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A Level Chemistry > Module 2 - Bonding And Structure > Flashcards

Module 2 - Bonding And Structure Flashcards

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1
Q

what is ionic bonding

A

electrostatic attraction between positive and negative ions

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2
Q

describe how ions are formed in ionic bonding

A

the transfer of electrons from metal atoms to non metal atoms from the outer electrons ( the ones in the highest occupied principal energy level )

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3
Q

why do ionic compounds have high melting and boiling points

A
  • ionic bonds are strong and a lot of heat is needed to break them (high energy)
  • giant ionic lattice
  • strong electrostatic attraction between + and - ions
  • larger ionic charges produce stronger ionic bonds, so more heat is required to break the ionic bonds
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4
Q

how do ionic compounds conduct electricity

A
  • as solids, can’t conduct because their ions are bonded together in the lattice
  • when molten, the ions can break free of the lattice and are able to move.
  • the ions are charged particles so can carry an electric current
  • usually soluble in water because water molecules have a slight electrical charge so can attract ions away from the lattice
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5
Q

what is covalent bonding

A

the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

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6
Q

why do simple covalent molecules have low boiling/ melting points

A

-covalent bonds within the molecules (intra molecular forces) are strong, but the forces of attraction between molecules (intermolecular forces) are weak
- not much energy is needed to overcome these forces of attraction

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7
Q

why don’t simple covalent molecules conduct electricity

A
  • no ions or free electrons present
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8
Q

why do giant covalent molecules have high boiling/ melting points

A
  • high energy needed to break because
    diamond- 4 strong covalent bonds between every c atom
    graphite - 3 strong covalent bonds between every c atom
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9
Q

why do giant covalent molecules conduct electricity

A
  • delocalised electrons which move to carry charge through the structure
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10
Q

when is a dative covalent bond formed

A

when one atom contributes both of the electrons needed for the covalent bond

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11
Q

what is needed for a dative covalent bond to form

A

one atom has to have a lone pair of electrons and the other atom must have a vacant orbital

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12
Q

how is a dative covalent bond represented

A

as an arrow, which shows the direction of the electron pair donation

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13
Q

what is a metallic bond

A

the electrostatic force of attraction between a lattice of positively charged ions and delocalised electrons

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14
Q

factors affecting strength of a metallic bond

A
  • size of the charge on the positive ions- the bigger, the stronger
  • the number of mobile electrons - the more, the stronger
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15
Q

why are metals good conductors of electricity when solid

A

delocalised electrons can carry charge through the structure

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16
Q

why do metals have high melting and boiling points

A
  • giant lattice
  • strong electrostatic attraction between the metal ions and delocalised electrons
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17
Q

define electronegativity

A

the ability of an atom to attract the pair of electrons in a covalent bond (electron density)

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18
Q

how is electronegativity measured

A

on the Pauling scale

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19
Q

what is the most electronegative on the Pauling scale

A

F (4.0)

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20
Q

what is the least electronegative on the Pauling scale

A

Fr (0.7)

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21
Q

describe how electronegativity changes on the Pauling scale

A

decreases left across the period
decreases down the group

22
Q

what are factors that determine the electronegativity of an atom

A

the size of the nuclear charge
the size of an atom

23
Q

describe how the size of the nuclear charge affects electronegativity

A

the bigger this is, the larger the attraction between the nucleus and the pair of electrons
the electronegativity goes up

24
Q

describe how the size of an atom affects electronegativity

A

as this increases, the pair of electrons are further from the nucleus, and there will be a shielding effect from the inner electrons
the electronegativity goes down

25
Q

explain non polar covalent compounds

A

if both of the atoms have equal electronegativity they will both have the same tendency to attract electrons
the bond is formed roughly halfway between the 2 atoms

26
Q

explain the bond for atoms with slightly different electronegativity values

A

if atom A is slightly more electronegative than atom B, A has greater attraction to the electron pair
the electron pair will be pulled towards atom a
the a end of the bond will have more electron density and will become slightly negatively charged
the b end of the bond will have less electron density and will become slightly positively charged

27
Q

what does the polarity of molecules depend on

A

it’s overall shape

28
Q

what shape makes a molecule non polar

A

molecules with overall symmetry
as the dipoles cancel out because overall the sharing of electrons is equal

29
Q

describe london forces

A

the weakest intermolecular attraction
form between neighbouring non polar molecules
only have an affect over a couple of nanometers

30
Q

describe the formation of london forces

A

electrons in an atom constantly move. at any time one end of the molecule might have low electron density compared to the other. A temporary dipole has been formed
the S+ end of the molecule will attract electron density to one end of a neighbouring molecule
the neighbouring molecule now has a dipole that has been induced by the first molecule - an induced dipole

31
Q

what happens to london forces as the molecules get bigger

A

more surface contact
more electrons
stronger london forces

32
Q

describe permanent dipole dipole forces

A

10x stronger than london forces
happen at the same time as london forces
form between two polar covalent molecules
force of attraction between opposite charged on neighbouring molecules
leads to highs boiling points compared with similar sized compounds that just have london forces

33
Q

describe the formation of permanent dipole dipole forces

A

the highly electronegative atom in a molecule is attracted towards the S+ atom in a neighbouring molecule as the S- tries to attract electron density towards itself

34
Q

describe hydrogen bonding

A

only happens when a H atom is covalently bonded to an N, O or F atom (very electronegative)
the strongest intermolecular force of attraction as 100x stronger than london forces
ionic bonds are 100x stronger than H bonds

35
Q

describe the formation of hydrogen bonds

A

a H atom is covalently bonded to an N, O or F atom
the bond becomes so polarised that the S+ H atom can also form a weak bond with another F, O or N atom in a neighbouring molecule

36
Q

steps for drawing hydrogen bonds

A
  • always show lone pairs
  • always show the polarity
  • H bond always from a lone pair
37
Q

what does the shape of a molecule depend on

A

the number of electron pairs around the central atom in the molecule

38
Q

describe the properties of a linear molecule

A

bond pairs - 2
lone pairs - 0
bond angle- 180

39
Q

what do double bonds count as when working out shapes of molecules

A

1 pair of bonded electricity

40
Q

properties of a trigonal planar molecule

A

bond pairs - 3
bond angle - 120

41
Q

properties of a tetrahedral molecule

A

bond pairs -4
bond angle- 109.5

42
Q

properties of a trigonal bipyramidal molecule

A

bond pairs -5
bond angle - 120 and 90

43
Q

properties of an octahedral molecule

A

bond pairs- 6
bond angle - 90

44
Q

when are irregular shaped molecules formed

A

if a molecule or ion has lone pairs on the central atom
as the shapes are slightly distorted away from the regular shapes
as there is extra repulsion caused by the lone pairs

45
Q

properties of a pyramidal molecule

A

bond pairs-3
lone pairs - 1
total pairs-4
angle-107

46
Q

properties of a v shaped/ non linear molecule

A

bond pairs-2
lone pairs-2
total pairs-4
angle-104.5

47
Q

steps for drawing shapes of ions

A
  1. draw outer shell electrons of central atom
  2. look at charge and either add or remove an electron/ electrons
  3. pair up the electrons
  4. work out shape and bond angles
48
Q

what are the properties of water

A

high boiling/ melting point for a simple molecule
ice less dense than water

49
Q

why does water have a high boiling/ melting point for a simple molecule

A

hydrogen bonding between molecules

50
Q

why is ice less dense than water

A

ice has an open lattice with hydrogen bonds holding the water molecules apart
when the ice melts the rigid hydrogen bonds collapse allowing the water molecules to move closer together
the distance the molecules are held apart is what makes ice less dense than water

A Level Chemistry (6 decks)

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