl5 and therefore has 10 electrons in its outer shell
S has 6 electrons in its outer shell and forms SF6 and therefore has 12 electrons in its outer shell

Question 14

What is a dative covalent bond (coordinate bond)

Answer 14

Like a normal covalent bond it involves a shared pair of electrons, however, in a dative covalent bond, both of the electrons originate from one atom.

Question 15

Give an example of an ion that has a dative covalent

Answer 15

NH4 (+) when NH3 and H+ are bonded together

Question 16

Why do dative covalent bonds occur?

Answer 16

When an atom, such as nitrogen in NH3, has a non-bonded pair of electrons

Question 17

How do you show a dative covalent bond?

Answer 17

• In a dot and cross diagram they are from the same species so are drawn as the same, dots or crosses
• In a stick diagram you draw an arrow from the species with the pair of electrons to the atoms that is accepting them

Question 18

What two factors does bond length depend on?

Answer 18

• The size of the atoms
• The number of shared electron pairs (As this increases the strength of the covalent/ionic bond gets greater)

Question 19

As bond strength decreases, bond length…

Answer 19

Increases

Question 20

What is pure ionic bonding?

Answer 20

Pure ionic bonding is when there is only a very little polarising effect of the negative ion by the positive ion so the ions are perfectly spherical. The electrons are shared equally between the two atoms (however, this rarely happens)

Question 21

What does pure ionic bonding require?

Answer 21

• A large positive ion with a low charge
• A small negative ion

Question 22

What is polarisation?

Answer 22

When a cation distorts the electron cloud of an anion

Question 23

Why do certain ions have a higher polarising effect?

Answer 23

As they have a high charge density e.g Al (3+), which will attract an anion with a large ionic radius

Question 24

Why are larger negative ions more easily polarised than smaller negative ions?

Answer 24

As their outer electrons are less attracted to their nucleus and so more easily distorted

Question 25

What is the definition of electronegativity?

Answer 25

A measure of the tendency of an atom to attract a bonding pair of electrons

Question 26

Electronegativity increases as…

Answer 26

You move right and up on the periodic table

Question 27

Where does pure covalent only occur?

Answer 27

If the two atoms are identical e.g N2, O2 and H2

Question 28

How do you draw an electron density map?

Answer 28

Draw circles around the separate atoms. In ionic bonding the circles don’t overlap where as, in covalent bonding they join in the middle and are symmetrical.

Question 29

Why does melting and boiling point increase across a period e.g Na, Mg, Al

Answer 29

- The ionic charge increases and the number of delocalised electrons increases - Therefore there is more attraction between the electrons and the nucleus so the ionic radius decreases - Therefore, there is a higher charge density and so the forces of attraction within the metallic bonds are stronger so require more energy to overcome.

Question 30

Why do the melting and boiling points of group 1 metals decrease as you go down the group?

Answer 30

- ionic radius increases - the ratio of atoms to electrons in the metallic structure stays the same - Charge density decreases, therefore the forces of attraction between positive ions and electrons decrease

Question 31

What is the order of repulsion strength between pairs of electrons?

Answer 31

Lone pair - Lone pair > Lone pair - Bonding pair > Bonding pair - Bonding pair

Question 32

What is the bond angle in linear molecules? Give examples.

Answer 32

180° as they have 0lp’s and 2bp’s that repel to the point of minimum repulsion e.g BeCl2, CO2, HCN and C2H2

Question 33

What is the bond angle in Triagonal planar molecules? Give examples.

Answer 33

120° as they have 0lp’s and 3bp’s that repel to the position of minimum repulsion e.g BCl3

Question 34

What is the bond angle in Tetrahedral molecules? Give examples.

Answer 34

109.5° as they have 0lp’s and 4bp’s that repel to the position of minimum repulsion. The central atom lies at the centre of the tetrahedron. The bond coming out of the is represented by a filled triangle and the bond into the page, by a dashed triangle.

Question 35

What is the bond angle in Trigonal bipyramidal molecules? Give examples.

Answer 35

The equatorial atoms that lie on the plane have a bond angle of 120°. The axial atoms that upand down make a bond angle of 90° with the equatorial atoms. The molecule has 5bp’s and 0 lp’s.

Question 36

What is the bond angle in Octahedral Molecules? Give Examples.

Answer 36

90° as they have 6bp’s and 0lp’s that repel to the point of minimum repulsion. Two bonds go into the page and two come out of the page.

Question 37

What is the bond angle in Triangular Pyramidal Molecules? Give Examples.

Answer 37

107° as they have 1lp and 3bp’s which repel to the position of minimum repulsion. One bond is on the plane, one into the page and one out of the page e.g NH3

Question 38

What is the bond angle in Non -linear molecules? Give examples.

Answer 38

104.5° as they have 2bp’s and 2lp’s that repel to the position of minimum repulsion. Both bonds are on the plane e.g H2O

Question 39

How can molecules that are electron deficient gain a full outer shell?

Answer 39

By sharing a pair of electrons with an atom which has a non-bonding or lone pair of electrons

Question 40

What is the bond angle in Square Planar Molecules? Give examples.

Answer 40

90° as they have 4bp’s and 2lp’s. Two bonds go into the plane and two bonds go out of the plane e.g XeF4

Question 41

Why are all the bond angles in Methane 109.5°?

Answer 41

- Methane contains 4 bonding pairs and 0 lone pairs of electrons. - Electron pairs repel to the position of minimum repulsion - The 4 bonding pairs of electrons will repel each other equally - Giving methane it’s tetrahedral shape with bond angle 109.5°

Question 42

What is the trend in ionic radius across period 3, from Na to Cl?

Answer 42

As you go from Na (+) → Al (3+) the ions begin already small and get smaller. As you go from P (3-) → Cl (-) the ions begin big and get smaller

Question 43

Why do the negative ions all have a greater ionic radius than the positive ones?

Answer 43

This is because the negative ions are gaining electrons, therefore they are gaining a whole extra layer making their radius larger.

Question 44

Why do the metallic ions decrease in radius from group 1 to group 3?

Answer 44

Whilst the ions are isoelectronic, as you go across a period the number of protons in the nucleus is increasing giving it a more positive charge. This in turn attracts the electrons more strongly causing a decrease in the ionic radius.

Question 45

Why do the non-metallic elements decrease in radius from group 5 to group 7?

Answer 45

The number of electrons is decreasing so the overall charge on the ion gets smaller. Therefore, there is less repulsion between the electrons so a smaller ionic radius.

Question 46

What is the definition of ionic radius?

Answer 46

Ionic radius is the distance from the nucleus to the outer edge of the electron cloud of an ion.

Question 47

What is the definition of ionic radius?

Answer 47

Ionic radius is the distance from the nucleus to the outer edge of the electron cloud of an ion.

Question 48

What is the definition of a cation?

Answer 48

A positively charged ion

Question 49

What is the definition of an anion?

Answer 49

A negatively charged ion

Question 50

When will a cation form a stronger ionic bond with an example?

Answer 50

When the ion is smaller with a higher charge the ionic bonds are stronger. for example, the size of the ions decreases in radius and increases in charge as you go from Na(+) to Al(3+)

Question 51

When will an anion form a stronger ionic bond?

Answer 51

When the ion has a larger ionic radius the electrons are more easily attracted by the positive nucleus of the other ion. For example, as you go from F(-) to O(2-) the strength of the ionic bond that forms increases.

Question 52

How do you draw a dot and cross diagram for the compound ion sulphate?

Answer 52

- The bottom 2 oxygen atoms are bonded to the sulphur by a double bond - The top 2 oxygen atoms are bonded to the sulphur by a single bond - The extra 2 electrons gained to form the 2- ion are shown by an asterisk - The dot and cross diagram is surrounded by brackets and a 2- charge in the top right

Question 53

How do you draw a dot and cross diagram for the compound ion hydroxide?

Answer 53

- The dot and cross diagram is surrounded by brackets and a - charge in the top right - The extra electron gained to form the - ion are shown by an asterisk

Question 54

How do you draw a dot and cross diagram for the compound ion carbonate?

Answer 54

- The top 2 oxygen atoms are bonded to the carbon by a single bond - The extra 2 electrons gained to form the 2- ion are shown by an asterisk - The bottom oxygen atom is bonded to the carbon by a double bond

Question 55

Why does H+ form dative covalent bonds?

Answer 55

As it only has 1 proton and no electrons, therefore needing 2 electrons for a full outer shell

Question 56

Why does the bond strength decrease from HF to HI, using the idea of larger atoms to explain your answer?

Answer 56

As you move from HF to HI the atomic radius of the halogen increases as it gains more shells. There is also increased shielding between the shared pair of electrons and the adjacent halogen nucleus. Therefore the forces of attraction between the nucleus and the electron cloud decrease.

Question 57

Why is C-C bond weaker than C=C bond?

Answer 57

This is because the C-C bond only has one shared pair of electron and has a larger bond length in comparison to C=C which has two shared pairs of electrons and a smaller bond length, so the electrostatic forces of attraction between the electrons and nuclei are stronger

Question 58

What does the statement ‘the ionic compound has covalent character’ mean?

Answer 58

When a cation distorts an anion there is an increase in electron density in the space between the ions, some sharing of electrons and partial covalency

Question 59

Explain why negative ions are polarised by positive ions?

Answer 59

Positive cations attract the negative electron cloud of the anion

Question 60

Explain why smaller, more highly charged positive ions have a greater polarising effect?

Answer 60

If the positive ion is smaller with a large charge they have a greater charge density so a greater attractive force on the electron cloud

Question 61

What happens to the shared pair of electrons when atoms of different electronegativities bond covalently?

Answer 61

The shared pair of electrons is attracted to the more electronegative element. This causes the bond to become polarised and have a delta(+) and delta(-) charge

Question 62

What is the definition of a dipole?

Answer 62

A bond or molecule who’s ends have opposite charges/ a separation of positive and negative charge between two covalently bonded atoms.

Question 63

Explain the change in bond angle from 109.5° to 107° to 104.5° in methane, ammonia and water.

Answer 63

- There is an increasing number of lone pairs - lone pair-lone pair repulsion is greater than lone pair-bonding pair repulsion - Therefore the bond angle is decreasing

Question 64

What is a polar bond?

Answer 64

A polar bond forms when the Electronegativities of the two atoms in the bond are different so a partial positive and a partial negative charge is gained

Question 65

Can a molecule have polar bonds but not be polar? If yes, why?

Answer 65

Yes, this is because if the polar bonds are symmetrical and act in opposite directions then the charges cancel out.

Question 66

How do you draw a dipole onto a diagram of the shape of molecule?

Answer 66

An arrow with a base from the partially positive to the partially negative atom

Question 67

What are London Forces?

Answer 67

They are the weakest type of intermolecular force that occurs between all molecules

Question 68

How do London Forces form?

Answer 68

- In all molecules the electrons are moving constantly and randomly - This leads to fluctuations in the electron density in side the molecule - This then can cause parts of the molecule to become more or less negative to from an instantaneous/temporary dipole - These temporary dipoles can cause dipoles in neighbouring molecules called induced dipoles - The London Force is the force between the two dipoles

Question 69

What are the two factors that effect the strength of London Forces?

Answer 69

- The number of electrons in a molecule - The shape of a molecule

Question 70

What is the definition of polarisability?

Answer 70

Polarisability is an indication of the extent to which the electron cloud in a molecule (ion) can be distorted by a nearby electric charge.

Question 71

How does an increasing number of electrons result in stronger London Forces?

Answer 71

Bigger molecules that have a large number of electrons have a higher polarisability. Therefore, the possibility for temporary induced dipoles is greater so there are stronger London Forces present.

Question 72

Why do the melting and boiling points of halogens increase down the group?

Answer 72

- As you go down the group the halogen molecules increase in size, therefore their are more electrons contributing to a larger electron cloud - The outer electrons become further from the nucleus so the electron cloud is more easily distorted forming more instantaneous dipoles - This leads to stronger London Forces being formed so more energy is required to overcome them

Question 73

Why do straight chain molecules have high melting and boiling points in comparison to branched chain molecules?

Answer 73

Straight chain molecules have a larger surface area in comparison to branched molecules which are more compact. Therefore, the straight chain molecules have a larger area over which London Forces can form so more energy is required to overcome them.

Question 74

Do London Forces act between all molecules?

Answer 74

Yes

Question 75

How do permanent dipole-dipole forces form?

Answer 75

These occur when polar molecules, ones that have permanent dipoles, attract each other. This is a more strong intermolecular force in comparison to London Forces.

Question 76

What are hydrogen bonds?

Answer 76

A strong type of permanent dipole-dipole interaction and the strongest type of intermolecular force.

Question 77

What 2 features must be present for hydrogen bonds to form between molecules?

Answer 77

- A hydrogen atom must be covalently bonded to a small, highly electronegative atom that is N, O or F - There must be a lone pair on a second electronegative atom in a neighbouring molecule

Question 78

What are the 2 reasons for why Hydrogen Bonding is particularly strong in comparison to other permanent dipole-dipole forces?

Answer 78

- There is a very large dipole, due to the difference in electronegativities, between the Hydrogen atom and N, O or F atom. This means that the covalent bond is extremely polarised. - The small size of the Hydrogen atom gives a high charge density from the partial positive charge that it has

Question 79

Why is the bond angle around the hydrogen, that is both covalently bonded to an atom and has a hydrogen bond with another molecule, 180°?

Answer 79

There are 2 pairs of electrons around the Hydrogen atom that are involved in the Hydrogen bond. These pairs of electrons repel to the position of minimum repulsion, as far apart as possible

Question 80

What are the elements you need to remember when drawing a diagram of hydrogen bonding?

Answer 80

- Partially positive and negative charges on the H atom and O/N/F atom - Lone pair(s) of electrons on the atom involved in the hydrogen bond - Bond angle of 180° around the hydrogen involved in the Hydrogen bond - Dashed lines to represent the Hydrogen bond

Question 81

What gives ammonia, water and hydrogen fluoride their relatively high boiling temperatures?

Answer 81

Hydrogen Bonding

Question 82

Why does the density of water increase as it goes from a solid to a liquid?

Answer 82

Ice has a lower density than water. This is because, in a solid state, the Hydrogen bonding holds the water molecules in an open structure. As the ice melts this structure collapses, therefore the molecules get closer together causing the density to increase.

Question 83

Why do the boiling points of the noble gases increase as you go down the group?

Answer 83

As you go down the group the number of electrons increases, therefore there is a large electron cloud that distorts more easily to form instantaneous dipoles. Therefore, stronger London Forces are created between molecules so the energy required to break these forces increases.

Question 84

Why does hydrogen fluoride have a higher boiling point than the other hydrogen halides?

Answer 84

As it is the only hydrogen Halide containing Hydrogen bonds

Question 85

Why does the boiling point increase gradually as you move from HCl to HI?

Answer 85

Whilst the difference in electronegativity decreases so the strength of the permanent dipole-dipole forces decreases, this is smaller than the increase in the strength of the London Forces due to an increased number of electrons. Therefore, more energy is required to overcome these attractive forces.

Question 86

Why are the boiling points of all the hydrogen halides much higher than the Noble Gases?

Answer 86

- Noble gases only contain London Forces - Hydrogen halides contain both London Forces and permanent dipole-dipole forces or hydrogen bonds. These types of intermolecular force and significantly stronger than London Forces so the addition of them means more energy is required to overcome the forces between molecules.

Question 87

When describing differences in boiling points, what points should you make?

Answer 87

London forces - affected by number of electrons and the shape of the molecule (branched or straight chain) Permanent dipole-dipole - If the molecule is permanently polar (differing electronegativities) Hydrogen bonds - strongest off all intermolecular forces Some molecules will contain a more than one type of intermolecular force

Question 88

What is the rule for the solubility of solutes?

Answer 88

“Like dissolves Like” This means that polar solvents dissolve polar solutes and vice versa

Question 89

Why are many ionic solids able to dissolve in water?

Answer 89

- The ions in ionic solids are strongly hydrated by polar water molecules - The water molecules cluster around the ions and bind to them - In ionic solids, the energy released when the water molecules bind to the ions is enough to overcome the ionic bonding between the ions

Question 90

Why are other salts, not ionic solid, insoluble in water?

Answer 90

As the energy that is released when the water molecules attach to the ion, is not large enough to overcome the existing forces between the ions.