Bonding And Structure

Question 1

How are ionic bonds formed, and what structure do they create?

Answer 1

Ionic bonds form through electron transfer from a metal to a non-metal, creating positive and negative ions. Oppositely charged ions are held together by strong electrostatic forces in a giant ionic lattice.

Question 2

What ions do Group 1 typically form?

Answer 2

Group 1: lose 1 electron → 1+ ions (e.g., Na+)

Question 3

What ions do Group 2 typically form?

Answer 3

Group 2: lose 2 electrons → 2+ ions (e.g., Mg2+)

Question 4

What ions do Group 6 typically form?

Answer 4

Group 6: gain 2 electrons → 2- ions (e.g., O2-)

Question 5

What ions do Group 7 typically form?

Answer 5

Group 7: gain 1 electron → 1- ions (e.g., Cl-)

Question 6

Name and provide the formulas for common compound ions.

Answer 6

Ammonium: NH₄⁺
Hydroxide: OH⁻
Nitrate: NO₃⁻
Carbonate: CO₃²⁻
Sulfate: SO₄²⁻

Question 7

Give examples of ionic compounds and their ion ratios.

Answer 7

MgO: Mg²⁺ + O²⁻ in 1:1 ratio
MgCl₂: Mg²⁺ + 2Cl⁻ (1:2)

Question 8

Describe the structure and bonding of a giant ionic lattice.

Answer 8

A giant ionic lattice is a Regular, repeating 3D structure (e.g., NaCl cube)
Held together by strong ionic bonds

Question 9

What are the physical properties of ionic compounds?

Answer 9

• High MP/BP: Due to strong ionic bonds
• Conduct electricity: Only molten or in solution (ions free to move)
• Soluble in water: Water’s polarity disrupts the lattice

Question 10

Covalent Bonding

How are covalent bonds formed and what do they achieve for atoms?

Answer 10

Covalent bonds form when atoms non-metals share electrons.
Atoms achieve full outer shells (noble gas configuration).

Question 11

What are the three main types of covalent bonds

Answer 11

Single bond: 1 pair shared
Double bond: 2 pairs shared (e.g., CO₂)
Triple bond: 3 pairs shared (e.g., N₂)

Question 12

What are the general physical properties of simple covalent compounds?

Answer 12

Simple covalent compounds have Low MP/BP due to Weak intermolecular forces and
Don’t conduct electricity
• Examples: H₂O, CO₂, CH₄

Question 13

Describe the bonding and structure in diamond, and relate this to its physical properties

Answer 13

Each C atom forms 4 bonds in a tetrahedral structure

Diamonds are Extremely hard, have very high MP, good thermal conductor, insoluble, doesn’t conduct electricity

Question 14

Describe the bonding and structure in graphite

Answer 14

Each C atom forms 3 bonds; has 1 delocalised electron in a planar hexagonal. Multiple stacked layers of hexagonal carbons sheets with weak intermolecular forces between layers

Question 15

What defines a dative covalent bond, and how is it represented?

Answer 15

A dative covalent bond is where both electrons in the bond come from one atom.
It is represented using an arrow (→) pointing from the electron donor to the electron acceptor.

Question 16

Give two common examples of molecules that contain dative covalent bonds and explain their formation.

Answer 16

NH₄⁺ (Ammonium ion):
The nitrogen atom in NH₃ has a lone pair, which it donates to an H⁺ ion to form the NH₄⁺ ion.
• H₃O⁺ (Hydronium ion):
The oxygen atom in H₂O donates a lone pair to a proton (H⁺) to form H₃O⁺.

Question 17

Shapes of Molecules & Electron Pair Repulsion

What is the basic principle behind the Electron Pair Repulsion Theory (VSEPR), and what are the key rules?

Answer 17

The Electron Pair Repulsion Theory states that electron pairs repel each other to maximise their separation in space.
Key ideas include:
• Bonding pairs and lone pairs of electrons repel each other.
• Lone pairs repel more strongly than bonding pairs, causing bond angles to adjust accordingly.

Question 18

Outline the steps used to determine the shape of a molecule or ion using VSEPR theory.

Answer 18

1. Identify the central atom.
   
   2. Count its outer shell electrons (use the group number).
   3. Add electrons from bonded atoms.
   4. Adjust for charge:
      • Add electrons for a negative charge
      • Subtract electrons for a positive charge
   5. Divide the total number of electrons by 2 to get the number of electron pairs.
   6. Compare with the number of bonding pairs to determine the number of lone pairs.
   7. Use this to determine the molecular shape and bond angles.

Question 19

Using the VSEPR model, complete the table below with molecular shapes and bond angles for different combinations of bonding and lone pairs

Answer 19

Linear- 2 EP, 0 LP, 180°
Trigonal planar- 3 EP, 0LP, 120°
Bent- 3 EP, 1LP, <120°
Tetrahedral- 4 EP, 0 LP, 109.5°
Trigonal pyramidal- 4EP, 1LP, 107°
Bent- 4EP, 2LP, 104.5°
Trigonal bipyramidal- 5EP, 0LP, 90°, 120°
Seesaw- 5EP, 1LP, 86°, 102°
T-shaped- 5EP, 2LP, 87.5°
Octahedral- 6EP, 0LP, 90 °
Square pyramidal- 6EP, 1LP, <90°
Square planar- 6EP, 2LP, 90°

Question 20

What are intermolecular forces, and how do they differ from intramolecular forces?

Answer 20

Intermolecular forces are forces of attraction between separate molecules, whereas intramolecular forces are the chemical bonds (ionic, covalent, or metallic) within a molecule or compound.
IMFs are weaker than covalent or ionic bonds,
IMFs significantly affect physical properties like melting point, boiling point, volatility, viscosity, and solubility.

Question 21

List the three main types of intermolecular forces

Answer 21

Van der Waals’ Forces
Permanent Dipole-Dipole
Hydrogen Bonding

Question 22

Describe van der waals’ forces and state its relative strength

Answer 22

Fluctuations of the electron distribution around the nucleus cause Temporary dipoles which induce neighbouring dipoles molecules (present in all substances)
And are the Weakest intermolecular forces

Question 23

Describe Permanent Dipole-Dipole forces and state its relative strength

Answer 23

Permanent Dipole-Dipole is the Electrostatic attraction between molecules with permanent dipoles and its relative strength is Intermediate

Question 24

Describe hydrogen bonding and state its relative strength

Answer 24

Hydrogen bonding is where the hydrogen atom is covalently bonded to a highly electronegative atom (such as nitrogen, oxygen, or fluorine) and a lone pair of electrons on another electronegative atom. Hydrogen bonding is the Strongest IMF (but weaker than covalent bonds)

Question 25

What causes van der Waals’ forces and in which substances are they present?

Answer 25

Van der Waals’ forces arise due to temporary dipoles formed by the uneven distribution of electrons in a molecule. These dipoles induce dipoles in nearby molecules, leading to induced dipole-dipole interactions. They are present in all molecules, but are the only type of IMF in non-polar molecules such as O₂ and CH₄.

Question 26

What factors increase the strength of van der Waals’ forces?

Answer 26

• Number of electrons: More electrons leading to larger fluctuations in electron density resulting in stronger temporary dipoles. • Surface area: Greater surface area increases contact between molecules, enhancing induced dipole interactions. Example: Iodine (I₂) forms a solid due to strong van der Waals’ forces between large molecules with many electrons.

Question 27

Why do noble gases liquefy at low temperatures?

Answer 27

Noble gases such as argon and neon are monoatomic and non-polar. They liquefy due to van der Waals’ forces, which are the only forces acting between these atoms at low temperatures.

Question 28

Permanent Dipole-Dipole Interactions What conditions are required for permanent dipole-dipole interactions to occur?

Answer 28

These interactions occur between polar molecules — molecules with a permanent dipole caused by differences in electronegativity and asymmetrical shape, resulting in an uneven distribution of charge.

Question 29

How do permanent dipole-dipole interactions influence physical properties?

Answer 29

These interactions result in higher boiling and melting points compared to substances of similar size that only have van der Waals’ forces. Example molecules: HCl, CH₃Cl

Question 30

Can a molecule have both van der Waals’ and permanent dipole-dipole forces?

Answer 30

Yes. All molecules experience van der Waals’ forces, and polar molecules additionally experience permanent dipole-dipole interactions.

Question 31

Hydrogen Bonding (Strongest IMF) What are the requirements for hydrogen bonding to occur in a substance?

Answer 31

Hydrogen bonding occurs when: 1. Hydrogen is covalently bonded to a highly electronegative atom: N, O, or F 2. A lone pair on another N, O, or F atom is available for hydrogen to interact with 3. Represented as a dashed line, e.g. O–H···O

Question 32

List common molecules that exhibit hydrogen bonding.

Answer 32

Water (H₂O) Ammonia (NH₃) Hydrogen fluoride (HF) Alcohols, carboxylic acids, primary and secondary amines

Question 33

How does hydrogen bonding explain anomalous properties of water?

Answer 33

• High boiling point: Due to strong hydrogen bonds between H₂O molecules • High surface tension: Hydrogen bonds create cohesive forces (water molecules attract eachother more strongly ) • Ice is less dense than water: Hydrogen bonding arranges molecules in an open lattice structure, increasing volume and reducing density

Question 34

Rank the intermolecular forces by strength

Answer 34

Hydrogen Bonding-Strongest IMF Permanent Dipole-Dipole-Moderate Van der Waals’-Weakest

Question 35

State the impact on melting/boiling points, Volatility , Solubility and viscosity based on the intermolecular forces strengths

Answer 35

Melting/Boiling Points: Stronger IMFs result to Higher melting/boiling points • Volatility: Stronger IMFs result to Lower volatility (less easily evaporated) • Solubility: Polar molecules dissolve better in polar solvents (like dissolves like) • Viscosity: Strong hydrogen bonding results to Higher viscosity (thicker liquid)

Question 36

How does molecular polarity affect solubility?

Answer 36

Polar substances dissolve in polar solvents (e.g., water dissolves ethanol) due to strong interactions like hydrogen bonding or dipole-dipole. • Non-polar substances dissolve in non-polar solvents due to van der Waals’ interactions. This is due to the energetics of IMF disruption and formation

Question 37

Bond Polarity & Molecular Polarity What is electronegativity and how is it measured?

Answer 37

Electronegativity is the ability of an atom to attract the shared pair of electrons in a covalent bond. It is measured on the Pauling scale, which is a relative scale with fluorine assigned the highest value (4.0).

Question 38

Trends in electronegativity across a period and down a group

Answer 38

Increases across a period (due to increasing nuclear charge and decreasing atomic radius) Decreases down a group (increased shielding and atomic radius reduce attraction)

Question 39

What is a polar covalent bond, and how does it form?

Answer 39

A polar covalent bond forms when two atoms with different electronegativities share electrons unequally. The more electronegative atom pulls the shared electrons closer, creating a partial negative charge (δ⁻), while the less electronegative atom becomes partially positive (δ⁺). This separation of charge forms a dipole within the bond.

Question 40

Provide examples of polar and non-polar covalent bonds.

Answer 40

H–Cl are Polar covalent because Cl more electronegative than H so dipole forms C–H are Slightly polar due to Small difference in electronegativity Cl–Cl are Non-polar covalent because Atoms have identical electronegativity O–H are Polar covalent because O is more electronegative than H

Question 41

What is a permanent dipole, and how does it differ from an instantaneous dipole?

Answer 41

A permanent dipole is a fixed separation of charge across a bond due to a constant difference in electronegativity between two atoms (e.g. in HCl). This differs from an instantaneous dipole, which is temporary and caused by fluctuating electron density, as seen in van der Waals’ forces.

Question 42

What determines whether a molecule is polar or non-polar overall?

Answer 42

A molecule’s overall polarity depends on: 1. The polarity of its bonds (i.e., if they have dipoles) 2. The symmetry of the molecule’s shape If the molecule is asymmetrical, the dipoles do not cancel out, resulting in a net dipole forming polar bonds If the molecule is symmetrical, dipoles cancel out making the molecule non-polar, even if it has polar bonds.

Question 43

Give examples of molecules with polar bonds that are non-polar overall.

Answer 43

CO₂ have polar bonds, but are non-polar due to its Linear shape where dipoles cancel each other. CCl₄ have polar bonds but are non-polar because they have a Tetrahedral shape so symmetrical dipoles cancel CH₃Cl have polar bonds and are polar molecules because they are arranged asymmetrically so dipoles do not cancel making molecule polar

Question 44

How does molecular polarity influence physical properties?

Answer 44

Boiling and melting points: Polar molecules experience permanent dipole-dipole interactions, which raise boiling/melting points compared to non-polar equivalents. • Solubility: Polar molecules are more soluble in polar solvents (like water), due to favourable dipole-dipole or hydrogen bonding interactions. • Miscibility: Polar and non-polar liquids often do not mix (immiscible) due to incompatible intermolecular forces.

Question 45

Why does water dissolve ionic and polar covalent substances?

Answer 45

Water is a highly polar solvent due to its bent shape and strong O–H dipoles. • It can surround and stabilize ions by ion-dipole interactions (e.g. Na⁺ and Cl⁻). • water can form hydrogen bonds with polar molecules (e.g. ethanol), increasing solubility. Thus, “like dissolves like” — polar substances dissolve well in polar solvents.

Question 46

What is the difference between bond polarity and molecular polarity?

Answer 46

Bond polarity refers to the dipole within a single bond due to a difference in electronegativity between bonded atoms. Molecular polarity is the overall dipole moment of the molecule and depends on both bond polarities and molecular geometry (symmetry).

Question 47

How does molecular geometry affect dipole cancellation?

Answer 47

In symmetrical molecular shapes (e.g., linear, tetrahedral with identical atoms), dipole moments cancel out due to even spatial distribution. In asymmetrical molecules (e.g., bent or trigonal pyramidal), dipoles do not cancel, resulting in a net molecular dipole. Example: • CH₄ (tetrahedral): Non-polar • NH₃ (trigonal pyramidal): Polar

Question 48

What are the properties of graphite

Answer 48

Conducts electricity, soft/slippery, low density- due to weak intermolecular forces between layers , insoluble high MP- due to strong covalent bonds