Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

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

2: Structure & Bonding Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

Define ionic bonding

A

Electrostatic attraction between two oppositely charged ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What difference in electronegativity causes ionic bonding?

A

Typically greater than 1.5

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

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

A

Atomic radius increases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Define covalent bonding

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Define a dative covalent bond

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

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

A

Linear shape

180°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

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

A

Trigonal planar

120°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

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

A

Tetrahedral

109.5°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

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

A

Trigonal bi-pyramid

90° + 120°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

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

A

Octahedral

90°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

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

A

Trigonal pyramid

107°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

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

A

Bent

104.5°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Why is ICl4- not tetrahedral?

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Define polarity

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Define electronegativity

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Define a dipole

A

Sum of all the polarities on a molecule

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What is the Pauling scale?

A

Measures electronegativity of elements from 4.0 to 0

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Name some trends in electronegativity

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

What is the electronegativity in covalent bonding?

A

Very small/zero electronegativity

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

What is the electronegativity in ionic bonding?

A

Large difference in electronegativity

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Why does the electronegativity increase across a period?

A

Nuclear charge increases

Atomic radius stays roughly the same

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

What is a polar covalent bond?

A

Electronegative difference not too small nor large

Covalent bond forms but electrons are attracted towards the more electronegative atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

What do polar covalent bonds form?

A

Slightly negative and slightly positive charge on the atoms

Bond is polarised so it has a slight ionic character

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

What causes a permanent dipole?

A

Shape of the molecules

Polarity of individual bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

What is a permanent dipole?

A

The overall molecule is polar

Meaning there must be a net dipole

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

What is a non-polar molecule made of?

A

All identical bonds

No lone electron pairs on central atoms

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

What is a polar molecule made of?

A

Different bonds or

Identical bonds and lone electron pairs on the central atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

What are the three types of dipole?

A

Permanent Dipole
Induced Dipole
Instantaneous Dipole

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

What is an induced dipole?

A

No permanent dipole (similar electronegativity)

If it is next to a molecule with a dipole, a dipole can be induced in it

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

What is an instantaneous dipole?

A

Electron cloud may not be evenly distributed leading to the formation of a temporary dipole

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

How can you determine if a liquid is polar?

A

Observing if a jet of liquid is deflected by a charged rod

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

What are London forces?

A

Instantaneous dipole - induced dipole

Caused by electrons constantly moving which can cause a charge

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

What affects the strength of instantaneous dipoles?

A

Number of electrons (the more electrons the stronger the dipole)
Proximity of molecules (closer they are the stronger)

38
Q

Why do straight chain molecules have stronger id-id forces?

A

More points of contact as they pack better

39
Q

What is a pd-pd interaction?

A

Two dipolar molecules are attracted to each other due to the different charges

40
Q

What is hydrogen bonding?

A

The strongest form of pd-pd interactions

10x weaker than covalent, 10x stronger than normal pd-pd

41
Q

What molecules can hydrogen bonding occur with?

A

X - H —Y

X & Y = N, O, F

42
Q

How does hydrogen bonding explain the density of ice?

A

Density of ice is lower

As hydrogen bonds create a crystalline structure in ice, which means the molecules aren’t close together

43
Q

Why can only some molecules form hydrogen bonds?

A

N, O, F
All very electronegative
Contain lone pairs

44
Q

Why is ethanol soluble in water?

A

Forms hydrogen bonds with water

45
Q

When will liquids mix (miscible)?

A

If it is energetically advantageous

46
Q

What is the explanation if two liquids are mixed and the temperature went down?

A

Inter-molecular forces of the liquids before are stronger than the attractions which would be formed in the mix
This would be immiscible

47
Q

What is the explanation if two liquids are mixed and the temperature stayed the same?

A

The bonding between the mix is the same as the individual liquids

48
Q

What is the explanation if two liquids are mixed and the temperature went up?

A

Inter molecular forces in the individual liquids is weaker than that in the mixture

49
Q

How are ionic solids soluble?

A

Ions surrounded by polar water molecules and energy is released therefore exothermic
(hydration enthalpy is negative)

50
Q

What are cations surrounded by when dissolved in water?

A

Slightly negative oxygen atoms

51
Q

What are anions surrounded by when dissolved in water?

A

Slightly positive hydrogen atoms

52
Q

What is hydration enthalpy?

A

Enthalpy change when a solution of ions is made from 1 mole of gaseous ions

53
Q

What is lattice enthalpy?

A

Enthalpy change when 1 mole of solid is formed from the separate ions, always exothermic

54
Q

What is the enthalpy of a solution?

A

Enthalpy change when one mole of an ionic solid dissolves in excess water
If negative, the solid will dissolve

55
Q

What is the equation of enthalpy for a solution?

A

Enthalpy of solution = Hydration enthalpy - lattice enthalpy

56
Q

How does distance affect electrostatic attraction?

A

Gets weaker with distance

57
Q

How does the small ionic radius of group 2 metals affect the electrostatic attraction?

A

As more closely packed they are able to have stronger ionic bonding than larger ions

58
Q

How does the ionic radius change going down a group?

A

Increases

59
Q

What are isoelectric ions?

A

Ions of different atoms with the same numbers of electrons

60
Q

What is proof for ionic lattices having charged particles?

A

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
Q

What is the bond length?

A

Distance between two nuclei in covalent bonding

62
Q

How does bond length affect the bond enthalpy?

A

The shorter the covalent bond, the higher the bond enthalpy

63
Q

What is electron density in relation to covalent bonding?

A

The number of electrons within the covalent bond

64
Q

How does electron density of covalent bonds affect bond enthalpy?

A

Higher the electron density
Shorter the bond
Therefore higher the enthalpy

65
Q

How much do lone pairs repel in comparison to bond pairs?

A

Lone pairs repel more than bonding pairs

66
Q

Between which electron pairs are the angles largest?

A

Lp/lp is largest angles
Lp/bp is second largest angles
Bp/bp is smallest angels

67
Q

How can lone pair repulsion affect bonding pairs?

A

Bonding pair angles reduced as they are pushed together by lone pair repulsion

68
Q

What halogen bonded to hydrogen has the shortest bond?

A

H - F

most electronegative

69
Q

Which of the following has the largest ionic radius?

S 2-, Cl -, K +, Ca 2+

A

S 2-

Same electrons but less protons so not as strong electrostatic pull on electrons

70
Q

Why does H2O have a bond angle of 104.5?

A

Electrons repel as far as possible
Lone pair has greater repulsion than bond pair
Therefore less than 109

71
Q

What is an ionic bond?

A

Electrostatic attraction between oppositely charged ions

72
Q

What is an ionic giant lattice?

A

3-D arrangement of ions which is regular

Alternating between positive and negative

73
Q

What properties do ionic compounds have?

A

Hard crystalline structures
High melting and boiling points
Generally high solubility in water
Good electrical conductivity when molten/dissolved

74
Q

Why do ionic compounds not conduct electricity as a solid?

A

Ions cannot move out of the regular shape, therefore it cannot conduct electricity

75
Q

What is a covalent bond?

A

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

76
Q

What is a simple molecule?

A

Non-metals bonded together covalentlyUsually diatomic

77
Q

What are the properties of simple covalent molecules?

A

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
Q

What is a giant covalent structures?

A

Many atoms joined together by very strong covalent bonds

79
Q

What is the structure of diamond?

A

Each C atom bonded to 4 others

Tetrahedral 3D arrangement

80
Q

What is the structure of graphite?

A

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
Q

What keeps layers of graphite together?

A

London/van der Waals forces

82
Q

What is graphene?

A

1 layer of graphite
Hexagon shape
Good electrical conductor
Relatively strong

83
Q

What is the structure of metals?

A

Giant metallic 3D lattices

Metal cations surrounded by a sea of deloacalised electrons, attracted by electrostatic forces

84
Q

What does the structure of metals cause?

A

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
Q

Why do ionic lattices form?

A

Each ion is electrostatically attracted in all directions to ions of the opposite charge

86
Q

Why can ionic compounds not be shaped?

A

Same charged ions repel each other, so they break when its stretched or hammered Brittle

87
Q

Why are ionic substances soluble in polar solvents?

A

Ions are pulled apart by polar molecules

88
Q

What are the properties and explanations of giant covalent structures?

A

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
Q

Why can graphite conduct electricity?

A

As only bonds to 3 other carbon atoms, there is one delocalised electron per atom
Electrons can carry the charge

90
Q

What bonding occurs in metallic bonding?

A

Electrostatic attraction between delocalised negative electrons and the positive metal ions

91
Q

How does the number of electrons affect metallic bonding?

A

Stronger bonding as there is a larger electrostatic force

92
Q

How do impurities affect the electrical conductivity of metals?

A

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)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions