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A Level - Chemistry > 2: Structure & Bonding > Flashcards

2: Structure & Bonding Flashcards

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1
Q

Define ionic bonding

A

Electrostatic attraction between two oppositely charged ions

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2
Q

What difference in electronegativity causes ionic bonding?

A

Typically greater than 1.5

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3
Q

How does the atomic radius change as you go down groups?

A

Atomic radius increases

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4
Q

How does the atomic/ionic radius of a group 1 metal compare to a group 2 metal?

A

Group 2 metals have a smaller atomic & ionic radius

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5
Q

Define covalent bonding

A

Electrostatic attraction between two nuclei and a shared pair of electrons between them

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6
Q

How are the electrons shared in normal covalent bonding?

A

Each atom provides one electron

Overlap of orbitals containing an electron from each

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7
Q

Define a dative covalent bond

A

Two atoms share a pair of electrons, with both electrons being donated by one atom

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8
Q

How does a dative covalent bond form?

A

An atom with a lone pair of electrons donates both electrons to another atom that needs 2 electrons to fill its outer shell

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9
Q

Name the shape and angles of a molecule with 2 electron pairs

A

Linear shape

180°

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10
Q

Name the shape and angles of a molecule with 3 electron pairs

A

Trigonal planar

120°

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11
Q

Name the shape and angles of a molecule with 4 electron pairs

A

Tetrahedral

109.5°

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12
Q

Name the shape and angles of a molecule with 5 electron pairs

A

Trigonal bi-pyramid

90° + 120°

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13
Q

Name the shape and angles of a molecule with 6 electron pairs

A

Octahedral

90°

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14
Q

Name the shape and angles of a molecule with 3 electron pairs + 1 lone pair

A

Trigonal pyramid

107°

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15
Q

Name the shape and angles of a molecule with 2 electron pairs and 2 lone pairs

A

Bent

104.5°

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16
Q

Why is ICl4- not tetrahedral?

A

6 electron groups
4 bonding pairs + 2 lone pairs
Square planar (90°)

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17
Q

Define polarity

A

In-balance of the sharing of electrons which causes a charge

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18
Q

Define electronegativity

A

A measure of the tendency of an atom in a molecule to attract the electrons in a covalent bond

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19
Q

Define a dipole

A

Sum of all the polarities on a molecule

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20
Q

What is the Pauling scale?

A

Measures electronegativity of elements from 4.0 to 0

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21
Q

Name some trends in electronegativity

A

Increases across a period
Decreases down a group
Non-metals tend to be more electronegative

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22
Q

What is the electronegativity in covalent bonding?

A

Very small/zero electronegativity

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23
Q

What is the electronegativity in ionic bonding?

A

Large difference in electronegativity

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24
Q

Describe ionic bonding using electronegativity

A

Atoms of significantly higher electronegativity (usually non-metals) form ionic bonds with atoms of significantly lower electronegativity (generally metals)

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25
Why does the electronegativity increase across a period?
Nuclear charge increases | Atomic radius stays roughly the same
26
What is a polar covalent bond?
Electronegative difference not too small nor large | Covalent bond forms but electrons are attracted towards the more electronegative atom
27
What do polar covalent bonds form?
Slightly negative and slightly positive charge on the atoms | Bond is polarised so it has a slight ionic character
28
What causes a permanent dipole?
Shape of the molecules | Polarity of individual bond
29
What is a permanent dipole?
The overall molecule is polar | Meaning there must be a net dipole
30
What is a non-polar molecule made of?
All identical bonds | No lone electron pairs on central atoms
31
What is a polar molecule made of?
Different bonds or | Identical bonds and lone electron pairs on the central atom
32
What are the three types of dipole?
Permanent Dipole Induced Dipole Instantaneous Dipole
33
What is an induced dipole?
No permanent dipole (similar electronegativity) | If it is next to a molecule with a dipole, a dipole can be induced in it
34
What is an instantaneous dipole?
Electron cloud may not be evenly distributed leading to the formation of a temporary dipole
35
How can you determine if a liquid is polar?
Observing if a jet of liquid is deflected by a charged rod
36
What are London forces?
Instantaneous dipole - induced dipole | Caused by electrons constantly moving which can cause a charge
37
What affects the strength of instantaneous dipoles?
Number of electrons (the more electrons the stronger the dipole) Proximity of molecules (closer they are the stronger)
38
Why do straight chain molecules have stronger id-id forces?
More points of contact as they pack better
39
What is a pd-pd interaction?
Two dipolar molecules are attracted to each other due to the different charges
40
What is hydrogen bonding?
The strongest form of pd-pd interactions | 10x weaker than covalent, 10x stronger than normal pd-pd
41
What molecules can hydrogen bonding occur with?
X - H ---Y | X & Y = N, O, F
42
How does hydrogen bonding explain the density of ice?
Density of ice is lower | As hydrogen bonds create a crystalline structure in ice, which means the molecules aren't close together
43
Why can only some molecules form hydrogen bonds?
N, O, F All very electronegative Contain lone pairs
44
Why is ethanol soluble in water?
Forms hydrogen bonds with water
45
When will liquids mix (miscible)?
If it is energetically advantageous
46
What is the explanation if two liquids are mixed and the temperature went down?
Inter-molecular forces of the liquids before are stronger than the attractions which would be formed in the mix This would be immiscible
47
What is the explanation if two liquids are mixed and the temperature stayed the same?
The bonding between the mix is the same as the individual liquids
48
What is the explanation if two liquids are mixed and the temperature went up?
Inter molecular forces in the individual liquids is weaker than that in the mixture
49
How are ionic solids soluble?
Ions surrounded by polar water molecules and energy is released therefore exothermic (hydration enthalpy is negative)
50
What are cations surrounded by when dissolved in water?
Slightly negative oxygen atoms
51
What are anions surrounded by when dissolved in water?
Slightly positive hydrogen atoms
52
What is hydration enthalpy?
Enthalpy change when a solution of ions is made from 1 mole of gaseous ions
53
What is lattice enthalpy?
Enthalpy change when 1 mole of solid is formed from the separate ions, always exothermic
54
What is the enthalpy of a solution?
Enthalpy change when one mole of an ionic solid dissolves in excess water If negative, the solid will dissolve
55
What is the equation of enthalpy for a solution?
Enthalpy of solution = Hydration enthalpy - lattice enthalpy
56
How does distance affect electrostatic attraction?
Gets weaker with distance
57
How does the small ionic radius of group 2 metals affect the electrostatic attraction?
As more closely packed they are able to have stronger ionic bonding than larger ions
58
How does the ionic radius change going down a group?
Increases
59
What are isoelectric ions?
Ions of different atoms with the same numbers of electrons
60
What is proof for ionic lattices having charged particles?
Migration of ions Green copper chromate on paper with cathode and anode Blue at cathode due to copper Yellow at anode due to chromate
61
What is the bond length?
Distance between two nuclei in covalent bonding
62
How does bond length affect the bond enthalpy?
The shorter the covalent bond, the higher the bond enthalpy
63
What is electron density in relation to covalent bonding?
The number of electrons within the covalent bond
64
How does electron density of covalent bonds affect bond enthalpy?
Higher the electron density Shorter the bond Therefore higher the enthalpy
65
How much do lone pairs repel in comparison to bond pairs?
Lone pairs repel more than bonding pairs
66
Between which electron pairs are the angles largest?
Lp/lp is largest angles Lp/bp is second largest angles Bp/bp is smallest angels
67
How can lone pair repulsion affect bonding pairs?
Bonding pair angles reduced as they are pushed together by lone pair repulsion
68
What halogen bonded to hydrogen has the shortest bond?
H - F | most electronegative
69
Which of the following has the largest ionic radius? | S 2-, Cl -, K +, Ca 2+
S 2- | Same electrons but less protons so not as strong electrostatic pull on electrons
70
Why does H2O have a bond angle of 104.5?
Electrons repel as far as possible Lone pair has greater repulsion than bond pair Therefore less than 109
71
What is an ionic bond?
Electrostatic attraction between oppositely charged ions
72
What is an ionic giant lattice?
3-D arrangement of ions which is regular | Alternating between positive and negative
73
What properties do ionic compounds have?
Hard crystalline structures High melting and boiling points Generally high solubility in water Good electrical conductivity when molten/dissolved
74
Why do ionic compounds not conduct electricity as a solid?
Ions cannot move out of the regular shape, therefore it cannot conduct electricity
75
What is a covalent bond?
Electrostatic attraction between two nuclei and a shared pair of electrons between them
76
What is a simple molecule?
Non-metals bonded together covalentlyUsually diatomic
77
What are the properties of simple covalent molecules?
Low melting point Non-conducting of heat or electricity More soluble in non-polar solvents (hexane not water) Dissolved solutions don't conduct electricity
78
What is a giant covalent structures?
Many atoms joined together by very strong covalent bonds
79
What is the structure of diamond?
Each C atom bonded to 4 others | Tetrahedral 3D arrangement
80
What is the structure of graphite?
Each C atom bonded to 3 others Hexagons within a layer 4 outer electrons, 3 used in bonding Free one delocalised and flow along layers
81
What keeps layers of graphite together?
London/van der Waals forces
82
What is graphene?
1 layer of graphite Hexagon shape Good electrical conductor Relatively strong
83
What is the structure of metals?
Giant metallic 3D lattices | Metal cations surrounded by a sea of deloacalised electrons, attracted by electrostatic forces
84
What does the structure of metals cause?
High melting point - strong electrostatic attractions High density - close packing Conducts electricity - delocalised electrons Malleable - bonding is non-directional so they can slide over each other Conducts heat - electrons can pass kinetic energy to each other Insoluble - strength of bonding
85
Why do ionic lattices form?
Each ion is electrostatically attracted in all directions to ions of the opposite charge
86
Why can ionic compounds not be shaped?
Same charged ions repel each other, so they break when its stretched or hammered Brittle
87
Why are ionic substances soluble in polar solvents?
Ions are pulled apart by polar molecules
88
What are the properties and explanations of giant covalent structures?
Very high melting points - need to break many strong bonds Extremely hard - same as above Good thermal conductors - Vibrations travel easily through the rigid structure Insoluble - no ions, atoms more attracted to neighbours in the lattice than the solvent Can't conduct electricity - no charged ions or free electrons
89
Why can graphite conduct electricity?
As only bonds to 3 other carbon atoms, there is one delocalised electron per atom Electrons can carry the charge
90
What bonding occurs in metallic bonding?
Electrostatic attraction between delocalised negative electrons and the positive metal ions
91
How does the number of electrons affect metallic bonding?
Stronger bonding as there is a larger electrostatic force
92
How do impurities affect the electrical conductivity of metals?
Reduce conductivity Reduces number of electrons that are free to move and carry charge As the electrons transfer to the impurity and forms anions

A Level - Chemistry (26 decks)

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