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A-level chemistry > Bonding and Structure > Flashcards

Bonding and Structure Flashcards

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1
Q

what type of ions will metals always form

A

cations

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2
Q

define ‘metallic bond’

A

the strong electrostatic attraction between metal ions and a sea of delocalised electrons

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3
Q

define ‘metallic radius’

A

half the distance between the centers of two adjacent metal ions in a metallic lattice

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4
Q

in terms of melting temperature, how does this change with metallic radii

A

the smaller the metallic radius, the greater the melting point

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5
Q

name 6 chemical properties of metals

A
  • smaller first ionisation energy than non-metals in the same period
  • have the lowest electronegativity levels in the periodic table
  • low electronegativity and first IE means its more energetically feasible to lose electrons in bonding
  • react with dilute acids to form hydrogen gas and a solution of a salt
  • react with water to form an alkaline solution of a metal hydroxide and hydrogen gas
  • a more reactive metal will reduce the ions of a less reactive metal
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6
Q

define an ionic bond

A

the strong electrostatic attraction between oppositely charged ions

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7
Q

what is the octet rule

A

when ions are formed they normally have a full outer shell (like that of a noble gas)

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8
Q

properties of ionic bonds (4)

A
  • high melting and boiling points
  • brittle and easily cleaved (split)
  • soluble in water
  • electrical and thermal conductors when molten or in solution
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9
Q

what happens when ionic bonds are hydrated

A

hydrogen is attracted to the negative ions and oxygen is attracted to the positive ions in the bond

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10
Q

why do ionic bonds have a high melting and boiling point

A

because of the strong electrostatic attraction between the oppositely charged ions, it requires lots of energy to overcome

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11
Q

what is the melting point determined by in ionic bonding

A

the strength of the ionic bonds

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12
Q

name three factors effecting the strength of ionic bonds

A
  • the size of the ionic radius (smaller size means a greater strength)
  • charge of the ions (higher charge, greater strength)
  • the surface area
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13
Q

why does ionic radius increase down a group

A

the number of shells the ion contains increases

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14
Q

what is it called when ions have identical electron configurations

A

isoelectronic

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15
Q

why does ionic radius decrease from N3- to Al3+

A

the number of protons increases this increases the nuclear charge. so, the electrons face a stronger force of attraction to the nucleus, decreasing the ionic radius

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16
Q

define a covalent bond

A

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

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17
Q

what is a dative covalent/coordinate bond

A

a bond formed where one atom donates two electrons to an atom or ion

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18
Q

give two examples of molecules with a dative bond

A

ammonium, carbon monoxide

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19
Q

what are dative covalent bonds represented with

A

an arrow

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20
Q

give an example of a molecule with a single covalent bond

A

H-F

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21
Q

give an example of a molecule with a double covalent bond

A

O=O

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22
Q

give an example of a molecule with a triple covalent bond

A

N=-N

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23
Q

define bond enthalpy

A

the energy required to break a particular covalent bond of one mole of a molecule in gaseous state (strength of a covalent bond)

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24
Q

shorter covalent bond. higher or lower bond enthalpy?

A

higher

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25
Q

what two forces determine bond length of a covalent bond

A

attraction (positive nuceli-negative electrons)
repulsion (two positive nuclei + negative electrons)

26
Q

how is a high bond enthalpy created

A

greater electron density between atoms results in stronger forces of attraction. atoms are pulled closer to each other, shorter bond=high bond enthalpy

27
Q

which has a higher bond enthalpy
H-F or N=-N

A

N=-N (greater electron density, shorter bonds)

28
Q

Name 4 giant covalent structures

A
  • graphite
  • diamond
  • graphene
  • silicone dioxide
29
Q

Properties of graphite

A
  • insoluble
  • good electrical conductor
  • low density
  • very high melting point
  • lubricant
30
Q

Why is graphite insoluble

A

Due to the strong covalent bonds between the carbon atoms

31
Q

Why does graphite have a low density

A

Due to its layers being far apart

32
Q

Why can graphite conduct electricity

A

Each carbon atom is bonded to three other carbon atoms. That leaves a delocalised electron as the fourth. This can carry a charge throughout the layers

33
Q

Why does graphite have a very high melting point

A

Due to the strong covalent bonds between the carbon atoms

34
Q

Properties of diamond (4)

A
  • good thermal conductor
  • poor electrical conductor
  • insoluble
  • high melting point
35
Q

Why is diamond a good thermal conductor

A

It’s tightly packed, rigid arrangement allows for thermal energy to pass easily

36
Q

Why is diamond a poor electrical conductor

A

Each carbon atom is covalently bonded to four others In a tetrahedral shape. This means there are no delocalised electrons so no charge can be carried.

37
Q

Why is diamond insoluble

A

Due to the strong covalent bonds between carbon atoms, they are too strong to break

38
Q

Why does diamond have a very high melting point

A

Contains strong covalent bonds between the carbon atoms

39
Q

What giant covalent structure has a similar shape and properties to diamond

A

Silicone IV dioxide

40
Q

What are the properties of graphene (3)

A
  • lightweight and transparent
  • good electrical conductor
  • high strength properties
41
Q

Why is graphene lightweight and transparent

A

It’s one layer of graphite made up of hexagonal rings but only one atom thick

42
Q

Why is graphene a good electrical conductor

A

Contains delocalised electrons (carbon covalently bonded to only three other carbons)

43
Q

Why does graphene have high strength properties

A

Delocalised electrons strengthen the covalent bonds between carbon atoms

44
Q

What is graphene used in (3)

A
  • aircraft shells
  • super computers/high speed computers
  • smart phone screens
45
Q

How do covalent bonds become polar

A

If the atoms attached to it have a difference in electronegativity

46
Q

In terms of polarity, a bigger difference in electronegativity=

A

A more polar bond

47
Q

Give two examples of polar molecules

A

Hydrogen chloride
Water

48
Q

Why is H-Cl polar

A

The chlorine atom is more electronegative, electrons shift closer to the chlorine
H—;Cl

49
Q

Is Cl-Cl polar and why

A

No because the chlorine atoms are bonded with the same electronegativity values so the shared electrons sit in the middle
Cl-;-Cl

50
Q

Why is water polar

A

Because it is unsymmetrical

51
Q

Is CO2 polar and why

A

No because it is a symmetrical molecule
O;=C=;O

52
Q

What makes covalent molecules non polar

A

If the molecule is bonded with the same/similar electronegativity values, the shared electrons will sit in the middle of the bond. Furthermore if the molecule is symmetrical.

53
Q

how strong is the repulsion between two bonding pair electrons

A

low

54
Q

how strong is the repulsion between a lone pair and a bonding pair electron

A

medium

55
Q

how strong is the repulsion between two lone pair electrons

A

high

56
Q

how do lone pair electrons affect the shape of the molecule

A

they push bonding pairs closer together, for every lone pair the bonding angle reduces by 2.5degrees

57
Q

what is electronegativity

A

the ability of an atom to attract an electron pair

58
Q

why do molecules with more electrons have larger London dispersion forces

A
  • they will have larger fluctuations in electron density
  • leads to larger temporary dipoles, and stronger dipole-dipole interactions
59
Q

what type of molecules do you find permanent dipole-dipole interactions between

A

polar molecules

60
Q

where do you find hydrogen bonds

A

when hydrogen atoms form a strong permanent dipole-dipole interaction between either nitrogen, oxygen or fluorine

61
Q

name three examples of liquids with hydrogen bonds

A

water
hydrogen fluoride
ammonia

62
Q

why do hydrogen bonds form

A

when hydrogen is bonded to an extremely electronegative element is forms a strong delta plus charge. due to its small size, hydrogen has a high charge density so it can form a strong bond with any highly delta negative atom

A-level chemistry (21 decks)

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