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chemistry > bonding and structure > Flashcards

bonding and structure Flashcards

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1
Q

definition of ionic bonding?

A

the electrostatic attraction between oppositely charged ions

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2
Q

what structure do ionic compounds adopt when in a solid state?

A

a giant lattice structure

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3
Q

why do ions not conduct electricity when in solid state

A

as ions are not free to move

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4
Q

properties of ionic compounds

A
  • high m.p and b.p
  • soluble
  • electrical conductors (l/aq)
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5
Q

why do ionic compounds have a high m.p and b.p and what alters it

A

as ionic bonds are strong - the greater the charge of the ion, the stronger the ionic bond, and therefore more energy required to break it - hence different ionic structures will have different melting/boiling points

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6
Q

why are ionic compounds soluble

A

the stronger the ionic bond, the less soluble the substance becomes

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7
Q

when are ionic compounds electrical conductors

A

when they are liquid or aqueous as ions are free to move

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8
Q

covalent bonding def.

A

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

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9
Q

which electron doesn’t reach its full outer shell?

A

boron

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10
Q

dative covalent bonding def.

A

a covalent bond in which the shared pair of electrons have been supplied by one of the bonding atoms only

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11
Q

result of 4BP + 0LP

A

tetrahedral

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12
Q

bond angle of tetrahedral shape

A

109.5

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13
Q

result of 3BP + 1LP

A

pyramidal

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14
Q

bond angle of pyramidal shape

A

107

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15
Q

result of 2BP + 2LP

A

non-linear

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16
Q

bond angle of non-linear shape

A

104.5

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17
Q

result of 3BP + 0LP

A

trigonal planar

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18
Q

bond angle of trigonal planar

A

120

19
Q

result of 5BP + 0LP (pf5)

A

trigonal pyramidal

20
Q

result of 6BP + 0LP (sf6)

A

octahedral

21
Q

bond angle of octahedral shape

A

90

22
Q

result of 2BP + 0LP (CO2)

A

linear molecule

23
Q

bond angle of linear shape

A

180

24
Q

electronegativity def.

A

a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond

25
Q

most electronegative element

A

fluorine

26
Q

least electronegative element

A

francium

27
Q

electronegativity trend across a period and down a group

A

increases as you go across a period, decreases as you go down a group

28
Q

non-polar bond def.

A

when the electronegativity difference is 0 or close to 0 eg diatomic molecules / with no charge separation across bond in a molecule

29
Q

intermolecular forces def.

A

weak attractions between dipoles on different molecules

30
Q

list types of IFS

A
  • induced dipole-dipole interactions (London forces)
  • permanent dipole-dipole interactions
  • hydrogen bonding
31
Q

when are London forces present

A

always present in all molecules and atoms

32
Q

how does boiling point and going down the group affect London forces

A
  • as we go down the group
  • boiling point increases
  • so no. of electrons increase
  • so London forces between molecules are stronger
33
Q

define a simple covalent lattice in solid state

A

weak London forces are holding the covalent bonds between molecules

34
Q

what is a simple covalent molecule

A

London forces between molecules in gas/liquid state

35
Q

properties of simple covalent molecules

A
  • low m.p. and b.p.
  • low density
  • cannot conduct electricity as no free moving electrons
  • the polar molecules are soluble in polar solvents, vice versa for non-polar molecules
36
Q

hydrogen bonding def.

A

a strong dipole-dipole attraction between an electron-deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom containing N, O or F on a different molecule

37
Q

hydrogen bonding criteria

A
  • electron deficient hydrogen
  • lone pair of electrons on N, O, or F
38
Q

how many hydrogen bonds can one water molecule form?

A

4

39
Q

anomalous properties of water

A
  • higher m.p/b.p compared to other simple covalent molecules
  • ice is less dense than liquid water
  • cohesive eg transpiration in plants
  • high surface tension (many h2 bonds)
40
Q

why is ice less dense than water?

A

when all the h2 bonds form, they hold the water molecules far apart creating an open lattice structure. when they’re further apart this means they’re less compact, therefore less dense

41
Q

metallic bonding def.

A

a strong electrostatic attraction between positive ions in a fixed position and delocalised electrons

42
Q

what is a giant metallic lattice for

A

all metals will adopt a giant metallic lattice

43
Q

properties of a giant metallic lattice

A
  • high m.p/b.p
  • conduct electricity due to delocalised/free moving electrons
  • insoluble - due to strong metallic bonds
44
Q

what makes a stronger metallic bond

A

the higher the charge of the ion

chemistry (8 decks)

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