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A Level Chemistry > Topic 2 - Bonding > Flashcards

Topic 2 - Bonding Flashcards

1
Q

Definition if ionic bonding

A

A strong electrostatic force of attraction between oppositely charged ions in a lattice

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2
Q

Why do chemical bonds form?

A

When atoms share or transfer electrons to achieve a more stable electron arrangement, often a full outer main level of electrons.

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3
Q

Formula of sulfate

A

SO4 (2-)

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4
Q

Formula of nitrate

A

NO3 (-)

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5
Q

Formula of Hydroxide

A

OH (-)

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6
Q

Formula of carbonate

A

CO3 (2-)

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7
Q

Formula of Ammonium

A

NH4 (+)

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8
Q

Properties of ionically bonded compounds …

A

Always solid at room temp

Giant structures = high melting points

Only conducts when molten or in aqueous solution as ions are not free when solid

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9
Q

Definition of metallic bonding

A

Strong electrostatic force of attraction between positive metal ions and sea of delocalised electrons

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10
Q

Structure in metallic bonding

A

Giant metallic lattice

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11
Q

Properties of metals x4

A

Good conductors of electricity as delocalised electrons can move throughout the structure.

High thermal conductivity due to delocalised electrons and the closely packed ions allowing vigorous vibrations.

Malleable and Ductile - arranged in layers which can slide and each metal ion will still be in the same environment.

High melting points

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12
Q

What does the strength of a metal depend on?

A

Charge on the ion - bigger charge = more delocalised electrons and the strong the force of attraction between electrons and metal ions.

Size of the ion - Smaller ion means the closer the electrons are to the positive nucleus = stronger bond

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13
Q

What is a covalent bond?

A

A strong electrostatic force of attraction between positive nuclei of the atoms involved and the shared pair of electrons

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14
Q

What’s the difference between a single covalent bond and multiple covalent bonds?

A

A single covalent bond contains one pair of shared electrons

Multiple covalent bonds contain multiple shared pairs

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15
Q

What is a dative covalent bond?

A

A co-ordinate (dative covalent) bond contains a shared pair of electrons with both electrons supplied by one atom.

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16
Q

Melting and boiling of ionic structure

A

High as lots of energy is required to overcome the strong electrostatic forces of attraction between oppositely charged ions.

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17
Q

Other word for giant covalent

A

Macromolecular

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18
Q

Diamond structure and properties x4

A

Each carbon strongly, covalently bonded to 4 other carbons in a giant lattice of atoms which are arranged tetrahedrally.

High melting and boiling

Does not conduct electricity as there are no freely moving charged particles

Very hard as strong covalent bonds hold atoms in fixed positions

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19
Q

Graphite structure and properties?

A

Each carbon atom strongly, covalently bonded to 3 others forming a giant lattice of atoms in a trigonal planar arrangement.

High melting and boiling point

Does conduct electricity as between layers there are delocalised electrons that can move freely and conduct - can only conduct electricity along hexagonal planes and not vertically throughout the structure.

Soft and slippery as there are weak forces of attraction (VDW forces) between the layers which are easily overcome to allow the layers to slide relative to one another

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20
Q

Temperature of any substance is ….

A

Directly proportional to the mean kinetic energy of the particles

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21
Q

Why do liquids cool as they evaporate?

A

As some of the KE is used to overcome the forces of attraction between particles to allow them to escape.

The leads to a lower mean KE and therefore a lower T

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22
Q

What happens when a solid is heated?

A

Particles vibrate faster
KE + T increase

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23
Q

What happens during melting?

A

Heat energy is used to overcome or partially overcome the forces of attraction between particles instead of increasing the KE. Therefore T is constant through melting.

As KE remains constant

24
Q

What happens during boiling.

A

Heat energy is used to overcome the forces or bonds between the particles rather than increase the KE.

Liquid boils but KE remains constant so T remains constant

25
Ice and Iodine are examples of?
Molecular crystals
26
Describe Ice
Each water molecule comprises of one oxygen atom and two hydrogens. Each hydrogen atom is covalently bonded to the oxygen. Although the covalent bonds within molecules are very strong, there are **weak intermolecular forces of attraction** between molecules leading to a low melting and boiling point. Does **not** conduct as there are no freely moving charged particles
27
Describe Iodine.
Each molecule of iodine comprises two Iodine atoms which are joined by **one covalent bond** Although the covalent bond within the molecule is very strong, **weak intermolecular forces of attraction occur between molecules** **Low sublimation point** as IMFs require little energy to break Does not conduct electricity as there are no freely moving charged particle.
28
Define electronegativity.
The power of an atom to withdraw electron density towards itself in a covalent bond
29
What 3 factors affect electronegativity
**No. of protons in the nucleus** - more protons = stronger attraction **Size of atom** - Smaller atom means the e- in the covalent bond is closer to the nucleus = stronger attraction **Shielding** - more electron shells between the nucleus and the shared e- = weaker attraction
30
What is a polar bond?
A covalent bond existing between atoms with different electronegativity. This leads to unsymmetrical electron distribution and a partial positive charge on the less electronegative and partial negative charge on the more electronegative atom involved. If the molecule is symmetrical eg. F2 then electron density is shared equally and the bond is non polar.
31
Why do dipoles form?
Due to uneven distributions of electrons due to different electronegativities of the atoms in a molecule.
32
Why do some molecules with polar bonds not have permanent dipoles?
When a molecule is **symmetrical** - dipoles cancel each other out leading to no overall permanent dipole When molecules have **similar electronegativities**
33
When are weak IMFs broken?
When molecular substances are vapourised
34
List the IMFs strongest to weakest
Hydrogen bonds Permanent dipole dipole forces Van der Waals forces
35
How do van der Waals forces work?
The electron cloud is in constant motion This leads to a temporary electron imbalance Leading to a temporary dipole in the molecule which induces a temporary dipole in another molecule. Leading to a temporary attraction between the temporary dipoles, this is the Van der Waals force
36
What are Intermolecular forces?
Weak electrostatic forces of attraction acting between molecules
37
What makes the Van der Waals force stronger how does it correlate to its melting point?
As a molecule gets larger so does the boiling point There are more electrons leading to a stronger VDW force. This then requires more energy to overcome the forces between molecules
38
Where are VDW forces present?
In all molecular substances at all times
38
How do dipoles form?
Due to uneven distribution of electrons in a molecule. This is due to differences in electronegativity of the atoms in a molecule.
39
When do permanent dipole - dipole forces form?
Dipole - dipole attractions in **polar molecules**
40
When do permanent dipole - dipole forces not form?
In **non polar molecules** Sometimes individual bonds are polar but if the shape is **symmetrical** then the polar bonds cancel each other out.
41
What 2 conditions are needed for a hydrogen bond to form?
A very electronegative atom (N, O, F) with an available lone pair of electrons And a Hydrogen atom directly bonded to a very electronegative atom
42
What atoms are sufficiently electronegative for a Hydrogen bond to form?
N, O and F
43
When do the H bonds break in solid, liquid and gas?
Solid - Bonds are permanent Liquid - Continously breaking and reforming Gas - Permanently broken
44
Why does ice have a lower density than water?
As the molecules are held apart in a tetrahedral structure of hydrogen and covalent bonds. This is different to liquid water where the bonds are constantly breaking and reforming. Much empty space between molecules. Larger volume with the same mass of water = lower density
45
How and why do pairs electrons in the outer shell of an atom arrange themselves?
As far apart as possible as to minimise repulsion
46
Why do ionic compounds tend to dissolve in water?
Water molecules are polar so parts of the molecule are negatively and positively charged. Charged parts pull ions away from the lattice which cause it to dissolve.
47
2 BP 0 LP Name Angle Diagram
Linear 180
48
3 BP 0LP Name Angle Diagram
Trigonal Planar 120
49
4 BP 0LP Name Angle Diagram
Tetrahedral 109.5
50
5 BP 0 LP Name Angle Diagram
Trigonal Bipyramidal 120 and 90
51
6 BP 0LP Name Angle Diagram
Octahedral 90
52
3 BP 1LP Name Angle Diagram
Pyramidal 107
53
2 BP 2 LP Name Angle Diagram
V shaped 104 .5
54
4 BP 2 LP Name Angle Diagram
Square planar 90

A Level Chemistry (12 decks)

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