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Chemistry Alevel - Module 2 > 2.2.2 - Bonding and Structure > Flashcards

2.2.2 - Bonding and Structure Flashcards

1
Q

What is the definition of ionic bonding?

A

Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.

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2
Q

What happens when metal atoms form ions?

A

Metal atoms lose electrons to form positive ions.

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3
Q

What happens when non-metal atoms form ions?

A

Non-metal atoms gain electrons to form negative ions.

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4
Q

What is the electronic configuration of Mg and Mg²⁺?

A

Mg: 1s² 2s² 2p⁶ 3s² → Mg²⁺: 1s² 2s² 2p⁶

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5
Q

What is the electronic configuration of O and O²⁻?

A

O: 1s² 2s² 2p⁴ → O²⁻: 1s² 2s² 2p⁶

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6
Q

What factors increase the strength of ionic bonding and melting point?

A

Smaller ion size and higher ionic charge.

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7
Q

Why does MgO have a higher melting point than NaCl?

A

Mg²⁺ and O²⁻ are smaller and more highly charged than Na⁺ and Cl⁻.

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8
Q

What is the structure of an ionic solid called?

A

A giant ionic lattice.

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9
Q

Do the sticks in lattice diagrams represent ionic bonds?

A

No, they help show ion arrangement, not actual bonds.

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10
Q

What attracts each ion in a giant ionic lattice?

A

Each ion is attracted to all surrounding oppositely charged ions.

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11
Q

Why do ionic compounds have high melting points?

A

Due to strong electrostatic forces between oppositely charged ions.

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12
Q

Why don’t solid ionic compounds conduct electricity?

A

Ions are fixed in the lattice and cannot move.

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13
Q

Why do molten or dissolved ionic compounds conduct electricity?

A

Ions are free to move and carry charge.

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14
Q

Are ionic compounds usually soluble in water?

A

Yes, they are usually soluble in aqueous solvents.

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15
Q

What is the definition of a covalent bond?

A

A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

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16
Q

What is a dative covalent bond?

A

It forms when both shared electrons come from one atom; also called a coordinate bond.

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17
Q

What are examples of compounds with dative covalent bonds?

A

NH₄⁺, H₃O⁺, NH₃BF₃

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18
Q

How is the direction of a dative covalent bond shown?

A

An arrow from the lone pair donor to the atom receiving the electrons.

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19
Q

Does a dative covalent bond behave like a normal covalent bond in shapes?

A

Yes, e.g., NH₄⁺ is tetrahedral.

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20
Q

What is average bond enthalpy?

A

A measurement of covalent bond strength — higher means stronger.

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21
Q

What is the bonding in Sodium chloride and Magnesium oxide?

A

Ionic bonding.

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22
Q

What is the bonding in Iodine, Ice, CO₂, H₂O, and CH₄?

A

Covalent (simple molecular with intermolecular forces).

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23
Q

What are examples of intermolecular forces?

A

Induced dipole–dipole, permanent dipole–dipole, hydrogen bonding.

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24
Q

What is the difference in boiling/melting point between ionic and simple molecular substances?

A

Ionic: high due to strong electrostatic forces.
Molecular: low due to weak intermolecular forces.

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25
What is the solubility of ionic vs molecular substances in water?
Ionic: generally good. Molecular: generally poor.
26
Do ionic or simple molecular substances conduct electricity when solid?
No.
27
Do they conduct electricity when molten?
Ionic: Yes. Molecular: No.
28
Describe ionic solids.
Crystalline solids.
29
Describe molecular substances.
Mostly gases and liquids.
30
How do lone pairs affect bond angles?
They repel more than bonding pairs, reducing angles by about 2.5° per lone pair.
31
How do you explain molecular shape?
State bonding/lone pairs, repulsion concept, lone pairs repel more, then state shape and bond angle.
32
What is the role of wedges in molecular representations?
Wedges are used to represent three-dimensional molecular shapes on flat paper.
33
What are the types of wedges, and what kind of bond do they show?
1. Solid line: A bond in the plane of the paper. 2. Solid wedge: A bond coming out of the plane of the paper. 3. Dotted wedge: A bond going into the plane of the paper.
34
How do lone pairs differ from bonded pairs in terms of repulsion?
Lone pairs are closer to the central atom and occupy more space than bonded pairs. Lone pairs repel more strongly than bonded pairs.
35
What is the order of repulsion strength?
Increasing repulsion: Bonded-pair/Bonded-pair < Bonded-pair/Lone-pair < Lone-pair/Lone-pair
36
How do lone pairs affect bond angles?
Lone pairs push bonded pairs closer together, reducing the bond angle by about 2.5° per lone pair.
37
How does the electron repulsion theory apply to multiple bonds?
In molecules with multiple bonds, each multiple bond is treated as a single bonding region.
38
What happens to bond angles as the number of electron pairs increases?
The greater the number of electron pairs, the smaller the bond angle due to increased repulsion.
39
How is the shape of a methane (CH₄) molecule described?
Methane (CH₄) is symmetrical with four C-H covalent bonds arranged in a tetrahedral shape. The bond angle is 109.5°.
40
What is the shape of the ammonium ion (NH₄⁺)?
The ammonium ion (NH₄⁺) has four bonded pairs around the central nitrogen atom, giving it a tetrahedral shape with bond angles of 109.5°, similar to methane.
41
What is electronegativity?
The relative tendency of an atom in a covalent bond to attract electrons.
42
What are the most electronegative elements?
F, O, N, Cl.
43
What increases electronegativity across a period?
More protons and smaller atomic radius.
44
What decreases electronegativity down a group?
Greater distance from nucleus and more shielding.
45
When is a bond purely covalent?
When electronegativity difference is small.
46
When does a polar covalent bond form?
When atoms have different electronegativities.
47
What does a polar bond result in?
A dipole with partial charges (δ⁺ and δ⁻).
48
Which atom is δ⁻ in a polar bond?
The one with higher electronegativity.
49
When is a bond ionic?
When the electronegativity difference is very large.
50
Why is CCl₄ non-polar and CH₃Cl polar?
CCl₄ is symmetrical (dipoles cancel), CH₃Cl is asymmetrical.
51
Can a molecule have polar bonds but still be non-polar?
Yes, if it is symmetrical.
52
What are induced dipole–dipole interactions also called?
London forces.
53
Where do induced dipole–dipole interactions occur?
Occur between all simple covalent molecules and the separate atoms in noble gases.
54
How do temporary dipoles form?
In any molecule the electrons are moving constantly and randomly. Because of this, Electron density fluctuates randomly, and parts of the molecule become more or less negative.
55
What causes induced dipoles in neighbouring molecules? What are they called?
A temporary dipole in one molecule. These are called induced dipoles.
56
What affects the strength of London forces (Induced dipole-dipole)?
Number of electrons: more electrons = stronger forces = higher boiling point.
57
What are permanent dipole–dipole forces?
Attractive forces between polar molecules.
58
How do permanent dipole-dipole forces compare to induced dipole-dipole interactions?
Permanent dipole–dipole forces are stronger than induced dipole–dipole interactions.
59
How do permanent dipole–dipole forces affect boiling points?
They increase boiling points because stronger intermolecular forces require more energy to overcome.
60
What bonds typically create permanent dipoles?
C–Cl C–F C–Br H–Cl C=O
61
Do permanent dipoles occur in addition to London forces?
Yes.
62
What causes the trend in boiling points of alkanes?
Increased electrons and surface area lead to stronger London forces.
63
Why does I₂ have a higher boiling point than Cl₂?
I₂ has more electrons → stronger London forces.
64
How does shape affect induced dipole forces?
Long chains have more surface contact → allows for more induced dipole-dipole interactions → stronger forces.
65
Do ionic substances have London or dipole forces?
No, only molecular and noble gases do.
66
What makes a molecule polar?
A molecule is polar if it is asymmetrical and contains a bond with a large difference in electronegativity between the atoms.
67
What are Van der Waals’ forces?
A collective term for induced and permanent dipole interactions.
68
What is hydrogen bonding?
It is a strong type of intermolecular force that occurs when hydrogen is bonded to nitrogen, oxygen, or fluorine.
69
What condition must N, O, or F meet to allow hydrogen bonding?
They must have a lone pair of electrons available for bonding.
70
Why does hydrogen bonding occur with N, O, or F?
Because there is a large electronegativity difference between hydrogen and N, O, or F.
71
What are examples of bonds where hydrogen bonding can occur?
Examples include O–H, N–H, and F–H bonds.
72
How does hydrogen bonding compare to other types of intermolecular forces?
Hydrogen bonding is stronger than both induced dipole–dipole and permanent dipole–dipole interactions.
73
Why do H₂O, NH₃ and HF have high boiling points?
Because of hydrogen bonding between their molecules.
74
What causes the trend from H₂S to H₂Te?
Increasing electrons → stronger London forces.
75
Which functional groups or compounds can hydrogen bond?
Alcohols, carboxylic acids, proteins, amides.
76
What atoms must be bonded to hydrogen for H-bonding?
Nitrogen, oxygen, or fluorine (with lone pairs).
77
What must be shown in diagrams of hydrogen bonding?
Lone pairs, δ⁺ and δ⁻ dipoles, and all partial charges.
78
How many hydrogen bonds can a water molecule form?
Two hydrogen bonds per molecule.
79
Why can water form two hydrogen bonds per molecule?
Because oxygen is very electronegative and has two lone pairs of electrons.
80
Why is ice less dense than liquid water?
In ice, the hydrogen bonds hold molecules further apart, creating an open structure with lower density.
81
What type of bonding exists between the iodine atoms in an I₂ molecule?
Covalent bonds hold the two iodine atoms together in an I₂ molecule.
82
What type of structure do iodine crystals have?
A regular arrangement of I₂ molecules.
83
What holds I₂ crystals together?
Weak induced dipole–dipole forces between I₂ molecules.
84
Why do giant ionic substances have high boiling and melting points?
Because they have a giant lattice of ions with strong electrostatic forces between oppositely charged ions.
85
Why do simple molecular substances have low boiling and melting points?
Because of weak intermolecular forces between molecules (e.g., induced dipole–dipole, hydrogen bonds).
86
How soluble are giant ionic substances in water?
Generally good – many dissolve well in water.
87
How soluble are simple molecular substances in water?
Generally poor – most do not dissolve well in water.
88
Why don’t giant ionic substances conduct electricity when solid?
Because the ions are fixed in a lattice and cannot move.
89
Why don’t simple molecular substances conduct electricity when solid?
Because they have no ions and the electrons are localised (not free to move).
90
Why do giant ionic substances conduct electricity when molten?
Because the ions are free to move and carry electrical charge.
91
Why don’t simple molecular substances conduct electricity when molten?
Because they still have no ions to carry charge.
92
What is the general physical state of giant ionic substances at room temperature?
Crystalline solids.
93
What is the general physical state of simple molecular substances at room temperature?
Mostly gases or liquids.
94
How do you explain the shape of a molecule?
1. State number of bonding pairs and lone pairs of electrons. 2. State that electron pairs repel and try to get as far apart as possible (or to a position of minimum repulsion.) 3. If there are no lone pairs state that the electron pairs repel equally. 4. If there are lone pairs of electrons, then state that lone pairs repel more than bonding pairs. 5. State actual shape and bond angle.

Chemistry Alevel - Module 2 (7 decks)

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