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A-level Chemistry AQA > 3.1.3 Bonding (physical chemistry) > Flashcards

3.1.3 Bonding (physical chemistry) Flashcards

1
Q

define metallic bonding

A

The strong electrostatic attraction between positive metal ions (cations) and a sea of delocalised electrons

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2
Q

draw a diagram for the bonding in:
-sodium
-aluminium
-magnesium

A
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3
Q

explain why the bonding in magnesium is stronger than the bonding in sodium

A

-Mg2+ ions have a higher charge than Na+ ions
-Magnesium has twice as many delocalised electrons per metal ion than sodium
-Mg2+ ions have a smaller ionic radius so have a higher charge density
-this means the attraction between the metal cations and the delocalised e- is stronger

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4
Q

4 properties of metals

A

1.conductivity
-good electrical and thermal conductors as the delocalised e- help transfer energy as they flow through the metal
2.strength
-very strong as there is a strong electrostatic attraction between cations and delocalised e-
3. malleable and ductile
-layers of metal ions can slide past one another
4.melting and boiling points
-directly linked to strength of metallic bond
-stronger the bonds, higher the melting point and boiling point

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5
Q

definition of covalent bonding

A

shared pair of electrons between 2 atoms

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6
Q

5 macromolecular structures

A

1.Diamond
2.Graphite
3.Graphene
4. Silicon
5.Silicon dioxide

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7
Q

allotropes

A

different structural forms of the same element

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8
Q

describe structure and bonding of diamond

A

-Each carbon atom is bonded to 4 other carbons by strong covalent bonds
-To melt, many strong covalent bonds need to be broken
-This requires alot of energy to overcome
-Melting point is very high
-Diamond is very strong
-No delocalised e- so doesnt conduct electricity

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9
Q

describe structure and bonding of graphite

A

-Each carbon is covalently bonded to 3 other carbon atoms - strong covalent bonds
-1 delocalised e-/ carbon atom and so can conduct heat and electricity as it can move through the structure
-Graphite is arranged in layers held together by weak intermolecular forces
-These layers can slide over eachother making graphite a soft structure

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10
Q

describe structure and bonding of graphene

A

-a single layer of graphite
-conducts electricity as it has delocalised e-
-very strong and has strong covalent bonds between atoms

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11
Q

simple molecular structures forces and properties

A

-low melting points and boiling points
-only relatively weak IMF need to be broken
-no delocalised electrons or ions so not conductors of electricity

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12
Q

Ionic bonding definition

A

The strong electrostatic attraction between oppositely charged ions
-e- are always transferred from the metal to the non-metal
-the metal will always form a cation and the non metal an anion

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13
Q

physical properties of ionic compounds

A

-high melting points
-strong electrostatic attraction between ions requires lots of energy to overcome

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14
Q

conductivity of ionic compounds

A

-as a solid no delocalised e- or mobile ions so doesnt conduct electricity
-when molten or aqueous ions mobile and can easily carry charge so conducts electricity

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15
Q

formula of ammonium

A

[NH4]+

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16
Q

formula of hydroxide

A

[OH]-

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17
Q

formula of nitrate

A

[NO3]-

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18
Q

formula of nitrite

A

[NO2]-

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19
Q

formula of hydrogencarbonate

A

[HCO3]-

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20
Q

formula of chlorate (I)

A

[ClO]-

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21
Q

formula of chlorate (V)

A

[ClO3]-

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22
Q

formula of carbonate

A

[CO3]2-

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23
Q

formula of sulfate

A

[SO4]2-

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24
Q

formula of sulfite

A

[SO3]2-

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25
Q

formula of dichromate

A

[Cr2O7]2-

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26
Q

formula of phosphate

A

[PO4]3-

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27
Q

define molecule

A

a group of atoms which are covalently bonded to one another

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28
Q

define coordinate bond

A

a shared electron pair which have both come from the same atom

29
Q

describe a coordinate bond

A

-coordinate bonds are always represented by an arrow
-the direction of the arrow points from the atom which donates the pair of electrons to the atom which accepts the pair
-once a coordinate bond has been formed it behaves in exactly the same way as a covalent bond i.e same length and strength

30
Q

electron repulsion strength

A

lone to lone pair>lone to bond pair>bond to bond pair
-the strength of the repulsions determines the bond angles between the bond to bond pair

31
Q

3 bonding pairs

A

triganol planor
120 degrees

32
Q

2 bonding pairs
1 lone pair

A

V-shaped
117.5 degrees

33
Q

4 bonding pairs

A

Tetrahedral
109.5 degrees

34
Q

3 bonding pairs
1 lone pair

A

triganol pyramidal
107 degrees
replace e with lone pair

35
Q

2 bonding pairs
2 lone pairs

A

V-shaped
104.5 degrees
replace e with lone pair

36
Q

5 bonding pairs

A

triganol bipyramidal
90 degrees
120 degrees

37
Q

4 bonding pairs
1 lone pair

A

see saw
87.5 degrees
117.5 degrees
replace e with lone pair

38
Q

3 bonding pairs
2 lone pairs

A

triganol planor
120 degrees

39
Q

6 bonding pairs

A

octahedral
90 degrees

40
Q

5 bonding pairs
1 lone pair

A

square pyramidal
87.5 degrees

41
Q

4 bonding pairs
2 lone pairs

A

square planor
90 degrees

42
Q

define electronegativity

A

the power of an atom to attract the pair of electrons in a covalent bond towards itself

43
Q

how is electronegativity measured

A

-the pauling scale
-each element assigned a number between 0 and 4
-the higher the number the more electronegative it is

44
Q

factors which affect how electronegative an element is

A

1.nuclear charge, as nuclear charge increases so does electronegativity
2. shielding, as shielding increases electronegativity decreases
3. atomic radius, as atomic radius decreases, electronegativity increases
(as a rule of thumb, the closer an element to fluorine, the more electronegative it is)

45
Q

explain electronegativity across period 2

A

-nuclear charge increases across the period
-atomic radius decreases across the period
-electronegativity increases across the period
-nuclear attraction for the shared pair of electrons increases

46
Q

polarity between Cl-Cl

A

-molecules made of atoms with no difference in electronegativity have their electrons distributed evenly

47
Q

polarity between H-Cl

A

-The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical
-this produces a polar covalent bond
-the atom which has a greater electron density has a delta- charge and the atom which has a lower electron density has a delta+ charge
-H is delta positive and Cl delta negative

48
Q

is BF3 polar

A

-all bonds polar
BUT
-molecule is symmetrical
-dipoles cancel out
-molecule not polar

49
Q

is NH3 polar

A

-all bonds are polar
-molecule is unsymmetrical
-dipoles do not cancel out
-molecule is polar

50
Q

is CHCl3 polar

A

-all bonds are polar
-molecule is unsymmetrical
-dipoles do not cancel out
-molecule is polar

51
Q

what makes a molecule symmetrical

A

-all outer atoms the same
-triganol planor
-linear
-tetrahedral
-triganol bipyramidal
-octahedral
-square planor

52
Q

three types of intermolecular forces

A

1.hydrogen bonding
2.permanent dipole dipole
3.induced dipole dipole

53
Q

when hydrogen bonding occurs

A

-intermolecular force
-lone pair on N,O or F
H-F bond
H-O bond
H-N bond

54
Q

when permanent dipole dipole intermolecular forces occur

A

-present in polar molecules

55
Q

when induced dipole dipole intermolecular forces occur

A

-non-polar molecules
-The more electrons, the stronger the force
-between all molecules and group 18 elements

56
Q

how does hydrogen bonding arise

A

-there is a very large difference in electronegativity between H and N/O/F
-this creates a large dipole on the bond
-The delta positive hydrogen strongly attracts a lone pair of electrons on the delta negative N/O/F in another molecule

57
Q

how to draw hydrogen bonding

A
  1. Draw all lone pairs and dipoles
  2. Dashed hydrogen bond between delta + H and delta - O from another molecule
  3. Bond angle must be 180 degrees
58
Q

how does permanent dipole dipole forces arise

A

-differences in electronegativity leads to polar bonds
-if dipoles do not cancel out due to symmetry there is an overall dipole
-delta positive on one molecule is attracted to the delta negative on another molecule

59
Q

how to draw permanent dipole dipole forces

A
60
Q

how does induced dipole dipole forces arise

A

-random movement of electrons leads to an uneven electron distribution and a transient/instantaneous/temporary dipole
-this induces a dipole in a neighbouring atom or molecule
-delta positive and delta negative dipoles attract one another

61
Q

how to draw induced dipole dipole forces

A
62
Q

hydrogen bonding in ice

A

-ice is less dense than liquid water
-2 hydrogen bonds per molecule
-holds molecules together in an open lattice structure
-more space between molecules than liquid water

63
Q

hydrogen bonding in proteins

A

-proteins are held in a complex 3D shape by Hydrogen bonds between different parts of the chain
-3D shape is important for a functional protein

64
Q

hydrogen bonding in DNA

A

-2 strands of DNA held together in a double helix by hydrogen bonds

65
Q

describe the general trend in melting and boiling points across period 3
Na to Mg to Al

A

-Na, Mg, Al are all giant metallic lattices
-Electrostatic attraction between positive ions and delocalised electrons
-Ionic charge increases from Na to Mg to Al (+1,+2,+3)
-Number of delocalised electrons increases
-Metallic bonding is stronger
-More energy is required to overcome this

66
Q

describe the general trend in melting and boiling points across period 3
Silicon

A

-Silicon is a macromolecular structure with strong covalent bonds between atoms
-Lots of energy is required to overcome the many strong covalent bonds

67
Q

describe the general trend in melting and boiling points across period 3
P to S to Cl to Ar

A

-P to S to Cl to Ar are all simple molecular structures
-They have weak induced dipole dipole forces between molecules
P4 contains 60 e-
S8 contains 128 e-
Cl2 contains 34 e-
Ar contains 18 e-
-The more electrons there are the stronger the induced dipole dipole forces between molecules and so the more energy required to overcome this
-The melting points are much lower than Na to Si as induced dipole dipole forces are much weaker than metallic and covalent bonds
-These require less energy to overcome

68
Q

Explain why Silicon has the highest melting point but Aluminium has the highest boiling point

A

-When Al is molten there is still attractions between cations and delocalised electrons
-To vaporise it, strong metallic bonds still need to be broken
-Alot of energy is required to overcome this
-When Si is molten the covalent bonds are already broken
-Less energy is required to vaporise it

A-level Chemistry AQA (5 decks)

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