3 Chemical Bonding

Question 1

Define (and describe) an ionic bond.

Answer 1

An ionic bond is an electrostatic attraction between cations and anions in an ionic lattice (They are non-directional as an ion attracts an oppositely charged ion in all directions without a preferred orientation).

Question 2

What indicates the strength of an ionic bond?

Answer 2

The greater the magnitude of lattice energy, the stronger the ionic bond. Lattice energy is the energy released when one mole of an ionic crystalline solid is formed from its constituent gaseous ions. It is proportional to the mod of q+q-/r+ + r- (where q represents charge and r represents the radius of the cation and anion. r+ + r- is also the inter-ionic distance).

Question 3

Explain how charges and inter-ionic distance affect the strength of an ionic bond.

Answer 3

The higher the charge, the greater the electrostatic attraction and strength of ionic bond. The shorter the inter-ionic distance between the ions, the greater the attraction, the stronger the ionic bond.

Question 4

Describe how ions are arranged in the giant ionic lattice structure.

Answer 4

In an ionic compound, the constituent ions are held in fixed positions in an orderly arrangement such that the attraction between the oppositely charged ions is a maximum and repulsion between similarly charged ions is a minimum.

Question 5

Describe how atoms are arranged in a giant molecular lattice structure.

Answer 5

The atoms are held together in an extensive network by covalent bonds (also known as a macromolecule).

• Recall sec school knowledge.

Question 6

Describe a simple molecular lattice structure.

Answer 6

Made up of molecules attracted to each other by weak intermolecular forces of attraction.

Question 7

State the physical properties of substances with simple molecular lattice structure.

Answer 7

(i) Low m.p. - liquids or gases at room temperature
(ii) Soluble in non-polar solvents (both id-id interaction) and insoluble in polar solvents
(iii) Do not conduct electricity in solid, molten state (no mobile charge carriers) but it may ionise in aqueous state to form ions

Question 8

Define (and describe) a metallic bond and describe what metals are made of.

Answer 8

A metallic bond is the electrostatic attraction between a lattice of positive ions and delocalised electrons. They are non-directional. Metals are considered to be composed of a rigid lattice of positive ions surrounded by a sea of delocalised electrons.

Question 9

State and explain the factors affecting the strength of metallic bonds.

Answer 9

(i) Number of valence electrons available (more, stronger)
(ii) Charge of cations (higher, stronger)
(iii) Size of cations (smaller, stronger due to greater charge density and thus greater electrostatic attraction between the lattice of positive ions and delocalised electrons)

Question 10

State the physical properties of metals.

Answer 10

(i) High electrical conductivity(delocalised electrons act as charge carriers and flow towards positive terminal)
(ii) Good thermal conductivity (Electrons take in thermal energy and move faster and more randomly to collide with other electrons, passing on the energy faster)
(iii) Malleable (beaten into shapes due to non-directional bonds) and ductile (drawn into wires)
(iv) High density (closely packed)
(v) High m.p. and b.p.
(vi) Insoluble in all solvents except liquid metals

Question 11

State the octet rule.

Answer 11

Atoms tend to lose, gain or share electrons until they have 8 electrons in their valence shell.

Question 12

What are the steps to drawing the Dot and Cross diagrams for ionic compounds?

Answer 12

Pg 10 of notes

Question 13

Define a covalent bond.

Answer 13

A covalent bond is the electrostatic attraction between the shared pair of electrons and the positive charged nuclei (represented by a straight line).

Question 14

Define a dative covalent bond. What are the requirements to form a dative covalent bond?

Answer 14

A dative covalent bond is formed when the shared pair of electrons is provided by only one of the bonding atoms (represented by an arrow). To form such a bond, one atom must have a lone pair of electrons for donation while another atom or ion must have a vacant, low-lying orbital to accept the pair of electrons (Vacant means empty, unoccupied by electrons, Low-lying refers to being energetically accessible by that element). This can result in the formation of an adduct or dimer.

Question 15

In terms of n, what is the noble gas configuration?

Answer 15

ns2np6

Question 16

What are the steps to drawing the Dot and Cross diagrams for species with covalent bonds?

Answer 16

Pg 13 of notes

Question 17

What are the steps to drawing the Dot and Cross diagrams for charged species?

Answer 17

Electrons are generally lost from the less electronegative atom (usually the central atom) and electrons are generally gained by the more electronegative atom (usually the side atom). Overall charge must be included.

Question 18

What are the requirements for an element to expand octet? Explain in terms of bonding and structure why FO2 does not exist but ClO2 does.

Answer 18

Molecules with more than 8 electrons in the expanded octet of the valence shell of an atom can do so when the atom is in period 3 and beyond.

Why O=F=O cannot exist: The structure for ClO2 involves an expansion of octet as the Cl atom forms double bonds with
the each of the O atoms. Cl has low-lying vacant 3d orbitals which can accommodate the
additional electrons from bonding.
F as a Period 2 element does not have low-lying, vacant orbitals which can be utilised for the
expansion of octet. Hence FO2 does not exist.

Why OO cannot exist: In addition, the formation of OF→O by which F forms two dative covalent bonds with O is not
possible as F is too electronegative to donate its lone pairs to O to form the dative bonds.

Question 19

Why is it that PCl5 exists and not NCl5?

Answer 19

3d orbitals available for expansion of octet unlike elements in period 2 which can only accomadate a maximum of 8 electrons.

Question 20

What are vacant, low-lying orbitals?

Answer 20

Vacant, low-lying orbitals are either in the same subshell or same shell as occupied valence orbitals.

Question 21

What is the VSEPR theory?

Answer 21

Valence shell electron pair repulsion theory. It is used to predict the molecular geometry of a species based on the theory that electron pairs (bond pairs and lone pairs) around the central atom of the molecules are arranged as far apart as possible in space so as to minimise their mutual repulsion.

Question 22

Rank the different types of electron pairs in increasing order in terms of their repulsion. Explain how bond angles deviate from the predicted bond angle as the number of lone pairs of electrons increases for the same electron pair geometry.

Answer 22

Bond pair-bond pair repulsion, lone pair-bond pair repulsion and lone pair-lone pair repulsion. As the number of lone pairs of electrons increases for the same electron pair geometry, the bond angles deviate more greatly from the predicted angle. This is because the closer the electron-pairs are to the central atom, the greater is the repulsion. About the same central atom, a lone pair exerts greater repulsion than a bond pair. The lone pair is attracted by only one positive nucleus and hence is closer to the central atom compared to the bond-pair electrons which are attracted by two nuclei.

Question 23

State the steps required to deduce the molecular shape of a given species.

Answer 23

Pg 16 of notes

Question 24

Recall the table of VSEPR theory.

Answer 24

Pg 17 of notes

Electron pair geometry, Shape and Bond Angle

Question 25

Define electronegativity.

Answer 25

Electronegativity is the relative ability of an atom in a molecule to attract the shared/bonding electrons in a covalent bond.

Question 26

Explain why repulsion between bond pairs increases when there is an increase in electronegativity of an element.

Answer 26

When the element is more electronegative, the element draws its bond-pair closer to itself and the bond-pair electrons are thus closer to the nucleus of the element and exert more repulsion (greater bond angle) than those in a less electronegative element.

Question 27

Explain how a non-polar covalent bond is formed. Describe how the electron-pair/bonding electrons are shared between two atoms.

Answer 27

When two atoms of the same element and hence, the same electronegativity form a covalent bond, the electron-pair (bonding electrons) are equally shared between the two nuclei and the bond is a non-polar covalent bond. The electron-pair is equidistant from the two atoms and there is an even distribution of electrons/electrons that are symmetrically arranged.

Question 28

Explain how a polar covalent bond is formed.

Answer 28

When two atoms of different elements and hence, the different electronegativities form a covalent bond, the electron-pair (bonding electrons) are unequally shared between the two nuclei. Partial charges arise on the two bonded atoms, forming a polar covalent bond. The more electronegative atom has a stronger tendency to attract the electrons and hence the shared electron-pair is closer to the more electronegative atom. The electron pair is not equidistant from the two atoms and there is an uneven distribution of electrons.

Question 29

Is there a relationship between electronegativity, bond dipole moment and polarity of a bond?

Answer 29

Yes, the degree of polarity of a bond is measured by its dipole moment (which is a vectorial quantity and is shown by an arrow that points towards the more electronegative element). The greater the difference in electronegativity between the two atoms, the greater the bond dipole moment and the more polar the bond.

If you have a molecule with a central atom bonded to all identical atoms and no lone pairs it is definitely non-polar. However, just because it does not fall into this category does not necessarily make a molecule polar –there is still a need to draw out the molecule and check for dipole moment(taking into account its shape)

Question 30

What is a polar molecule?

Answer 30

A polar molecule has a distinct partially positive and partially negative region in its structure, resulting in a permanent dipole (an overall dipole moment that is not zero).

Question 31

How do we determine if a molecule is polar?

Answer 31

To determine whether the molecule is polar or has a net dipole moment, we must first look at the polarity of the bonds and the molecular geometry to find the vector sum of all the bond dipole moments.

Question 32

What is a non-polar molecule?

Answer 32

A molecule is non-polar when both the atoms are made of similar electronegativities since the bond dipole moments and vector sum of dipole moments is zero.

Question 33

Describe and explain how polar molecules will behave when subjected to an electric field. Give an example of such instances.

Answer 33

The polar molecule will orientate in a direction opposite to that of the field to minimise the electrostatic energy of the molecules. For example, the molecules of a polar compound become aligned between the plates of a capacitor with the more electronegative atom being aligned with the positive plate. Stream of liquid being deflected towards the charged rod.

Question 34

State the three types of intermolecular forces in increasing order in terms of their strength.

Answer 34

1) Instantaneous dipole-induced dipole forces
2) Permanent dipole-permanent dipole forces
3) Hydrogen bonds

Question 35

Explain how instantaneous dipole-induced dipole forces arise. Where can they be found?

Answer 35

Electrons are constantly moving around in a particle. At any given moment, the electron density of a particle can be unsymmetrical, resulting in an instantaneous dipole which can induce a short-lived dipole in a neighbouring molecule, resulting in instantaneous dipole-induced dipole interactions.

Since instantaneous dipole-induced dipole interactions are caused by momentary movements of electron charge in atoms, they are present between all particles.

Question 36

Why are instantaneous dipole-induced dipole forces short-lived and overall attraction weak?

Answer 36

Such interactions are short-lived and the overall attraction is weak because the electrons keep moving and the dipoles vanish and reform.

Question 37

State and explain the factors that affect the strength of instantaneous dipole-induced dipole interactions.

Answer 37

(i) Number of electrons/electron cloud size: As the number of electrons in a molecule increases, the strength of the interaction increases. Larger electron clouds are more easily polarised/distorted than smaller electron clouds, resulting in greater ease of formation of instantaneous dipoles and induced dipoles.
(ii) Surface area for molecular interaction: Straight-chained hydrocarbons have greater surface area for intermolecular interactions compared to branched isomers and so they have stronger interactions.

Question 38

Explain the trend of boiling points of halogens down the group.

Answer 38

As the number of electrons of the halogens increases down the group, the ease of polarisability of electron clouds and strength of id-id interaction also increases down the group. Since more energy will be required to break the stronger intermolecular interactions, the boiling points of halogens increase down the group

Question 39

Explain how permanent dipole-permanent dipole interactions arise. Where can they be found?

Answer 39

(Polar molecules have permanent dipoles in their structures, these molecules tend to align such that the partially positive end of one molecule is near the partially negative end of the other molecule.) The electrostatic attraction between the partially positive and partially negative end of the molecules gives rise to permanent dipole-permanent dipole interactions.

Question 40

State and explain the factor(s) that affect the strength of permanent dipole-permanent dipole interactions.

Answer 40

The strength of permanent dipole-permanent dipole interactions increases with increasing dipole moments. (e.g. The difference in electronegativity between H and Fin the H-F bond is greater than the difference in electronegativity between H and Clin the H-Cl bond, so the H-F bond is more polar and thus has a greater dipole moment.) For molecules with the same number of electrons, the strength of their instantaneous dipole-induced dipole interactions is similar. However, the polar molecule will form stronger permanent dipole-permanent dipole interactions than non-polar molecules which form only instantaneous dipole-induced dipole interactions.

Question 41

What is a hydrogen bond? What are the requirements for a hydrogen bond to be formed? Explain how hydrogen bonds arise.

Answer 41

A hydrogen bond is a type of permanent dipole-permanent dipole interactions that occurs between molecules containing a hydrogen atom covalently bonded to a very small, highly electronegative atom with lone electron pairs, principally F, O and N.

There must be a hydrogen atom (DIRECTLY) covalently bonded to an F, O or N atom and a lone pair of electrons on the F, O, N atom in a neighbouring molecule (bearing a partially negative charge which can attract the partially positive charge on the H atom).

Hydrogen bonds arise when the highly electronegative atoms F, O and N cause the hydrogen atom to have a high partially positive charge which allows it to form a particularly strong attraction with a lone pair of electrons on an adjacent molecule, creating an intermolecular force known as the hydrogen bond. The small sizes of the F, O and N atoms also allow the lone pair on the other F, O, N atom to come close to the protonic H atom.

Question 42

How do we determine the extensiveness of hydrogen bonding in a substance? How does it affect the boiling point of substances?

Answer 42

To calculate the average number of H-bonds per molecule, find the number of lone pairs on F, O, N and the number of hydrogen atoms attached to F, O, N. The average number of H-bonds per molecule is the smaller of the two numbers. Greater extensiveness, high boiling point as more energy is required to break more hydrogen bonds.

Question 43

Explain how some hydrogen bonds can be stronger than others. Give an example.

Answer 43

The strength of a hydrogen bond varies with the dipole moments of bonds. The greater the difference in electronegativity, the higher the dipole moment and the stronger the hydrogen bond. The hydrogen bonds between water molecules are stronger than that in ammonia molecules due to the higher dipole moment in the H-O bond than the N-H bond. The dipole moment of the H-F bond is also greater than the H-N bond as F is more electronegative than N. Thus, hydrogen bonds between HF molecules are stronger than those between NH3 molecules and hence a greater amount of energy is required to overcome the hydrogen bonds between HF molecules than the ammonia molecules. Thus, HF has a higher boiling point than NH3.

Question 44

Explain why ice is less dense than water.

Answer 44

Pg 28 and 29 of notes.

Question 45

Explain why the molar mass of some carboxylic acids in the vapour phases is twice that of its molar mass calculated from its molecular formula. Explain why carboxylic acids do not form dimers in aqueous solutions.

Answer 45

Carboxylic acids exist as dimers in the vapour phase. It is made up of two carboxylic acids bonded to each other by two hydrogen bonds forming an 8-membered ring. However, in aqueous solutions, the water molecules are in abundance and the carboxylic acids will form hydrogen bonds with water instead of with one another.

Question 46

Explain the differences in boiling point and solubility of 2-nitrophenol and 4-nitrophenol.

Answer 46

Pg 30 of notes.

Question 47

State the factors that affect whether a substance dissolves.

Answer 47

(i) Amount of energy is needed to overcome the intermolecular forces within the solute and within the solvent
(ii) Amount of energy released to, compensate for the above energy when intermolecular forces of attraction are formed between the solvent and solute molecules in the solution.

Question 48

What are the three possible cases of dissolution?

Answer 48

Case 1: Simple molecules with the same type of intermolecular forces mix well
Case 2: Ionic solids tend to dissolve in water
Case 3: A chemical reaction occurs (Pg 33 of notes)

Question 49

How do we determine whether dissolution is favourable/unfavourable?

Answer 49

1) Identify the solute-solute interaction
2) Identify the solvent-solvent interaction
3) Identify the type of solute-solvent interaction possible

Dissolution will be favourable when the solute-solvent interaction is similar to the solute-solute and solvent-solvent interaction. This is because there is sufficient energy to overcome the solute-solute and solvent-solvent interactions.

Dissolution will not be favourable when the solute-solvent interactions are weaker than the solute-solute and/or solvent-solvent interactions. This is because the energy released in forming the solute-solvent interactions cannot compensate for the breaking of the e.g hydrogen bonds in water (solvent-solvent interactions in this case).

Question 50

What are ion-dipole forces? Explain why ionic solids tend to dissolve in water/polar solvents.

Answer 50

An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a polar molecule (a neutral molecule that has a dipole). The anion attracts the partially positive end of a neutral polar molecule.

Ionic solids tend to dissolve in water/polar solvents because of the large amount of energy released from ions forming strong ion-dipole interactions with the polar water molecules (i.e. strong ion-dipole interactions). This can compensate for the energy required to overcome the strong ionic bonds in the crystal lattice which must be broken down in order to dissolve.

Question 51

What factors affect the strength of ion-dipole forces?

Answer 51

1) Charge and size of an ion (greater charge, smaller size, greater attraction for polar molecule)
2) Magnitude of the polar moment and size of the polar molecule

Question 52

When does an ionic solid not dissolve in water?

Answer 52

When the ionic bonds in the lattice are very strong and require a lot of energy to overcome such as in MgO.

Question 53

Explain when covalent bonding between 2 atoms occurs.

Answer 53

When the valence orbitals of the two atoms overlap and the two orbitals are either 1) both singly occupied with an electron (covalent bond) or 2) one is filled with a pair of electrons while the other is an empty orbital (dative covalent bond/co-ordinate bond).

Question 54

Explain the variation of potential energy with the inter-nuclear distance between two hydrogen atoms. Explain how and when a covalent bond is formed with regard to potential energy.

Answer 54

When the inter-nuclear distance between the two hydrogen atoms is small, the potential energy is high due to inter-nuclear repulsion as the atoms are too close to one another. As the distance is too far apart, the atoms are separated and not even sharing electrons.

A covalent bond is formed when the molecule is most stable when its potential energy is at a minimum. Here, the two atoms approach each other and their orbitals overlap and the electron pair spreads out over both orbitals, forming a covalent bond. The overlap of the atomic orbitals is effective as the electron density in the region between the nuclei is a maximum.

Question 55

Describe the two ways covalent bonds are classified. Which type of bond is stronger?

Answer 55

1) A sigma bond is formed when two orbitals overlap head-on. The electron density is concentrated between the nuclei of the two bonding atoms, along the nuclear axis.
2) A pi bond is formed when atomic orbitals overlap collaterally (side-to-side overlap) The orbitals have electron clouds above and below the nuclear axis but with zero electron density along the axis. This type of bonding takes place when atoms undergo multiple bonding where one of them must be a sigma bond, while the rest are pi bonds.

Sigma bond is stronger than a pi bond as the orbital overlap in sigma bond is more effective.

Question 56

Define covalent bond length. Explain how it affects the strength of a covalent bond.

Answer 56

Covalent bond length is the distance between the nuclei of two bonded atoms. The shorter the bond length, the stronger the covalent bond. This is because the length of the bond is determined by the number of bonded electrons (the bond order). The higher the bond order, the stronger the pull between the two atoms and the shorter the bond length.

Question 57

State the factors that affect covalent bond strength.

Answer 57

1) Number of bonds between atoms (single vs double vs triple bonds)
2) Effectiveness of overlap of orbitals
3) Differences in the electronegativities of bonding atoms (bond polarity)
4) Type of hybridisation of the orbitals of the bonding atoms (Chem bonding II)

Question 58

Define bond energy. State what it can be used for.

Answer 58

Bond energy is defined as the average energy absorbed when one mole of a particular bond is broken in the gaseous state. It is used to measure the strength of a covalent bond (not intermolecular forces of attraction).

Question 59

Explain how the number of bonds between atoms (single vs double vs triple bonds) affects covalent bond strength.

Answer 59

For the same number of bonding atoms (e.g. 2 carbon atoms), an increase in the number of bonds increases the number of shared electrons between the two atoms and hence increases electron density. The increased electrostatic force of attraction between the bond pairs and the two nuclei increases bond strength.

Question 60

Explain how the effectiveness of overlap of orbitals affects covalent bond strength.

Answer 60

The more effective orbital overlap, the stronger the bond. In halogens, as the halogen increases in size, the valence orbital used in bonding is more diffused and the overlap of the orbitals is less effective and bond energy decreases.

Question 61

Explain how the differences in the electronegativities of bonding atoms (bond polarity) affect bond strength.

Answer 61

Other than the electrostatic force of attraction between shared pair of electrons and the two nuclei involved in the covalent bonds, there is an additional electrostatic force of attraction between the two partial charges of the molecule. The greater the difference in electronegativity, the more polar the covalent bond (greater the bond polarity), the greater the ionic character, hence the stronger the bond. As compared to a non-polar covalent bond, there is an increase in electrostatic attraction in a polar covalent bond due to the electrostatic forces of attraction between two partial charges, which leads to increased bond strength.

Question 62

What are the two types of intermediate bond types?

Answer 62

Ionic bond with some covalent character

Covalent bond with some ionic character

Question 63

Explain how ionic bonds with covalent character arise.

Answer 63

The attraction of the cation for the valence shell electrons of the anion causes polarisation (distortion) of the electron cloud in the anion, pulling it into the region between the nuclei, resulting in some form of orbital overlap i.e. covalent character as there is a presence of electron density/degree of covalency between the ions.

Question 64

Explain the factors which affect the degree of the covalent character of an ionic bond.

Answer 64

The more polarised the anion, the higher the degree of the covalent character.

The degree of the covalent character depends on the:

1) Polarising power of the cation which refers to the cation’s ability to polarise the anion, depends on its charge density. The higher the charge and the smaller the cation, the stronger is its polarising power.
2) Polarisability of the anion which refers to the ease of being polarised, depends on its charge and size. Anions of lower charge density are more polarizable. For anions having the same charge, the larger anion is polarised to a greater extent (F– < Cl– < Br– < I–) as electron clouds of large anions are easier to distort than those of small anions.

Question 65

Explain why AlCl3 is covalent (for reference AlF3 is ionic).

Answer 65

If AlCl3 were ionic, the high charge density of Al3+ ion would distort the electron cloud of the Cl- ions to such an extent that electron sharing becomes predominant. The difference in electronegativities between Al and Cl is sufficiently small hence it exists as a covalent compound with id-id interactions between its molecules. (This is because The larger the value of the electronegativity, the greater the atom’s strength to attract a bonding pair of electrons. You have a non-polar covalent bond anytime the two atoms involved in the bond are the same or anytime the difference in the electronegativities of the atoms involved in the bond is very small.)

Question 66

Explain why graphite conducts electricity, but not diamond.

Answer 66

Graphite has a layered structure in which 3 of the 4 valence electrons of carbon are used for bonding, leaving the 4th electron being delocalised over the layer, forming an extended pi-electron cloud. These delocalised electrons are able to move freely and act as mobile charge carriers. For diamond, all 4 valence electrons of carbon are used for covalent bonding with the other carbon atoms. Hence, it has no free mobile charge carriers to conduct electricity.

Question 67

Cl2 and I2 are soluble in non-polar solvents but ICl is insoluble. Explain why ICl is insoluble in non-polar solvent.

Answer 67

ICl is a polar molecule with strong permanent dipole-permanent dipole (pd-pd) interactions between the molecules. The energy released from the weak instantaneous dipole-induced dipole interactions between the ICl molecules and the non-polar solvent molecules are insufficient to compensate for the energy required to overcome the strong pd-pd interactions between ICl molecules. (Note that iodine is the partially positive atom)

Question 68

Explain why salicylic acid and its isomer, 4-hydroxybenzoic acid, differ in their boiling points.

Answer 68

Molecules of salicylic acid are able to form both intramolecular and intermolecular hydrogen bonds while molecules of 4-hydroxybenzoic acid are able to form only intermolecular hydrogen bonds. The formation of intramolecular hydrogen bonds in salicylic acid reduces the extensiveness of hydrogen bonding in salicylic acid. During boiling, energy is absorbed to only break the intermolecular hydrogen bonds in order to separate the molecules. Hence, more energy is required to break the more extensive intermolecular hydrogen bonds in 4-hydroxybenzoic acid than that in salicylic acid. Therefore, 4-hydroxybenzoic acid has a higher boiling point than salicylic acid.

Question 69

Describe the physical properties of a giant ionic lattice structure

Answer 69

lol type again I accidentally deleted

Question 70

Suggest a reason why IF7 exists, but ClF7 does not exist.

Answer 70

Chlorine is a smaller atom compared to iodine, so it cannot ‘pack’ as many fluorine atoms
around itself, due to repulsion of electron clouds between the F atoms / overcrowding /
steric factors / steric hindrance

Cannot use inability to expand octet as reason as
Cl is in period 3 – link to size of central atom

Question 71

Explain why compound A is a liquid at room temperature even though it is an ionic
compound

Answer 71

Given ionic compound, we must link to the lattice energy –
incorrect to attribute to this being predominantly covalent
(cation weak polarizing power + question states ionic) – must
talk about weak ionic bonds / small magnitude of LE

Due to the large size and low charge of the cation and anion, the lattice energy (LE) is
not very exothermic OR the magnitude of LE is small, so the ionic bond is rather weak.

Question 72

A: CH3OH (methanol)
B: CH3CH2CH2CH2CH2OH (pentan-1-ol)
Why is A not soluble in hexane but B soluble in hexane? [2]

Answer 72

: CH3OH (methanol)
B: CH3CH2CH2CH2CH2OH (pentan-1-ol)
Why is A not soluble in hexane but B soluble in hexane? [2]
For A, instantaneous dipole – induced dipole (id-id) interactions formed between A and
hexane are weaker than intermolecular hydrogen bonding between molecules of A. Energy
released from formation of id-id interactions between hexane and A cannot compensate for
energy absorbed to break the intermolecular hydrogen bonds between molecules of A. [1]
For B, the predominant intermolecular interactions between molecules is id-id due to the
relatively long non-polar hydrocarbon/alkyl chain. Energy released from the formation of
id-id interactions between the non-polar hydrocarbon chain and hexane can compensate
for energy absorbed in breaking intermolecular forces between molecules in B and
hexane (there is both id-id (predominant) and hydrogen bonding in B).

Question 73

Water, ammonia and hydrogen fluoride have relative molecular masses which are very close, but they have very different boiling points. Explain why this is so.

Answer 73

H2O, NH3 and HF all have intermolecular hydrogen bonding between molecules. However,
H2O can form 2 hydrogen bonds per molecule whereas NH3 and HF can only form 1 hydrogen bond per molecule. Therefore more energy is needed to overcome the more extensive hydrogen bonding between water molecules, resulting in its highest boiling point. As F is more electronegative than N, the hydrogen bond between HF molecules is stronger
than that between NH3 due to the greater partial negative charge formed on F and partial positive charge on H atoms on HF molecule. More energy is therefore needed to separate HF molecules resulting in HF having a higher boiling point than NH3.

Question 74

Explain with a suitable diagram why 2-hydroxybenzoic acid has a lower solubility in water than
4-hydroxybenzoic acid

Answer 74

Due to the close proximity of the –OH to the –COOH
groups, 2-hydroxybenzoic acid form intramolecular
hydrogen bonding as shown in the diagram on the right.
Thus it has less sites available for the formation of
intermolecular hydrogen bonding with water
molecules.
Hence 2-hydroxybenzoic acid forms less extensive
intermolecular hydrogen bonding with water molecules
compared to 4-hydroxybenzoic acid, resulting in lower
solubility in water

Question 75

State the answering technique for structure and bonding questions.

Answer 75

Answering technique for structure and bonding questions:
1. State the structure of the compounds (giant ionic structure)
2. Make reference to the type of bonding in the compound (ionic bonds
between oppositely charged ions)
3. Compare the factors affecting the strength of bond if they are of same type

Question 76

Explain in terms of bonding and structure, why the boiling point of PCl3 (76.1 °C) is higher than that of SiCl4 (57.3 °C), but lower than that of
SiBr4 (153 °C).

Answer 76

SiCl4 is a non-polar molecule, and its intermolecular forces consist of the instantaneous dipole–induced dipole interactions only. PCl3 is a polar molecule, and the intermolecular forces consist of permanent dipole–permanent dipole and instantaneous dipole–induced dipole interactions, which are overall stronger than the intermolecular forces in SiCl4. Hence the boiling point of PCl3 is higher than SiCl4 that of since more energy is required to separate the PCl3 molecules during boiling.

SiBr4 is non-polar but it has a much larger electron cloud compared to PCl3. A larger electron cloud is more easily polarised, hence, the intermolecular instantaneous dipole-induced dipole interactions found in SiBr4 is significantly stronger than the intermolecular forces in PCl3 which has a much smaller electron cloud. This results in a higher boiling point of SiBr4, even though SiBr4 does not have pd-pd interactions.