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as level chemistry > 1.5 solid structures > Flashcards

1.5 solid structures Flashcards

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1
Q

what is a crystal structure?

A

a description of how ions, molecules or atoms are distributed in a solid or crystalline structure

(e.g giant ionic is a type of crystal structure)

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2
Q

what are the 4 main types of crystal structure you need to know?

A
  1. giant ionic
  2. giant covalent (/macromolecular)
  3. simple covalent (/molecular)
  4. metallic
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3
Q

what are the two giant ionic compounds you need to know?

A
  1. sodium chloride
  2. caesium chloride
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4
Q

what is the crystal coordination number of sodium chloride? why?

A

6.6

cause each sodium ion is bonded to 6 chloride ions and each chloride ion is bonded to 6 sodium ions

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5
Q

in a sketch of sodium chloride, which ion is drawn bigger? (chloride ion or sodium ion?)

A

chloride ion is drawn bigger than the sodium ion

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6
Q

what is the crystal coordination number of caesium chloride? why?

A

8.8

cause each caesium ion is bonded to 8 chloride ions and each chloride ion is bonded to 8 caesium ions

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7
Q

in a sketch of caesium chloride, which ion is drawn bigger? (chloride ion or caesium ion?)

A

chloride ion is drawn bigger than caesium ion

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8
Q

why is the crystal coordination number for CsCl larger than for NaCl?

A

Cs+ is larger than sodium so can create more bonds

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9
Q

what are the physical properties of giant ionic structures?

A
  • high melting points
  • generally very soluble
  • generally brittle
  • only conduct electricity when molten or dissolved
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10
Q

why do giant ionic structures have high melting points?

A

due to strong electrostatic forces of attraction between oppositely charged ions

(also there are lots of these bonds due to each bind being part of a giant lattice structure)

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11
Q

why are giant ionic structures generally very soluble?

A

as water is a polar molecule and each ion forms bonds with surrounding water molecules

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12
Q

why are giant ionic structures generally brittle substances?

A

they will smash when a force is applied to them as they have static bonds between ions

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13
Q

why do giant ionic structures only conduct electricity when in a molten or a dissolved state?

A

as ions are onlu free to move when in these states and hence carry a charge

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14
Q

what are the 2 giant covalent/macromolecular structures you need to know

A
  1. diamond
  2. graphite
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15
Q

what does diamond consist of? (bond? atom?)

A

covalently bonded carbon atoms

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16
Q

how many strong covalent bonds does each carbon make? in what arrangement?

A

each carbon makes 4 strong covalent bonds to a carbon atom in a tetrahedral arrangement

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17
Q

why is diamond described as a ‘hard’ structure?

A

cause it is arranged in a tetrahedral arrangement

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18
Q

does diamond conduct electricity? why?

A

no - bc all electrons are used up in bonding

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19
Q

why does diamond have an extremely high melting point?

A

bc lots of energy is required to break all of the strong C-C bonds within the giant covalent structure

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20
Q

is diamond soluble or insoluble? why?

A

insoluble
due to the absence of any polarity within the structure

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21
Q

what does graphite consist of? (bond? atom?)

A

covalently bonded carbon atoms

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22
Q

how many strong covalent bonds does each carbon make in graphite? in what arrangement?

A

each carbon makes 3 strong covalent bonds to a carbon atom in a trigonal planar arrangement

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23
Q

does graphite conduct electricity? why?

A

yes - because each carbon only makes 3 strong covalent bonds so there is one free moving electron which is delocalised per carbon atom

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24
Q

does graphite have a high melting point? why?

A

yes - bc lots of energy is required to break all of the strong C-C bonds within the layers

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25
Q

why is graphite considered a ‘soft’ material?

A

weak Van der Waals forces exist between layers - these are easily broken

26
Q

is graphite soluble or insoluble? why?

A

insoluble
due to the absence of any polarity within the structure

27
Q

what are the 2 simple covalent/molecular structures you need to know?

A
  1. iodine
  2. ice
28
Q

in what state does iodine exist at at room temperature?

A

as a solid

29
Q

what is the structure of iodine?

A

a crystal of iodine would have alternating layers of iodine molecules

30
Q

what forces are found between molecules in iodine?

A

weak Van der Waals forces

31
Q

why is iodine described as ‘volatile’?

A

the weak Van der Waal’s forces between molecules are easily broken and it means iodine molecules can exist in the gas phase with little heating

(if heated, goes from solid straight to gas)

32
Q

what does volatile mean?

A

it can easily go into a gas state
- the bond between gases is weak

33
Q

does iodine conduct electricity? why?

A

no bc it doesn’t have delocalised electrons in any physical state

34
Q

why are simple covalent molecules ofyen described as ‘weak, crumbly or powdery’ substances?

A

due to the weak forces between I2 molecules
(VdW)

35
Q

does iodine have a high or low melting and boiling point?

A

low

36
Q

what is the shape of how water molecules are held in place in ice?

A

water molecules are held tetrahedrally in place

37
Q

how many hydrogen bonds does each water molecule in ice form?

A

each water molecules in ice forms 2 hydrogen bonds between its lone pairs on the oxygen and 2 hydrogen bonds which involve its hydrogen atome

38
Q

why is ice less dense than water?

A

the molecules are held further apart than in liquid water

ice has ‘gaps’ (filled with air/ trapped air)

39
Q

what are the physical properties of simple covalent/molecular structures?

A
  • generally low melting and boiling points
  • can be soluble, depending on the molecule
  • often described as ‘weak, crumbly or powdery’ substances
  • generally do not conduct electricity
40
Q

why do simple covalent structures generally have low melting and boiling points?

A

due to weaker IMF forces between molecules

41
Q

what type of simple covalent molecules are soluble?

A

if a molecule can form hydrogen bonds or has a permanent dipole, then it is likely to be soluble
non-polar molecules are insoluble

42
Q

why do simple covalent structures generally not conduct electricity?

A

due to the absence of any free moving electrons or ions

43
Q

what is a metallic bond?

A

the attraction between the negative electrons in the delocalised electron sea and their attraction to the positive metal ions in the lattice

44
Q

do metallic structures conduct electricity?

A

yes - due to free moving electrons which can carry charge

45
Q

do metallic structures have high or low melting points? why?

A

they have high melting point
due to the strong electrostatic forces of attraction between the electrons and the ions in the lattice

46
Q

how are metallic structures ductile and malleable?

A

as the metal ion layers are able to alide over each other without the structure snapping off due to the electron sea which acts like a ‘glue’ holding the structure together

47
Q

why do group 2 metals have a higher melting point?

A
  • the group 2 metal ions gave a higher charge density (2+)
  • this results in a stronger attraction between the delocialised electrons and the positive metal ion
  • this results in a higher melting point as more energy is requird to break the metallic bond
48
Q

what is said about a charge of 2+ compared to a charge of +?

A

2+ has a higher charge density

49
Q

why are group 2 metals better electrical conductors than group 1 metals?

A

as they donate 2 valence electrons per ion to the electron sea
- therefore there are more electrons in the electron sea, resultung in more electrons being ablr to carry a charge

50
Q

why are ionic solids made up of oppositely charged ions packed around each other? (lattice)

A

as this maximises electrostatic attraction brtween the oppositely charged ions and minimises repulsion between ions with the same charge, ensuring maximum bond energy

51
Q

what are the properties of ionic compounds determined by?

A

the strength of the electrostatic forces between the ions

52
Q

what does the melting and boiling points of metals depend on?

A

the number of delocalised electrons per atom

generally the more delocaised electrons per atom, the higher the melting and boiling points

53
Q

describe how the structures of sodium chloride and caesium chloride are similar

A

they are both cubic

54
Q

describe how the structures of sodium chloride and caesium chloride are different?

A
  • caesium ions are bigger than sodium ions
  • co-ordination number 6.6 for sodium chloride and 8.8 for caesium chloride
55
Q

name the force or bond that is being overcome when the solid changes into a gas : iodine

A

Van der Waals

56
Q

name the force or bond that is being overcome when the solid changes into a gas : diamond

A

covalent bond

57
Q

which out of caesium chloride and sodium chloride has the higher melting point?

A

sodium chloride has a higher melting point than caesium chloride

58
Q

which out of iodine and ice has the higher melting point?

A

iodine has a higher melting point than ice

59
Q

why does ice and water have different densities at 0.C?

A
  • in ice hydrogen bonds hold the moleculee together in an open (tetrahedral) structure
  • the molecules are further apart in ice than they are in water (so ice is less dense than water)
60
Q

on formjn ice at 0.C, 50cm^3 of water expands to occupy 54.5cm^3. calculate the density of ice at 0.C?

A

volume is 50cm^3 so mass is 50g
density = 50/54.5 = 0.917 g cm^-1

61
Q

explain why:
aluminium has a higher melting temperature than sodium.
you should refer to the nature of the bonding [3]

A
  • bonding is metallic
  • this is attraction between the sea/delocalised electrons and the positive ions
  • Al3+ has more electrons in the sea than Na+ / Al3+ has a higher charge density than Na+
  • more energy is needed to overcome forces in Al

as level chemistry (16 decks)

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