1.4 Bonding

Question 1

ionic bonding

Answer 1

one atom donates one or more electrons to the other atom(s), resulting in a positive ion (cation) and negative ion (anion). Both ions have full outer shells of electrons.

Question 2

eg of ionic bonding sodium chloride

Answer 2

Question 3

eg of ionic bonding calcium bromide

Answer 3

Question 4

covalent bonds

Answer 4

the atoms share a pair of electrons to form a single covalent bond.
Each atom gives one electron to the bonding pair.
If two pairs of electrons are shared, a double bond is formed.
Each of the atoms in the molecule usually has a full outer shell of electrons.

Question 5

eg covalent
chlorine and oxygen

Answer 5

Question 6

coordinate bond

Answer 6

A coordinate bond is a covalent bond, but one of the atoms provides both electrons in the shared pair.

Question 7

eg coordinate bond
ammonia

Answer 7

Question 8

Answer 8

Question 9

In a covalent (and coordinate) bond, there is a

Answer 9

strong electrostatic attraction between the shared pair of electrons and the nuclei of both atoms. This outweighs the repulsion between the electrons in the shared pair. Also, both electrons in the bond have opposing spins to minimise this repulsion.

Question 10

In an ionic bond, there is

Answer 10

an electrostatic attraction between two oppositely charged ions. In fact, the positive ions and negative ions are arranged so that each positive ion is surrounded by several negative ions and vice versa, to maximise the attractive forces and minimise the repulsive forces.

Question 11

Answer 11

Question 12

Answer 12

the same wrong

change to: opposed

Question 13

Answer 13

wrong = minimise

change to = maximise

Question 14

Answer 14

wrong = an attractive

change to = a repulsive

Question 15

define electronegativity

Answer 15

Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. It is measured on the Pauling scale. Some values are given below.

Question 16

how does electronegativity change across and peoriod and down a group

Answer 16

Electronegativity increases across a period and decreases down a group. This means that fluorine is the most electronegative element and caesium is the least. The greater the electronegativity value, the stronger the attracting power of the element for the bonding pair of electrons.

Question 17

order of electronegativity

Answer 17

more protons, more electronegative

Question 18

Suggest why noble gases do not have electronegativity values.

Answer 18

Their outermost shells are full, and therefore they do not have a tendency to gain or attract electrons.

Question 19

factors affecting electronegativity

Answer 19

number of protons
distance from nucleus
screening from inner electrons

Question 20

Why do polar molecules form

Answer 20

A polar bond forms due to a difference in electronegativity, creating a permanent dipole.

The greater the difference, the more polar the bond.

If atoms are the same, they share electrons equally → bond is non-polar.

Question 21

Answer 21

Question 22

intermolecular forces

Answer 22

Intermolecular forces are forces between molecules. These intermolecular forces are much weaker than covalent, ionic or metallic bonds.

Question 23

three types of intermolecular forces

Answer 23

induced dipole-induced dipole forces, permanent dipole-dipole forces and hydrogen bonds

Question 24

Polar molecules have dipoles

Answer 24

One end has a slightly positive charge, the other a slightly negative charge due to a difference in electronegativities between the atoms in the molecule. If these dipoles arrange themselves so that the negative region of one molecule is close to the positive region of another molecule, there will be an attraction between them.

These are called permanent dipole-dipole interactions and are an example of van der Waals forces.

Question 25

describe forces

Answer 25

Permanent dipole-dipole interactions between HCl molecules.

Question 26

Induced dipole-induced dipole forces

Answer 26

🔹 1. Induced Dipole–Induced Dipole (London Forces) Weakest IMF Occurs in all molecules (polar & non-polar) Caused by temporary shifts in electron density Stronger in larger molecules (↑ electrons, ↑ surface area) More London forces = higher boiling point also called Van der Waals

Question 27

Permanent Dipole–Dipole Interactions

Answer 27

🔸 2. Permanent Dipole–Dipole Interactions Occurs between polar molecules Caused by differences in electronegativity Greater difference = stronger dipole If electronegativity is equal → non-polar

Question 28

Hydrogen Bonding

Answer 28

🌟 3. Hydrogen Bonding Strongest IMF (but still weaker than covalent bonds) Occurs when H is bonded to N, O, or F H atom (delta+) attracts a lone pair on a nearby electronegative atom Requires a large electronegativity difference Leads to higher boiling points and solubility in water The δ+ hydrogen atom is sandwiched between two electronegative atoms. It is covalently bonded to one and hydrogen-bonded to the other.

Question 29

hydrogen bonding eg water

Answer 29

Question 30

🧪 Boiling Point Trend (Strength of Intermolecular Forces):

Answer 30

🧪 Boiling Point Trend (Strength of Intermolecular Forces): Induced Dipole < Permanent Dipole < Hydrogen Bonding < Covalent

Question 31

can dipole be formed in non-polar molecules

Answer 31

A dipole can still be formed in non-polar molecules. This is due to the constant movement of electrons around the nuclei where sometimes more electrons are concentrated on one side of the atom at any one time, causing a temporary dipole.

Question 32

The strength of induced dipole-induced dipole forces increases with

Answer 32

The strength of induced dipole-induced dipole forces increases with increasing number of electrons in the molecule. The more electrons in a molecule, the greater the fluctuation in the electron cloud around the nuclei, and the larger the temporary and induced dipoles created. This means stronger forces of attraction between the molecules.

Question 33

Boiling Point Trends in Group Hydrides: Group 4

Answer 33

Group 4 hydrides: Boiling points increase down the group due to more electrons → stronger induced dipole forces (van der Waals).

Question 34

Boiling Point Trends in Group Hydrides: Group 5 and 6

Answer 34

Group 5 and 6 hydrides (e.g. NH₃, H₂O) have abnormally high boiling points due to hydrogen bonding, which is stronger than van der Waals forces. For the rest of Group 5 and 6 hydrides, boiling points follow the same trend as Group 4: ↑ molar mass = ↑ boiling point due to stronger van der Waals forces.

Question 35

Boiling Point Trends in Group Hydrides: Group 7

Answer 35

Group 7 hydrides (e.g. HF) also show high boiling points due to hydrogen bonding.

Question 36

💧 Solubility: eg hydrogen bonding

Answer 36

Molecules like H₂O and NH₃ can form hydrogen bonds with water, making them soluble. Compounds like CH₄, which only have van der Waals forces, cannot hydrogen bond with water → insoluble.

Question 37

Answer 37

HF HCl – ✖ H is bonded to Cl (chlorine), which is not electronegative enough and too large to form hydrogen bonds effectively. HF – ✅ H is bonded to fluorine, which is very electronegative with lone pairs → can form strong hydrogen bonds. CH₃CH₃ (ethane) – ✖ Only contains C–H bonds → no hydrogen bonding. CH₃OCH₃ (dimethyl ether) – ✖ Has oxygen with lone pairs, but no hydrogen directly bonded to O → cannot hydrogen bond itself, though it can accept H-bonds from water.

Question 38

Answer 38

Question 39

diagram of bond angles up to 1-6 bonding pairs

Answer 39

Question 40

Covalent molecules and molecular ions have fixed shapes. The VSEPR principle is used to predict these shapes. The principle states: 4 things

Answer 40

1. The shape of a molecule or molecular ion is determined by the number of electron pairs in the valence (outer) shell around the central atom. 1. Electron pairs can be bonding pairs, (shared pair of electrons), or lone pairs, (unshared pair of electrons). 1. All electron pairs repel each other (since electrons are all negatively charged). The shape that is formed is one that enables the electron pairs to keep as far away from each other as possible so that repulsion is minimised. 1. Bonding pairs are spread out between the two bonding atoms, but lone pairs stay close to the central atom. As a result, lone pairs repel more than bonding pairs. This leads to the following sequence of repulsion for electron pairs: lone pair-lone pair repulsion > lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion.

Question 41

How to find the number of electron pairs

Answer 41

Write down the number of electrons in the outer shell of the central atom. (This will be the same as the periodic table group number.) Add one electron for each bond being formed. (The formula will give the number of bonds.) Allow for any ion charge on the central atom. (If the ion has a charge of 1-, then add one electron. For a 1+ charge, deduct one electron.) Divide the total number of electrons by 2 to find the number of electron pairs. Compare the number of electron pairs with the number of bonds to find the number of bonding pairs and lone pairs. (Use the formula to find the number of bonding pairs.) For a compound with double bonds, remember for step 2 that each double bond donates two electrons, and for step 5 a double bond counts as two bonds.

Question 42

Eg electron pairs Carbon dioxide, CO2

Answer 42

Carbon is in group 4, so has 4 electrons in its outer shell. Oxygen forms 2 double bonds with carbon, so 4 electrons are added from the oxygen atoms. CO2 is a molecule, not an ion, so has no charge. The total number of electrons is 8, therefore there are 4 electron pairs. All the electron pairs are bonding because a double bond contains 2 electron pairs.

Question 43

Eg electron pairs Nitrogen trichloride, NCl3

Answer 43

Nitrogen is in group 5, so has 5 electrons in its outer shell. Chlorine forms 3 bonds with nitrogen, so 3 electrons are added from the chlorine atoms. NCl3 is a molecule, not an ion, so has no charge. The total number of electrons is 8, therefore there are 4 electron pairs. Three electron pairs are involved in bonding with the chlorine atoms, so there is 1 lone pair.

Question 44

name

Answer 44

For example, phosphorus pentachloride:

Question 45

name

Answer 45

sulfur hexafluoride

Question 46

Answer 46

Question 47

Answer 47

Question 48

eg of different bond angles

Answer 48