Brainscape's Knowledge GenomeTM

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


See full index

A-Level Chemistry > Mod 2 Chapter 6 > Flashcards

Mod 2 Chapter 6 Flashcards

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

Electron pair repulsion theory

A

Model used in chemistry to explain and predict the shapes of molecules and poly atomic ions - electron pairs surrounding central atom determine the overall shape and the pairs of electrons repel each other so they are as far apart as possible to minimise repulsion (holding the bonded atoms in a definite shape)

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

4 bonded regions

A

CH4 - methane
Repel SS far apart as possible (3D)
Tetrahedral shape 4 equal C-H bond angles of 109.5

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

How to represent 3D shapes on a 2D plane?

A

Solid line - same plane as the paper
Dotted wedge - going into the paper
Solid wedge - coming out of the paper
HELPS TO VISUALISE THE STRUCTURES

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

Why do lone pairs repel more than bonded pairs?

A

They are slightly closer to the central atom and occupy more space therefore have a larger negative charge thus repelling more strongly than a bonded pair

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

Ammonia

A 
NH3
3 bonded regions
1 lone pair
Pyramidal ; 109.5-2.5 = 107 degrees
1 dotted, 1 solid w, 1 solid l
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Water

A 
H2O
2 bonded regions
2 lone pairs
Non-linear/bent
109.5-2.5-2.5 = 104.5 degrees
1 solid line and 1 solid wedge
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What do lone pairs do?

A

Repel bonded pairs CLOSER TOGETHER THUS DECREASING THE BOND ANGLE (angle between the bonded pairs of electrons)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What to do if the molecule contains multiple bonds?

A

Each set of bonds are treated as a bonding region - does not make a difference

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

Carbon dioxide

A 
CO2
2 double bonds (bonding regions)
Linear 
180 degrees
REPEL AS FAR APART AS POSSIBLE - gives it a linear shape with all 3 atoms in a straight line
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Boron trifluoride

A 
BF3
3 bonded pairs around central
No lone pairs
Maximum repulsion = 360/3 = 120 degrees
Trigonal planar
All 3 dotted lines
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Sulfur hexafluoride

A 
6 bondein regions
Oxtahedral shape (as each atom acts as the corner of an octahedron)
All 90 degrees
2 solid lines at top and bottom
2 dashed lines top right and left
2 solid wedges bottom right and left
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Ammonium ion

A

NH4+ ; 3 single bonds and 1 dative bond
Tetrahedral shape 109.5 degrees
Same 3D shape as CH4
Surrounded by ionic brackets

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

CO32-

A 
3 religions of electron density
NOTE ON NO3- DATIVE BOND FORMED INSTEAD OF DOUBLE BOND
No lone pairs
120 degrees trigonal planar
Both drawn flat
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

So42- ions

A

4 regions of electron density
109.5 degrees
Tetrahedral
No lone pairs

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

What does electron pair repulsion theory allow us to do?

A

Predict the arrangement of electron pairs surrounding the central atom

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

SO2

A

2 double bonds
1 lone pair
Extra electron density of double bonds cancels out the extra repulsion of lone pair creating a trigonal planar 120 degrees drawn flat (basically taking the lone pair as another bond)

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

What is a covalent bond?

A

Electrostatic attraction between hthe nucleus of the bonded atoms and the shared pair of electrons

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

Describe pure molecules of elements

A

All the diatomic molecules are the same element and therefore attract the bonded electrons evenly (shared equally) - PURE covalent bond

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

Difference when different atoms are bonded together

A

One atom (nucleus of the atom) will attract the shared electrons more strongly than the other

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

Factors that electro negativity depends on

A

Nuclear charges of the nucleus - more protons means larger positive charge thus more attraction
Radius/size of atom - if radius is smaller then it is more electronegative as there is a smaller distance to attract the electrons and thus less shielding (no extra repelling of electrons)

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

Electronegative thank definition

A

The attraction of a bonded atom for the pair of electrons in a covalent bond (measure of the atom’s ability to attract electrons)

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

Charge density

A

Charge density takes into account both the nuclear charge (atomic number) relative to the size of the atom - even if nucleus is really positive if atomic radius is too large then no difference made

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

Pauling scale

A

Used to compare the electrongetativty values of the atoms of different elements

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Trends on the Pauling scale

A

As you move across the table (groups) - the nuclear charge increases thus electronegativity also increases (number of protons)
As you move up the groups the electronegativity increases as the atomic radius decreases

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

Fluorine

A

Most electronegative - 4 - perfect combination of soze and nuclear charge

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

Which one is more effective?

A

Going up (reducing size of atom) has more effect than going across (just adding one more proton)

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

Describe parts of the scale

A

Non metals - nitrogen oxygen and fluorine are the most electronegative
Group 1 elements have the least electronegative

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

What can Pauling electronegative scale be used for?

A

To make predictions about the type of bonding - if you have a electronegativity difference of 0 = pure covalent
Electronegativity difference of 0-1.8 = polar covalent
Electronegative difference of >1.8 = ionic
If less than 0.5 (such as C-H) it can be ignored

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

If electronegativity difference is too large…

A

One bonded atom will have a much greater attraction to the electrons than the other - thus would have taken over and formed an ionic bond (transfer of electrons)

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

Non polar bond

A

Bonded electron pair is shared equally between the bonded atoms - this occurs in a pure covalent bond (both atoms the same) and when bonded atoms have the same/similar electronegativity
C-H is non-polar as they ave similar electronegativity values ; all hydrocarbons are non polar solvents and do not mix with water

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

Polar bonds

A

Where the bonded electron pair is shared unequally between the bonded atoms - a bond will be polar when the bonded atoms are different and have different electronegativity values ; polar covalent

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

H-Cl

A

H - 2.1
Cl - 3.0
Cl is more electronegative than hydrogen
Chlorine has a greater attraction for the bonded pair of electrons than the hydrogen atom (resulting in a polar covalent)
POLARISED BOND - small partial charge on hydrogen and negative charge on chlorine (delta = small)
Partial charges are much smaller than actual ionic charges

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

What is a dipole?

A

The separation of opposite partial charges within a polar covalent molecule

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

Permanent dipole

A

Small charge difference that does not change across a bond with delta+ and delta- partial changes on the bonded atmosphere - result of the bonded atoms having different electronegativity values

35
Q

Polar molecules

A

Depending on the shape of the molecule the polar bonds may cancel each other out (if acting in opposite directions) or reinforce the polarity

36
Q

H20 polar molecule?

A

Two O-H bonds have a permanent dipole BUT they act in different directions (but not opposite)g therefore the oxygen end retains the delta negative charge and the hydrogen end is positive - therefore the molecule is polar

37
Q

CO2 molecule polar?

A

Two C=O bonds have a permanent dipole but they act in opposite directions and opposite each other therefore cancelling each other out and creating a non polar molecule

38
Q

Polar solvents and solubility

A

Any ionic compounds break down in polar water with the positive ion attracted towards the oxygen and the negative end towards the hydrogen ; breaks down/dissolves in water

39
Q

Intermolecular forces definition

A

They are weak interactions between dipoles of different molecules

40
Q

3 main intermolecular forces categories

A 
London forces (induced dipole-dipole interactions)
Permanent dipole-dipole interactions
Hydrogen bonding
41
Q

Physical vs chemical properties

A

Intermolecular forces are responsible for physical properties such as melting/boiling points - whereas covalent bonds WITHIN molecule determine the identity/chemical reactions of molecules

42
Q

Bond enthalpy

A 
The amount of energy required to break a bond
London forced (least energy)
Hydrogen bonds (most)
Covalent bonds the highest
43
Q

London forces

A

Weak intermolecular forces that exist between ALL molecules whether polar or non polar
An instantaneous dipole is produced by the random movement of electrons - creating dipoles that then induce dipoles onto neighbouring molecules which continues on in a chain (position of dipoles constantly shifting) ; causes weak attraction between molecules
These induced dipoles are only temporary - next instant of time these may disappear causing the process to take place amongst other molecules

44
Q

As size of molecule increases

A

More electrons
Stronger instantaneous dipoles
Larger induced dipoles
Stronger the attractive forces between molecules
More energy needed to overcome these intermolecular London forces thus increasing the boiling point

45
Q

Van der Waal’s forces

A

Takes into account both London forces and permanent dipole dipole interactions

46
Q

Permanent dipole dipole interactions

A

Act as IMF between polar molecules with permanent dipoles

47
Q

Compare F2 and HCl

A

Fluorine molecules are non-polar (no difference in electronegativity) therefore only have London forces
Hydrogen chloride molecules are polar and have London forces and pd-pd interactions
THEREFORE extra energy required to break the additional Pd-pd interactions between HCl molecules - higher BP than F2

48
Q

Simple molecular substances

A

These consist of the intermolecular forces and are made up of molecules which have a definite number of atoms/molecular formula
In solid state - they are held together in a regular lattice and IMF such as London forces and (if polar) pd-pd interactions hold the molecules in place
STRONG COVALENT BONDS HOLD ATOMS WITHIN MOLECULE IN PLACE

49
Q

All simple molecular substances are….

A

COVALENTLY BONDED - room temperature they may exist as solids/liquids/gases ; may be solidified into a simple molecular lattice by slightly reducing the temperature

50
Q

IMF in simple molecular lattice

A

They can be easily broken at low temperatures - very little energy required ; low melting and boiling points
iMF are broken apart not the covalent bonds

51
Q

Solubility of non polar simple molecular substances

A

Non polar + non polar forms intermolecular forces between them (London forces) and these weaken the intermolecular forces in the lattice thus breaking the IMF and allowing compound to dissolve

52
Q

Non polar solubility in polar

A

Little interaction between molecules in lattice and solvent therefore intermolecular bonding within polar solvent is too strong to be broken so simple molecular substances are insoluble in polar
Non polar molecules only create induced dipole dipole interactions - they need Pd-pd to dissolve in polar solvents like water

53
Q

Solubility of polar simple molecular substances

A

Polar covalent substances dissolve in polar solvents as they can form Pe-pd forces ; they attract each other

54
Q

Dissolving of an ionic compound

A

Sugar dissolves in water - sugar is a polar covalent compound with many O-H bonds which attract and bond with polar water molecules ; this can also occur to liquids and gases (H-Cl gas dissolves in water to create Hydrochloric acid)

55
Q

Factor solubility depends on

A

Strength of dipole

56
Q

C2H5OH

A

Contains both polar (O-H) and non-polar (carbon) parts so can dissolve in both polar and non-polar solvents

57
Q

Hydrophilic part of biological molecules

A

POLAR (contain electronegative atoms - O2) that can interact with oxygen

58
Q

Hydrophobic part

A

Non-polar (carbon chain)

59
Q

Are simple molecular substances conductor?

A

There are no mobile charged particles therefore there is nothing to complete the circuit so they are NON CONDUCTORS OF ELECTRICITY

60
Q

Hydrogen bond

A

It is a type of permanent dipole dipole interactions between a lone pair of electrons (on an electronegative atom) and a hydrogen atom of an electronegative atom

61
Q

Hydrogen bond strength

A

Strongest type of intermolecular attractions

62
Q

Shape around hydrogen atom in a hydrogen bond

A

Always linear - 180 degrees

63
Q

Anomalous properties of water

A

Ice is less dense than the liquid - hydrogen bonds hold water molecules apart in an open lattice structure and the water molecules in ice are further apart than in water so solid ice is less dense than liquid water and floats

64
Q

Ice is less dense than water

A

So it floats forming an insulating layer and preventing water below from freezing solid

65
Q

Why is water perfect for hydrogen bonding?

A

Perfect ratio of lone pairs to hydrogen atoms - can form 4 hydrogen bonds (2 lone pairs + 2 hydrogen atoms)

66
Q

What does water form in the solid state?

A

The hydrogen bonds extend outwards holding water molecules apart in a tetrahedral lattice full of holes - hydrogen bond angle involved in bonding = 180 degrees

67
Q

What does tetrahedral lattice mean?

A

Decrease the density of water on freezing (larger volume) - when ice melts the ice lattice collapse and molecules move closer together

68
Q

Anomalous property of water 2

A

Relatively high melting point and boiling point

69
Q

Why does water have high MP and BP?

A

Hydrogen bonds are extra forces on top of London forces (which exist around all molecules) therefore water needs a lot of energy to break the extra hydrogen bonds ; when ice lattice breaks the rigid arrangement of hydrogen bonds collapse and when water boils the hydrogen bonds break completely

70
Q

Without hydrogen bonds…

A

Only have London forces therefore water would have a very low boiling point (exist as a gas(

71
Q

Other anomalous properties

A

Surface tension

Viscosity

72
Q

Other anomalous properties

A

Surface tension

Viscosity

73
Q

Without hydrogen bonds…

A

Only have London forces therefore water would have a very low boiling point (exist as a gas(

74
Q

Why does water have high MP and BP?

A

Hydrogen bonds are extra forces on top of London forces (which exist around all molecules) therefore water needs a lot of energy to break the extra hydrogen bonds ; when ice lattice breaks the rigid arrangement of hydrogen bonds collapse and when water boils the hydrogen bonds break completely

75
Q

Anomalous property of water 2

A

Relatively high melting point and boiling point

76
Q

What does tetrahedral lattice mean?

A

Decrease the density of water on freezing (larger volume) - when ice melts the ice lattice collapse and molecules move closer together

77
Q

What does water form in the solid state?

A

The hydrogen bonds extend outwards holding water molecules apart in a tetrahedral lattice full of holes - hydrogen bond angle involved in bonding = 180 degrees

78
Q

Why is water perfect for hydrogen bonding?

A

Perfect ratio of lone pairs to hydrogen atoms - can form 4 hydrogen bonds (2 lone pairs + 2 hydrogen atoms)

79
Q

Ice is less dense than water

A

So it floats forming an insulating layer and preventing water below from freezing solid

80
Q

Anomalous properties of water

A

Ice is less dense than the liquid - hydrogen bonds hold water molecules apart in an open lattice structure and the water molecules in ice are further apart than in water so solid ice is less dense than liquid water and floats

81
Q

Shape around hydrogen atom in a hydrogen bond

A

Always linear - 180 degrees

82
Q

Hydrogen bond strength

A

Strongest type of intermolecular attractions

83
Q

Hydrogen bond

A

It is a type of permanent dipole dipole interactions between a lone pair of electrons (on an electronegative atom) and a hydrogen atom of an electronegative atom

A-Level Chemistry (28 decks)

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