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Chemistry > Topic 2: Bonding > Flashcards

Topic 2: Bonding Flashcards

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1
Q

Define Ionic Bonding

A

the strong electrostatic attraction between the oppositely charged ions

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2
Q

what does charge density depend on?

A

ionic charge and ionic radius

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3
Q

what is proof for the existence of ions

A

migration of ions and physical properties

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4
Q

migration of ions example

A

Copper (II) Chromate
copper 2+ ions migrate to negative electrodes (blue)
chromate ions migrate to the positive electrode (yellow)

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5
Q

define covalent bonding

A

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

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6
Q

what is dative covalent bonding

A

instead of one electron coming from each atom, the electrons come from a single atom

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7
Q

below which period in the periodic table can atoms expand their octet

A

below period 2

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8
Q

define bond energy

A

measure of how strong a bond is

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9
Q

what does a double bond consist of?

A

a sigma bond and pi bond between 2 atoms

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10
Q

angle in a tetrahedral molecule

A

109.5

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11
Q

angle in a pyramidal molecule

A

107

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12
Q

angle in a bent molecule

A

104.5

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13
Q

define electronegativity

A

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

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14
Q

what are London forces

A

when a temporary dipole is attracted to an induced dipole

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15
Q

define metallic bonding

A

strong electrostatic attraction between metal ions and the delocalised electrons

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16
Q

tetrachloromethane, CCl₄, contains polar bonds but is not a polar molecule, why?

A

the centres of charges are in the same place therefor it is non polar

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17
Q

explain why the melting temp of magnesium is much higher than that of sodium

A

magnesium ion has a larger charge density
magnesium ions are smaller than sodium ions
magnesium ions have more delocalised electrons
magnesium ions have a greater attraction for the delocalised sea of electrons

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18
Q

the electrical conductivity of pure silicon is very low. explain why this is in terms of the bonding.

A

silicon’s outer electrons are fixed in covalent bonds

therefore silicons electrons are not free to move

19
Q

explain why sulphur forms an expanded octet but oxygen does not

A

to expand the octet an electron as to be promoted to a higher energy level
for sulfur the next energy level is fairly close, for oxygen the energy jump is too large

20
Q

describe the shape of TeF₆

A

octahedral. 90 degree angle

21
Q

describe the shape of AsF₅

A

trigonal bipyramidal. 120 and 90 degrees

22
Q

describe the shape of PH₃

A

trigonal pyramidal. 107.5 degrees

23
Q

state and explain whether you would expect magnesium fluoride to have a higher or lower melting point than magnesium oxide

A

magnesium fluoride has lower
fluoride has a lower charge than oxide
less energy needed to overcome/ break bonds

24
Q

suggest why copper (II) chromate ions dissolve in water

A

ions are hydrated

25
Q

state and explain the trend in ionic radius from K+ to Ti4+

A

ionic radius decreases
ions are isoelectronic
nuclear charge increases
attraction between electrons and nucleus increases

26
Q

aluminium fluoride (AlF₃) melts at a much higher temp than either aluminium bromide or aluminium chloride, but none of the compounds conduct electricity when molten. compare the structure and bonding present in the three compounds. (4)

A

AlF₃ bonding is covalent
structure is giant
other bonding is covalent
structure is simple molecular

27
Q

explain why hydrogen halides are gases at room temp

A

bp lower than room temp
weak intermolecular forces
little energy needed to overcome the forces between the molecules

28
Q

explain why less energy is needed to break the covalent bond in hydrogen iodide than in hydrogen bromide (3)

A

iodine atoms are larger than bromine atoms
bond is longer in HI than HBr
attraction between nuclei and shared pair of electrons is weaker

29
Q

state and explain the trend in bond polarity in the hydrogen halides as atomic number of the halogen increases

A

bond polarity decreases with atomic number
halogen atom becomes less electronegative
smaller difference in electronegativity between H and halogen

30
Q

explain why the F-P-F bond angle in PF₃ (107) is different to the F-Si-F bond angle in SiF₄ (109.5)

A

electron pair repulsion pushing bonds closer together

lone pairs repelling more than bonding pairs (PF₃)

31
Q

Explain why CH₄ molecules are tetrahedral and state the H-C-H bond angle

A

electron pairs repel
shape provides minimum repulsion
angle is 109.5

32
Q

Explain why the H-O-H bond angle in water is smaller than the H-N-H bond angle in NH₃

A

water has one more lone pair of electrons than NH₃
lone pairs repel more than bonding pairs
bonding pairs are pushed closer

33
Q

explain why HF has a higher melting point than HCl

A

HF forms hydrogen bonds
HCl only forms London forces
hydrogen bonds are stronger

34
Q

explain why HF dissolves better in water than hexane

A

both involve breaking hydrogen bonds (between HF molecules)
hydrogen bonds are reformed when HF dissolves in water
only London forces formed when HF dissolves in hexane

35
Q

describe how HF molecules form London forces

A 
random movement of electrons 
temporary dipole forms 
induced dipole (in adjacent molecules)
36
Q

explain why CH₃CH₂CH₂CH₃ has a higher boiling point than (CH₃)₃CH

A

branched alkane/ (CH₃)₃CH has weaker London forces

molecules have lower SA/ cant be as close together

37
Q

why are alkanes insoluble in water

A

they cant form hydrogen bonds

38
Q

Explain why methanethiol (CH₃SH) cant form hydrogen bonds

A

sulphur is less electronegative than oxygen

S-H bond is not polar enough (to attract S: from another molecule)

39
Q

state and explain which of methanol and methane is more volatile

A

methane
cant form hydrogen bonds
hydrogen bonds are stronger (than London forces)

40
Q

Explain why methanol molecules are polar and describe a lab test to prove it

A

(C-O and) O-H bonds polar
not symmetrical
test: charged object near jet of methanol
result: jet deflected

41
Q

describe the hydration of the chloride ions formed when HCl dissolves in water

A

partial positive H atom of water (1) is attracted to (full) negative charge of chlorine ions

42
Q

suggest why HCl dissolves well in methylbenzene (C₇H₈)

A

similar intermolecular forces

London forces

43
Q

why is ice less dense than water?

A

more open/more space between molecules (making it less dense) due to (3 Dimensional) lattice/ ring structure in ice
hydrogen bonds longer than the covalent bonds

44
Q

Aluminium fluoride and aluminium chloride are both crystalline solids at room temperature. Aluminium fluoride sublimes to form a gas at 1291qC (1564 K), whilst aluminium chloride sublimes at 178qC (451 K).

Use the Pauling electronegativity values in the Data Booklet to explain these differences in sublimation temperature.

A 
aluminium and chlorine electronegativity difference 1.5 AND aluminium and fluorine electronegativity  difference 2.5 
aluminium chloride (mostly) covalent / (small) molecule 
aluminium fluoride (bonds) more polar  
aluminium chloride molecular so weak(er) intermolecular forces / London forces 
aluminium fluoride is a giant structure/ strong electrostatic forces of attraction between the ions    
more energy needed to break the stronger bonds to cause sublimation in aluminium fluoride

Chemistry (8 decks)

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