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Edexcel Chemistry Flashcards > topic 2: bonding and structure > Flashcards

topic 2: bonding and structure Flashcards

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1
Q

what is the shape and bond angle of a molecule with 2 bonding pairs and 0 lone pairs and some examples

A

shape: linear
bond angle: 180
examples: CO2, CS2, HCN, BeF2

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2
Q

what is the shape and bond angle of a molecule with 3 bonding pairs and 0 lone pairs and some examples

A

shape: trigonal planar
bond angle: 120
examples:BF3, AlCl3, SO3, NO3-, CO3 2-

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3
Q

what is the shape and bond angle of a molecule with 4 bonding pairs and 0 lone pairs and some examples

A

shape: tetrahedral
bond angle: 109.5
examples: SiCl4, SO42-, ClO4-, NH4+

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4
Q

what is the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pairs and some examples

A

shape: trigonal pyramidal
bond angle: 107
examples: NCl3 ,PF3,ClO3
,H3O+

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5
Q

what is the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs and some examples

A

shape: bent
bond angle: 104.5
examples: OCl2, H2S, OF2, SCl2

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6
Q

what is the shape and bond angle of a molecule with 5 bonding pairs and 0 lone pairs and some examples

A

shape: trigonal bipyramidal
bond angle: 120 and 90
example: PCl5

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7
Q

what is the shape and bond angle of a molecule with 6 bonding pairs and 0 lone pairs and some examples

A

shape: octahedral
bond angle: 90
example: SF6

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8
Q

what is ionic bonding

A

strong electrostatic force of attraction
between oppositely charged ions formed by electron transfer

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9
Q

why are positive ions smaller in comparison to their atoms

A

because it has one less shell of electrons and the ratio of protons to electrons has increased so there is greater net force on remaining electrons holding them more closely

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10
Q

why are anions larger than their atoms

A

the negative ion has more electrons than the corresponding atom but the same number of protons. so the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger

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11
Q

does the ionic radius increase or decrease from N3- to Al3+ and why

A

decrease because there are increasing numbers of protons from N to F and then Na to Al but the same number of electrons. the effective nuclear attraction per electron therefore increases and ions get smaller

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12
Q

does the ionic radius increase or decrease down the group and why

A

increases down group
because as one goes down the group the ions have more shells of electrons

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13
Q

what are the physical properties of ionic compounds

A

*high melting points ( there are strong attractive forces between the ions)
*non conductor of electricity when solid (ions are held together tightly and can not move)
*conductor of electricity when in solution or molten. ( ions are free to move)
*brittle / easy to cleave apart

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14
Q

what is covalent bond

A

strong electrostatic attraction between 2 nuclei and the shared/ bonding pair of electrons

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15
Q

what is the strength of covalent bond demonstrated by

A

high melting points

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16
Q

why do giant covalent structures have high melting points

A

because they contain many
strong covalent bonds in a macromolecular structure. it takes a lot of energy to break the many
strong bonds

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17
Q

what is the effect of multiple bonds on bond strength and length

A

nuclei joined by multiple (i.e. double and triple) bonds have a greater electron density between them. this causes an greater force of attraction between the nuclei and the electrons between them, resulting in a
shorter bond length and greater bond strength

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18
Q

what is a dative bond and examples

A

when the shared pair of electrons in the covalent bond come from only one of the bonding atoms. a dative covalent bond is also called co-ordinate bonding
eg. NH4+, H30+, NH3BF3

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19
Q

what is electronegativity

A

the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself is measured on the Pauling scale

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20
Q

why do electronegativity increases across a period

A

the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more

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21
Q

why do electronegativity decreases down the group

A

because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases

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22
Q

what does a small electronegative difference mean

A

purely covalent

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23
Q

what would the electronegative difference when ionic

A

very large electronegativity difference (> 1.7)

24
Q

describe the formation of a permanent dipole

A

a polar covalent bond forms when the elements in the bond have different electronegativities . (Of around 0.3 to 1.7)
When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends

25
Q

what would a symmetrical molecule do to the polar

A

a symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar
the individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule
there is no net dipole moment: the molecule is non-polar

26
Q

when do London forces occur

A

occur between all molecular substances and noble gases
they do not occur in ionic substances

27
Q

what are induced dipoles

A

temporary dipoles can cause dipoles to form in neighbouring molecules

28
Q

what is the main factor affecting size of London forces

A

the more electrons there are in the molecule the higher the chance that temporary dipoles will form
this makes the London forces stronger between the molecules and more energy is needed to break them so boiling points will be greater

29
Q

why do the boiling points increase down group 7

A

increasing number of electrons in the bigger molecules causing an increase in the size of the London
forces between the molecules

30
Q

why do the boiling points increase through the alkane homologous series

A

increasing number of electrons in the bigger molecules causing an increase in the size of the London forces between molecules

31
Q

describe permanent dipole - dipole forces

A
  • permanent dipole-dipole forces occurs between polar molecules
  • it is stronger than London forces and so the compounds have higher boiling points
  • polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)
  • polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms
32
Q

describe hydrogen bonding

A

occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons. e.g. a –O-H -N-H F- H bond. there is a large electronegativity difference between the H and the O,N,F

33
Q

what is metallic bonding

A

the electrostatic force of attraction between the positive metal ions and the delocalised electrons

34
Q

what are the 3 factors that affect metallic bonding

A

number of protons
number of delocalised electrons per atom
size of ion

35
Q

why do metals have a high melting point

A

because the strong electrostatic forces between positive ions
and sea of delocalised electrons require a lot of energy to break

36
Q

why can metals conduct electricity

A

because the delocalised electrons can move through the structure

37
Q

why are metals malleable

A

because the positive ions in the lattice are all identical. so the planes of ions can slide easily over one another. the attractive forces in the lattice are the same whichever ions are adjacent

38
Q

describe the structure of diamond

A

Tetrahedral arrangement of carbon atoms. 4 covalent
bonds per atom diamond cannot conduct electricity because all 4 electrons per carbon atoms are involved in covalent bonds. they are localised and cannot move

39
Q

describe the structure of graphite

A

planar arrangement of carbon
atoms in layers. 3 covalent bonds
per atom in each layer. 4th outer
electron per atom is delocalised.
delocalised electrons between
layers
graphite can conduct electricity well between layers because one electron per carbon is free and delocalised, so electrons can move easily along
layers
it does not conduct electricity from one layer to the next because the energy gap between layers is too
large for easy electron transfer

40
Q

describe the structure of graphene

A

graphene is a new substance that is a one layer of graphite .i.e. 3 covalent bonds per atom and
the 4th outer electron per atom is delocalised
these have very high tensile strength because of the strong structure of many strong covalent
bonds
graphene can conduct electricity well along the structure because one electron per carbon is free
and delocalised, so electrons can move easily along the structure

41
Q

describe the structure of carbon nanotubes

A

These have very high tensile strength because of the strong
structure of many strong covalent bonds.
Nanotubes can conduct electricity well along the tube
because one electron per carbon is free and delocalised,
so electrons can move easily along the tube
nanotubes have potentially many uses. one being the potential to use as vehicles to deliver drugs to cell
there are delocalised electrons in buckminsterfullerene

42
Q

what are the properties of ionic

A

high melting and boiling points -because of giant lattice of ions with
strong electrostatic forces between
oppositely charged ions
solubility in water - generally good
poor conductivity when solid: ions
can’t move/ fixed in lattice
good conductivity when molten: ions can move

43
Q

what are properties of molecular simple molecule

A

low melting and boiling point - because of weak intermolecular
forces between molecules (specify type e.g London forces/hydrogen
bond)
solubility in water - generally poor
poor conductivity when solid: no ions to conduct and electrons are
localised (fixed in place)
poor conductivity when molten: no ions

44
Q

what are the properties of a macromolecular molecule

A

high melting and boiling point - because of many strong covalent bonds in macromolecular structure. take a lot of energy to break the many strong bonds
insoluble in water
conductivity when solid - diamond and sand: poor, because electrons can’t move (localised) graphite: good as free delocalised electrons between layers
poor conductivity when molten

45
Q

what are the properties of a metallic molecule

A

high melting and boiling point - strong electrostatic forces between positive ions and sea of delocalised electrons
insoluble in water
good conductivity when solid: delocalised electrons can move through structure
good conductivity when molten

46
Q

what are the possible reasons why there are 2 widely different values for compressive strength of graphite (both are valid)

A

lower values relate to the weak London forces
high values relate to the strong covalent bonds

47
Q

what is a possible reason as to why the density of iron is much higher than the density of graphite

A

iron atoms have a greater mass than carbon atoms
iron atoms pack closer together than carbon atoms

48
Q

ammonia and boron trifluoride react to form a compound NH3BF3 which contains a dative
covalent bond. Each of the molecules, NH3 and BF3, has a different feature of its electronic
structure that allows this to happen. Use these two different features to explain how a dative
covalent bond is formed

A

donation of lone pair from nitrogen to boron which has only has 6 electrons in outer shell

49
Q

during this reaction, the bond angles about the nitrogen atom and the boron atom
change. state the new H—N—H and F—B—F bond angles

A

both 109.5 bc there is now 4 BP and 0LP

50
Q

what is electronegativity

A

the relative ability of an atom to attract the bonding electrons in a covalent bond

51
Q

explain why both water and carbon dioxide molecules have polar bonds but only
water is a polar molecule

A
  • O is more electonegative than H and C
  • which results in a polar bond with O dipole -ve so C and H dipole +ve
  • CO2 is a symmetrical molecule and so the dipole moments cancel out
  • the V shape of water molecule/ the lone pairs of electrons of O means that the dipole moments don’t cancel
52
Q

explain, in terms of the structure and bonding of each element, the difference between these
values - silicon and chlorine

A

Si is a giant covalent structures and contains many strong covalent bonds
Cl2 is a simple molecular, diatomic molecule and contains weak London forces
covalent bonds in Si are stronger than London forces in chlorine so covalent bonds take more energy to break than London forces

53
Q

explain why diamond has a much higher melting temperature than iodine

A

iodine is simple molecular
- molecules are held together by weak London forces
diamond is a giant covalent structure with each carbon covalently bonded to 4 other carbons
- carbon atoms in diamond are held together by strong covalent bonds
strong covalent bonds require more energy to break than intermolecular forces

54
Q

Aluminium fluoride and aluminium chloride are both crystalline solids at room temperature.
Aluminium fluoride sublimes to form a gas at 1291°C (1564 K), whilst aluminium chloride
sublimes at 178°C (451 K).
Use the Pauling electronegativity values in the Data Booklet to explain these differences in
sublimation temperature

A

al and cl electronegative difference 1.5
al and f electronegative difference 2.5
aluminium chloride mostly covalent
aluminium fluorine bonds are more polar
aluminium fluoride has strong electrostatic forces of attraction between the ions
aluminium chloride is molecular so has weak London forces
more energy needed to break the strong bonds to cause sublimation in aluminium fluoride

55
Q
A

Edexcel Chemistry Flashcards (16 decks)

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