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A-Level Chemistry > Mod 2 Chapter 6 > Flashcards

Mod 2 Chapter 6 Flashcards

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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)

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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

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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

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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

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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
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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
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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)

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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

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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
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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
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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
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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

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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
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14
Q

So42- ions

A

4 regions of electron density
109.5 degrees
Tetrahedral
No lone pairs

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15
Q

What does electron pair repulsion theory allow us to do?

A

Predict the arrangement of electron pairs surrounding the central atom

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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)

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17
Q

What is a covalent bond?

A

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

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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

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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

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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)

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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)

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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

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23
Q

Pauling scale

A

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

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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

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25
Fluorine
Most electronegative - 4 - perfect combination of soze and nuclear charge
26
Which one is more effective?
Going up (reducing size of atom) has more effect than going across (just adding one more proton)
27
Describe parts of the scale
Non metals - nitrogen oxygen and fluorine are the most electronegative Group 1 elements have the least electronegative
28
What can Pauling electronegative scale be used for?
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
29
If electronegativity difference is too large…
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)
30
Non polar bond
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
31
Polar bonds
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
32
H-Cl
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
33
What is a dipole?
The separation of opposite partial charges within a polar covalent molecule
34
Permanent dipole
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
Polar molecules
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
H20 polar molecule?
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
CO2 molecule polar?
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
Polar solvents and solubility
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
Intermolecular forces definition
They are weak interactions between dipoles of different molecules
40
3 main intermolecular forces categories
``` London forces (induced dipole-dipole interactions) Permanent dipole-dipole interactions Hydrogen bonding ```
41
Physical vs chemical properties
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
Bond enthalpy
``` The amount of energy required to break a bond London forced (least energy) Hydrogen bonds (most) Covalent bonds the highest ```
43
London forces
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
As size of molecule increases
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
Van der Waal’s forces
Takes into account both London forces and permanent dipole dipole interactions
46
Permanent dipole dipole interactions
Act as IMF between polar molecules with permanent dipoles
47
Compare F2 and HCl
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
Simple molecular substances
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
All simple molecular substances are….
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
IMF in simple molecular lattice
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
Solubility of non polar simple molecular substances
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
Non polar solubility in polar
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
Solubility of polar simple molecular substances
Polar covalent substances dissolve in polar solvents as they can form Pe-pd forces ; they attract each other
54
Dissolving of an ionic compound
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
Factor solubility depends on
Strength of dipole
56
C2H5OH
Contains both polar (O-H) and non-polar (carbon) parts so can dissolve in both polar and non-polar solvents
57
Hydrophilic part of biological molecules
POLAR (contain electronegative atoms - O2) that can interact with oxygen
58
Hydrophobic part
Non-polar (carbon chain)
59
Are simple molecular substances conductor?
There are no mobile charged particles therefore there is nothing to complete the circuit so they are NON CONDUCTORS OF ELECTRICITY
60
Hydrogen bond
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
Hydrogen bond strength
Strongest type of intermolecular attractions
62
Shape around hydrogen atom in a hydrogen bond
Always linear - 180 degrees
63
Anomalous properties of water
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
Ice is less dense than water
So it floats forming an insulating layer and preventing water below from freezing solid
65
Why is water perfect for hydrogen bonding?
Perfect ratio of lone pairs to hydrogen atoms - can form 4 hydrogen bonds (2 lone pairs + 2 hydrogen atoms)
66
What does water form in the solid state?
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
What does tetrahedral lattice mean?
Decrease the density of water on freezing (larger volume) - when ice melts the ice lattice collapse and molecules move closer together
68
Anomalous property of water 2
Relatively high melting point and boiling point
69
Why does water have high MP and BP?
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
Without hydrogen bonds…
Only have London forces therefore water would have a very low boiling point (exist as a gas(
71
Other anomalous properties
Surface tension | Viscosity
72
Other anomalous properties
Surface tension | Viscosity
73
Without hydrogen bonds…
Only have London forces therefore water would have a very low boiling point (exist as a gas(
74
Why does water have high MP and BP?
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
Anomalous property of water 2
Relatively high melting point and boiling point
76
What does tetrahedral lattice mean?
Decrease the density of water on freezing (larger volume) - when ice melts the ice lattice collapse and molecules move closer together
77
What does water form in the solid state?
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
Why is water perfect for hydrogen bonding?
Perfect ratio of lone pairs to hydrogen atoms - can form 4 hydrogen bonds (2 lone pairs + 2 hydrogen atoms)
79
Ice is less dense than water
So it floats forming an insulating layer and preventing water below from freezing solid
80
Anomalous properties of water
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
Shape around hydrogen atom in a hydrogen bond
Always linear - 180 degrees
82
Hydrogen bond strength
Strongest type of intermolecular attractions
83
Hydrogen bond
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)

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