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AQA A Level Chemistry > bonding [P1] PAPER 1+2 > Flashcards

bonding [P1] PAPER 1+2 Flashcards

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1
Q

why do lone pairs of electrons influence the shape/bond angles in a molecule?

A

• electron pairs repel to be as far as possible
• lone pairs repel more than bond pairs

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2
Q

define the term ‘electronegativity’

A

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

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3
Q

how do permanent dipole-dipole forces between molecules arise?

A

• difference in electronegativity leads to bond polarity
• the dipoles don’t cancel each other out, so the molecule has an overall permanent dipole
• there is an attraction between the δ+ on one molecule and the δ- on another

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4
Q

how do van der Waals’ forces between molecules arise?

A

• electron movement in the first molecule (OR temporary dipole in the first molecule) induces a dipole in another molecule
• induced-temporary attraction OR δ+ attracts δ- in adjacent molecules

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5
Q

put the following intermolecular forces in order of strength, from most to least strongest: dipole-dipole, van der Waal’s, hydrogen bonding

A

• hydrogen bonding
• dipole-dipole forces
• van der Waal’s forces

⚠️ ONLY simple covalent molecules have IMF forces

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6
Q

describe the bonding in metals

A

giant metallic lattice held together by strong electrostatic forces of attraction between positive metal ions and delocalised electrons (metallic bonds)

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7
Q

why do metals have a high melting point?

A

• giant metallic lattice held together by strong electrostatic forces of attraction between positive metal ions and delocalised electrons
• lots of energy required to overcome these forces

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8
Q

why are metals ductile and malleable?

A

layers of ions are able to slide over each other (while still being held together by delocalised electrons)

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9
Q

draw a diagram to show the arrangement of particles in a crystal of a metal

A

• regular pattern of identical positive ions (draw minimum of 6)
• delocalised electrons (same number of them as the overall positive charge, e.g. if the ion is 2+ then two delocalised electrons per ion)

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10
Q

why are metals good conductors of heat?

A

the tightly packed ions efficiently pass vibrational energy on effectively through the lattice structure

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11
Q

why are metals good conductors of electricity?

A

free/mobile delocalised electrons are able to move through the lattice structure

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12
Q

what does the strength of a metallic bond depend on?

A

• charge of ion
• size of ion

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13
Q

explain why magnesium has a higher melting point than sodium

A

[Mg and Na are in different groups, so…]
• same shielding
• magnesium ions have a higher charge than sodium ions
• so stronger forces of attraction between delocalised electrons and positive Mg ion OR metallic bonding is stronger in Mg

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14
Q

explain why strontium has a higher melting point than barium

A

[they are in the same group, so…]

• both ions have the same charge
• strontium ions are smaller ions
• strontium ions have less shielding
• stronger electrostatic forces of attraction between the positive ions and delocalised electrons in a strontium ion, so stronger metallic bonding (so more energy needed to overcome these forces)

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15
Q

describe the bonding in an ionic compound

A

giant ionic lattice held together by many strong electrostatic forces of attraction between oppositely charged ions

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16
Q

why do ionic compounds have high melting points?

A

• giant ionic lattice held together by many strong electrostatic forces of attraction between oppositely charged ions
• lots of energy required to overcome these forces

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17
Q

in what state can ionic compounds conduct electricity and why?

A

• molten or dissolved
• ions are free to move and carry charge

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18
Q

deduce why the type of bonding in nitrogen oxide is covalent rather than ionic

A

small difference in electronegativity

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19
Q

suggest why sodium iodide has a lower melting point than sodium bromide

A

iodide is a larger ion (so weaker attraction to the sodium ion)

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20
Q

common polyatomic ions - ammonium, cyanide, hydroxide, sulfate, nitrate, carbonate, phosphate, chlorate

A
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21
Q

what is a covalent bond?

A

shared pairs of electrons

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22
Q

what is a dative covalent bond?

A

shared pair of electrons with both electrons supplied by one atom

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23
Q

name the type of bond formed when a molecule of BF4- reacts with an F- ion and explain how

A

• dative covalent
• lone pair of electrons on F- are donated to BF4-

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24
Q

how is a dative covalent bond represented on a diagram?

A

arrow from atom donating electron pair to other atom

e.g. lone pair of electrons is donated from N of NH3 to H+ to form NH4+ :

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25
Q

how do hydrogen bonds arise?

A

attraction between lone pairs of electrons on N/O/F (that is bonded to H) and H δ+ (bonded to N/O/F) of the adjacent molecule

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26
Q

what must be present for hydrogen bonding to occur?

A

hydrogen bonded with nitrogen, oxygen or fluorine in the molecule

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27
Q

hydrogen bonding diagram

A

• the ‘ : |||| H — X ’ must be in a straight line [X being N, O or F]
• all partial charges and lone pairs shown

e.g. two molecules of water interacting:

28
Q

how do you determine the shape of a molecule?

A

(1) group number of central atom + charge of molecule + number of bonds in the molecule
(2) ANS to (1) divided by 2 ( = number of electron pairs - this tells you the structure)
(3) ANS to (2) - number of bonds = number of lone pairs
(3) work out name of shape and bond angles (for every lone pair, -2.5°)

e.g. METHANE, tetrahedral, bond angles = 109.5° :
• central atom = C = group 4
• charge = 0
• 4 bonds in the molecule
• SO, tetrahedral, bond angles = 109.5°

29
Q

name and bond angles of a molecule with a total of 2 electron pairs - 2 being bond pairs and 0 being lone pairs?

A

• linear
• 180°

e.g.

30
Q

name and bond angles of a molecule with a total of 4 electron pairs - 2 being bond pairs and 2 being lone pairs?

A

• bent
• arrangement = tetrahedral as there’s a total of 4 pairs -> 109.5° - 2.5° - 2.5° [2 LP] = 104.5°

e.g.

31
Q

name and bond angles of a molecule with a total of 3 electron pairs - 3 being bond pairs of electrons and 0 being lone pairs?

A

• trigonal planar
• 120°

e.g.

32
Q

name and bond angles of a molecule with a total of 4 electron pairs - 3 being bond pairs of electrons and 1 being a lone pair?

A

• trigonal pyramidal
• (arrangement = tetrahedral as there’s a total of 4 pairs -> 109.5° - 2.5° [1 LP] =) 107°

33
Q

name and bond angles of a molecule with a total of 4 electron pairs - 4 being bond pairs of electrons and 0 being lone pairs?

A

• tetrahedral
• 109.5°

e.g.

34
Q

name and bond angles of a molecule with a total of 5 electron pairs - 5 being bond pairs of electrons and 0 being lone pairs?

A

• trigonal bipyramidal
• 90°
• 120°

e.g.

35
Q

name of a molecule with a total of 5 electron pairs - 4 being bond pairs of electrons and 1 being a lone pair?

A

• seesaw

e.g.

36
Q

name and bond angles of a molecule with a total of 5 electron pairs - 3 being bond pairs and 2 being lone pairs?

A

• T shaped
• 90°, 180°

e.g.

37
Q

name and bond angles of a molecule with a total of 5 electron pairs - 2 being bond pairs of electrons and 3 being a lone pair?

A

• linear
• 180°

e.g.

38
Q

name and bond angles of a molecule with a total of 6 electron pairs - 6 being bond pairs of electrons and 0 being lone pairs?

A

• octahedral
• 90°

e.g.

39
Q

name and bond angle of a molecule with a total of 6 electron pairs - 5 being bond pairs of electrons and 1 being a lone pair?

A

• square pyramidal
• <90°
e.g.

40
Q

name and bond angles of a molecule with a total of 6 electron pairs - 4 being bond pairs of electrons and 2 being lone pairs?

A

• square planar
• 90°

e.g.

41
Q

factors that affect electronegativity?

A

• nuclear charge
• atomic radius
• shielding

42
Q

trend in electronegativity down a group and why?

A

• decreases
• atomic radius and shielding increase (outweighs the increase in nuclear charge)
• weaker attraction between the nucleus and pair of electrons

43
Q

trend in electronegativity across a period and why?

A

• increases
• atomic radius decreases (so atom gets smaller)
• nuclear charge increases
• same shielding
• stronger attraction between the nucleus and pair of electrons

44
Q

when is a molecule polar?

A

• polar bonds present
• (the molecule isn’t symmetrical so) dipoles do not cancel each other out

45
Q

why is ice less dense than water?

A

• water has hydrogen bonding
• there are spaces in the structure OR water molecules are held further apart OR water molecules are more spread out

46
Q

why do macromolecules (such as diamond or silicone dioxide) have high melting points?

A

• giant covalent lattice held together by many strong covalent bonds
• lots of energy needed to break these bonds

47
Q

how may a polar bond be drawn?

A

• partial charges
• arrow thing

48
Q

when do polar bonds arise?

A

when the atoms have a significant difference in electronegativity

49
Q

why are macromolecules (such as diamond and silicon dioxide) rigid?

A

many strong covalent bonds present which form a strong giant covalent lattice

50
Q

why can graphite conduct electricity?

A

free/mobile delocalised electrons which are able to move through the structure

51
Q

why do simple molecular compounds have low melting and boiling points?

A

• covalently bonded molecules held together by weak van der Waals forces
• not much energy required to overcome the vdW forces

52
Q

why are simple covalent molecules poor conductors of electricity?

A

no charged particles present that are mobile and able to carry charge

53
Q

even though it is a simple covalent molecule, why does water have an unusually higher boiling point?

A

• hydrogen bonding is present
• strongest intermolecular force
• more energy required to overcome these forces

54
Q

why do differently shaped objects of titanium have similar strengths?

A

same metallic bonding OR same bond strength

55
Q

why does a larger hydrocarbon have a higher boiling point than a smaller hydrocarbon?

A

larger one has stronger van der Waals’ forces between molecules as the molecules have a larger surface area and are packed closer together

56
Q

Explain the polarity of a polar bond.

A

• ___ is more electronegative than ___
• So [more electronegative element] is δ- and [other one] is δ+

57
Q
A
58
Q

Explain the bond angles in a molecule where bond angles are equal.

(e.g. tetrahedral molecule - ALL bond angles are 109.5° - why?)

A

electron pairs repel equally

59
Q
A

• B
• Na+ and F- have the same electron configuration
• Sodium ion has more protons than a fluoride ion, so has a greater attraction to the other electrons (so is the smaller ion)

60
Q

How do lone pairs of a specific molecule influence the bond angle in it?

A

• Lone pairs repel more than bond pairs
• So, the bond angle will be lower

61
Q

Why may molecules of the same substance NOT have dipole-dipole forces between them, despite each individual molecule having polar bonds?

e.g. CCl4

A

Symmetrical molecule, so dipoles cancel out

62
Q

Types of bonding in ammonium chloride?

A

Covalent, dative covalent, and ionic

63
Q

Suggest why magnesium is a liquid over a much greater temperature range compared to bromine.

A

Forces of attraction between the positive magnesium ions and delocalised electrons are stronger in liquid state

64
Q

Explain how the value of the Cl-Al-Cl bond angle in AlCl3 changes, if at all, on formation of the compound H3NAlCl3.

• NH3 and AlCl3 form a dative covalent bond (electron pair donated from N of NH3 to Al) to produce H3NAlCl3
• Shapes of NH3 and AlCl3:

A

• Aluminium is now surrounded by 4 bonds, so H3NAlCl3 is tetrahedral
• So, Cl–Al–Cl bond angle decreases to 109.5° in H3NAlCl3

65
Q

Use your knowledge of structure and bonding to explain why sodium bromide has a melting point that is higher than that of sodium, and higher than that of sodium iodide.

A

• Sodium has a giant metallic lattice structure and has metallic bonding – there is attraction between the positive ion and delocalised electrons
• NaBr and NaI have giant ionic lattice structures, and have ionic bonding – there is attraction between oppositely charged ions
• The ionic bonds are stronger than the metallic bonds
• Since a Br- ion is smaller than a I- ion, there is stronger attraction between the oppositely charged ions in NaBr than in NaI

66
Q

By reference to the forces between molecules, explain why ammonia is very soluble in water.

A

hydrogen bonding between water and ammonia

AQA A Level Chemistry (36 decks)

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