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chemistry > *Topic 2 - Bonding and structure > Flashcards

*Topic 2 - Bonding and structure Flashcards

1
Q

metal atoms lose electrons to form (1) ions

A

1 - positive

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2
Q

non metal atoms gain electrons to form (1) ions

A

2 - negative

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3
Q

State the electronic configuration of Magnesium as an element as an ion

A

Element - 1s2 2s2 2p6 3s2
Ion (Mg2+) - 1s2 2s2 2p6

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4
Q

What is the structure of ionic crystals

A

giant ionic lattices

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5
Q

Define ionic bonding

A

Ionic bonding is the strong electrostatic force of attraction between oppositely charged ions formed by electron transfer.

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6
Q

Explain what could cause ionic bonding to be stronger (in some compounds vs others)

A

Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges. E.g. MgO has a higher melting point than NaCl as the ions involved (Mg2+ & O2- are smaller and have higher charges than those in NaCl , Na+ & Cl- )

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7
Q

Comment on the ionic radius of positive ions compared to their atoms

A

Positive ions are smaller compared to their atoms because it has one less shell of electrons and the ratio of protons to electrons has increased so there is greater net force on remaining electrons holding them more closely.

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8
Q

Comment on the ionic radius of negative ions compared to their atoms

A

The negative ions formed from groups five to seven are larger than the corresponding atoms.
The negative ion has more electrons than the corresponding atom but the same number of protons. So the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger.

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9
Q

What happens to the ionic radii down a group

A

Within a group the size of the ionic radii increases going down the group. This is because as one goes down the group the ions have more shells of electrons.

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10
Q

List 4 physical properties of ionic compounds

A
  • high melting point (strong attractive forces between ions)
  • non conductor of electricity when solid (ions are held together tightly and can not move)
  • Conductor of electricity in solution or molten (ions are free to move)
  • Brittle / easy to cleave apart (a little force will push the ions along and ions will be next to similar ions; there will be a force of repulsion between like ions, pushing the layers apart)
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11
Q

Describe the migration of ions in CuCrO4 near electrodes

A

CuCrO4 is black, after a while:

  • blue colour of Cu2+ ions migrate to the negative electrode
  • yellow colour of CrO4^2- ions migrate to the positive electrode
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12
Q

What is a covalent bond

A

A covalent bond is strong and is caused by the electrostatic attraction between the bonding shared pair of electrons and the two nuclei.

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13
Q

What compounds can demonstrate the strength of the covalent bond and explain what property demonstrates this strength

A

The strength of covalent bond can be demonstrated by the high melting points of giant atomic structures like diamond and graphite. They have high melting points because they contain many strong covalent bonds in a macromolecular structure. It takes a lot of energy to break the many strong bonds.

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14
Q

In a covalent compound there is significant electron density (1) the atoms

A

1 - between

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15
Q

What is the effect of multiple bonds on bond strength and lengths

A

Nuclei joined by multiple (i.e. double and triple) bonds have a greater electron density between them.
This causes an greater force of attraction between the nuclei and the electrons between them, resulting in a shorter bond length and greater bond strength.

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16
Q

What is a dative bond and give an example

A

A dative covalent bond forms when the O
shared pair of electrons in the covalent bond come from only one of the bonding atoms. A dative covalent bond is also called co-ordinate bonding.

eg. NH4+, H3O+, NH3BF3

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17
Q

What is the direction of the arrow in a dative bond

A

The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient

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18
Q

Linear

Bonding pairs
Lone pairs
Diagram
Bond angle
Example

A

Bonding pairs - 2
Lone pairs - 0
Diagram - 3 elements in straight line
Bond angle - 180
Example - CO2, HCN, BeF2

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19
Q

Trigonal planar

Bonding pairs
Lone pairs
Diagram
Bond angle
Example

A

Bonding pairs - 3
Lone pairs - 0
Diagram - 3 elements around one element evenly spaced out, no lone pairs (peace sign)
Bond angle - 120
Example - BF3, AlCl3, SO3, NO3-, CO3^2-

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20
Q

Tetrahedral

Bonding pairs
Lone pairs
Diagram
Bond angle
Example

A

Bonding pairs - 4
Lone pairs - 0
Diagram - tetrahedral shape, two lines flat, one going out and one in
Bond angle - 109.5
Example - SiCl4, SO4^2-, ClO4^-, NH4^+

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21
Q

Trigonal pyramidal

Bonding pairs
Lone pairs
Diagram
Bond angle
Example

A

Bonding pairs - 3
Lone pairs - 1
Diagram - bottom three of the tetrahedral, one straight, one out of the page, one into the page and then a lone pair at the top
Bond angle - 107
Example - NCl3, PF3, ClO3, H3O^+

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22
Q

Bent

Bonding pairs
Lone pairs
Diagram
Bond angle
Example

A

Bonding pairs - 2
Lone pairs - 2
Diagram - two straight at the bottom and two lone pairs at the top like water
Bond angle - 104.5
Example - OCl2, H2S, OF2, SCl2

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23
Q

Trigonal bipyramidal

Bonding pairs
Lone pairs
Diagram
Bond angle
Example

A

Bonding pairs - 5
Lone pairs - 0
Diagram - one straight up and down and then 3 on the plane, one straight, one out and one in
Bond angle - 120 on the plane, 90 between the top and bottom and plane
Example - PCl5

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24
Q

Octahedral

Bonding pairs
Lone pairs
Diagram
Bond angle
Example

A

Bonding pairs - 6
Lone pairs - 0
Diagram - one straight top and bottom, 4 on the plane two in and 2 out
Bond angle - 90
Example - SF6

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25
Q

What are the marking steps to explaining shape of molecules

A

1 - state number of bonding pairs and lone pairs
2 - State that electrons repel to a point of minimum repulsion
3 - If there are no lone pairs, electrons repel equally
4 - If there are lone pairs, lone pairs repel more than bonding pairs
5 - state the actual shape and bond angle

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26
Q

Define electronegativity

A

Electronegativity is the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself.

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27
Q

What are the most electronegative atoms?

A

F, O, N, Cl

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28
Q

What is the single most electronegative element and what’s its value on the Pauling’s scale

A

Fluorine - has a value of 4.0

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29
Q

What factors affect electronegativity

A

Increases across a period as the number of protons increase and the atomic radius decreases because the electrons in the same shell are pulled in more

Decreases down a group because the distance between the nucleus and outer electrons increases and shielding of inner shell electrons increases

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30
Q

Which compounds will purely be covalent

A

compounds containing elements of similar electronegativity, small electronegativity difference

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31
Q

Which compounds will be ionic

A

Compounds containing elements of very different electronegativity, very large electronegativity difference

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32
Q

Explain the formation of a permanent dipole

A

A polar covalent bond forms when the elements in the bond have different electronegativities.

When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation. (dipole) δ+ δ- ends

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33
Q

state the dipoles in a molecules of HCl

A

H - δ+
Cl - δ-

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34
Q

Can symmetric molecules be polar?

A

A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar. (The dipoles will cancel - no net dipole movement: non polar molecule)

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35
Q

In what substances do London forces occur

A

London forces occur between all molecular substances and noble gases, no ionic substances.

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36
Q

Explain how London forces arise

A
  • electrons move constantly and randomly
  • electron density fluctuates, temporary dipoles form
  • these temporary dipoles cause dipoles to form in neighbouring molecules (induced dipoles)
  • induced dipole in always the opposite sign to the original one
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37
Q

What is the main factor affecting London Forces

A

Number of electrons
More electrons; increased chance of formation of temporary dipoles; makes the London forces stronger between the molecules; more energy is needed to break them; higher boiling point

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38
Q

How does the shape of a molecule affect London Forces

A

Long straight chain alkanes will have a larger surface area of contact between molecules for London forces to form than compared to spherical shaped branched alkanes and so have stronger London forces

39
Q

Permanent dipole dipole forces occur (1) to London Forces

A

1 - in addition to

40
Q

Describe the properties of permanent dipole-dipole forces

A
  • occur between polar molecules
  • stronger than London forces; higher boiling points
  • Polar molecules have a permanent dipole eg. C-Cl, C-F, C-Br, H-Cl
  • Polar molecules are asymmetrical and have a bond where there is significant difference in electronegativity between atoms
41
Q

Between which elements can hydrogen bonding arise

A

Hydrogen -> F, O or N (must have an available lone pair)
Because there is large electronegativity difference

42
Q

What two things must you label when drawing hydrogen bonding

A
  • lone pair of electrons of F, O or N
  • dipoles
43
Q

Hydrogen bonding occurs (1) London forces

A

1 - in addition to

44
Q

What is the bond angle of the hydrogen in hydrogen bonding

A

180

45
Q

List some functional groups which can form hydrogen bonds

A

alcohols, carboxylic acids, proteins, amides etc…

46
Q

Alcohols can form hydrogen bonding? How does this mean their properties compare to that of alkanes

A

alcohols have higher boiling points and relatively low volatility compared to alkanes with a similar number of electrons.

47
Q

How many hydrogen bonds can each water molecule form

A

2 -> there are 2 oxygens with lone pairs; stronger hydrogen bonding; relatively high boiling point

48
Q

Why does ice have a lower density

A

the molecules are held further apart by the hydrogen bonds than in liquid water and this explains the lower density of ice

49
Q

state the order of strength of the 3 intermolecular forces

A
  1. hydrogen
  2. permanent dipole-dipole forces
  3. London dispersion forces
50
Q

What causes the general increase in boiling points from H2S to H2Te

A

Increase in number of electrons, stronger London forces

51
Q

What is solubility of a solute in a solvent

A

a complicated balance of energy required to break the bonds in the solute and solvent against energy given out making new bonds between the solute and solvent

52
Q

Describe the process by which an ionic lattice dissolves in water

A
  • breaking the bonds in the lattice
  • forming new bonds between the metal ions and water molecules
  • negative ions are attracted to δ+ hydrogens on the polar water molecule
  • position ions are attracted to the δ- oxygen on the polar water molecules
53
Q

The higher the charge density the greater the (1) (e.g. smaller ions or ions with larger charges) as the ions attract the water molecules more strongly.

A

1 - hydration enthalpy

54
Q

how does the solubility of alcohols vary with size

A

The smaller alcohols are soluble in water because they can form hydrogen bonds with water. The longer the hydrocarbon chain the less soluble the alcohol.

55
Q

Which molecules will be insoluble in water

A

Compounds that cannot form hydrogen bonds with water molecules, e.g. polar molecules such as halogenoalkanes or non polar substances like hexane will be insoluble in water.

56
Q

Compounds which have similar (1) to those in the solvent will generally dissolve

A

1 - intermolecular forces

57
Q

Give an example in which a non polar solute will dissolve in a non polar solvent

A

iodine which has only London forces between its molecules will dissolve in a non polar solvent such as hexane which also only has London forces.

58
Q

Describe how Propanone can be considered to have both polar and non polar characteristics

A

It can form London forces with some non polar substances such as octane with its CH3 groups. Its polar C=O bond can also hydrogen bond with water.

59
Q

What do metals consist of

A

Metals consist of giant lattices of metal ions in a sea of delocalised electrons

60
Q

What is the definition of metallic bonding

A

Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons

61
Q

What are the 3 main factors that affect the strength of metallic bonding

A
  • number of protons / strength of nuclear attraction; the more protons the stronger the bond
  • number of delocalised electrons per atom; the more delocalised electrons the stronger the bond
  • size of an ion; the smaller the ions, the stronger the bond
62
Q

Why do metals have high melting points

A

the strong electrostatic forces between the positive ions and sea of delocalised electrons require a lot of energy to break

63
Q

Explain why Mg has stronger metallic bonding than Na and hence. a higher melting point

A

The metallic bonding gets stronger because in Mg there are more electrons in the outer shell that are released to the sea of electrons. The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is needed to break bonds.

64
Q

Why can metals conduct electricity well

A

Metals can conduct electricity well because the delocalised electrons can move through the structure.

65
Q

Why are metals malleable

A

Metals are malleable because the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another. The attractive forces in the lattice are the same whichever ions are adjacent.

66
Q

Giant lattices are present in:

A

i ionic solids (giant ionic lattices)
ii covalently bonded solids, such as diamond, graphite and silicon(IV) oxide (giant covalent lattices) iii solid metals (giant metallic lattices)

67
Q

Describe the structure of a giant ionic lattice of sodium chloride

A

giant ionic lattice showing alternate Na+ and Cl- ions

68
Q

Describe the structure of a giant covalent structure of diamond

A

Tetrahedral arrangement of carbon atoms. 4 covalent bonds per atom

69
Q

Describe the structure of a giant covalent structure of grpahite

A

Planar arrangement of carbon atoms in layers. 3 covalent bonds per atom in each layer. 4th outer electron per atom is delocalised. Delocalised electrons between layers.

70
Q

Why do graphite and diamond have very high melting points

A

Both these macromolecular structures have very high melting points because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds

71
Q

How do you draw a metallic structure diagram

A

Giant metallic lattice 2D showing close packing metal IONS

72
Q

How do you draw a molecular diagram of eg. Iodine

A

regular arrangement if I2 molecules had together by weak London forces

73
Q

How do you draw a molecular diagram of ice

A

a central water molecule with two ordinary covalent bonds and two hydrogen bonds in a tetrahedral arrangement
The molecules are held further apart than in liquid water and this explains the lower density of ice

74
Q

Why can diamond not conduct electricity

A

Because all 4 electrons per carbon atoms are involved in covalent bonds. They are localised and cannot move.

75
Q

Why can graphite conduct electricity well between layers but not from now layer to the next

A

because one electron per carbon is free and delocalised, so electrons can move easily along layers.
It does not conduct electricity from one layer to the next because the energy gap between layers is too large for easy electron transfer.

76
Q

Why do the two macromolecular structures of diamond and graphite have very high melting points

A

because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds

77
Q

Describe the structure of graphene

A

Graphene is a new substance that is a one layer of graphite .i.e. 3 covalent bonds per atom and the 4th outer electron per atom is delocalised.

78
Q

Why dies graphene have very high tensile strength

A

These have very high tensile strength because of the strong structure of many strong covalent bonds

79
Q

Why can graphene conduct electricity along the structure

A

because one electron per carbon is free and delocalised, so electrons can move easily along the structure.

80
Q

why do carbon nanotubes have very high tensile strength

A

because of the strong structure of many strong covalent bonds.

81
Q

why can carbon nanotubes conduct electricity

A

Nanotubes can conduct electricity well along the tube because one electron per carbon is free and delocalised, so electrons can move easily along the tube.

82
Q

What is a potential use of carbon nanotubes

A

Nanotubes have potentially many uses. One being the potential to use as vehicles to deliver drugs to cells.

83
Q

Do buckminsterfullerene contain delocalised electrons

A

yes

84
Q

Give the definition, structure and an example of ionic bonding

A

Definition - electrostatic force of attraction between oppositely charged ions
Structure - giant ionic lattice
Example - Sodium chloride, Magnesium oxide

85
Q

Give the definition, structure and an example of covalent bonding

A

Definition - shared pair of electrons
Structure - simple molecular; with intermolecular forces (London forces, permanent dipoles, hydrogen bonds) between molecules
Example - Iodine, ice, carbon dioxide, water, methane

85
Q

Give the definition, structure and an example of covalent bonding (simple molecular)

A

Definition - shared pair of electrons
Structure - simple molecular; with intermolecular forces (London forces, permanent dipoles, hydrogen bonds) between molecules
Example - Iodine, ice, carbon dioxide, water, methane

ONLY TALK ABOUT MOLECULES AND INTERMOLECULAR FORCES WITH SIMPLE MOLECULAR SUBSTANCES

86
Q

Give the definitions, structure and an example of covalent bonding (macromolecular)

A

Definition - shared pair of electrons
Structure - macromolecular; giant molecular structures
Example - Diamond, Graphite, Silicon dioxide, Silicon

87
Q

Give the definitions, structure and an example of metallic bonding

A

Definition - electrostatic force of attraction between the metal positive ions and the delocalised electrons
Structure - Giant metallic lattice
Example - Magnesium, sodium (all metals)

88
Q

Comment on the boiling and melting points of ionic, simple molecular, macromolecular and metallic compounds

A

Ionic: high - because of giant lattice of ions with strong electrostatic forces between oppositely charged ions.

Simple molecular: low - because of weak intermolecular forces between molecules (specify type e.g London forces/hydrogen bond)

Macromolecular: high - because of many strong covalent bonds in macromolecular structure. Take a lot of energy to break the many strong bonds

Metallic: high - strong electrostatic forces between positive ions and sea of delocalised electrons

89
Q

Comment on the solubility in water of ionic, simple molecular, macromolecular and metallic compounds

A

Ionic: generally good

Simple molecular: generally poor

Macromolecular: insoluble

Metallic: insoluble

90
Q

Comment on the conductivity when solid of ionic, simple molecular, macromolecular and metallic compounds

A

Ionic: poor - ions can’t move/fixed in lattice

Simple molecular: poor - no ions to conduct and electrons are localised (fixed in place)

Macromolecular: diamond and sand - poor - electrons can’t move (localised). Graphite - good - free delocalised electrons between layers

Metallic: good - delocalised electrons can move through structure

91
Q

Comment on the conductivity when molten of ionic, simple molecular, macromolecular and metallic compounds

A

Ionic: good - ions can move

Simple molecular: poor - no ions

Macromolecular: poor

Metallic: good

92
Q

Comment on general description of ionic, simple molecular, macromolecular and metallic compounds

A

Ionic: crystalline solids

Simple molecular: mostly gases and liquids

Macromolecular: solids

Metallic: shiny metal
Malleable as the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another
-attractive forces in the lattice are the same whichever ions are adjacent

chemistry (9 decks)

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