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Chemistry A Level AQA > 3.1.3- Bonding (PAPER 1+2) > Flashcards

3.1.3- Bonding (PAPER 1+2) Flashcards

1
Q

Write the formulas for phosphoric acid, nitric acid, sulfuric acid and hydrogen peroxide.

A

H3PO4
HNO3
H2SO4
H202

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2
Q

Write the formulas of these ions: nitrate, nitrite, phosphate, sulfate, sulfite, silicate.

A

NO3 -
NO2 -
PO4 3-
SO4 2-
SO3 2-
SIO4 2-

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3
Q

Write the formulas of these ions: borate, ammonium, chlorate, arsenate, iodate, selenate.

A

BO3 3-
NH4+
ClO3-
AsO4 3-
IO3 -
SeO4 2-

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4
Q

Write the formulas of a carbonate ion and a hydrogencarbonate ion.

A

CO3 2-
HCO3 -

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5
Q

Write the formulas of a hydride ion, a hydrogen ion a hydroxide ion and ammonia.

A

H-
H+
OH-
NH3

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6
Q

What is ionic bonding?

A

The electrostatic attraction between oppositely charged ions in an ionic lattice.

Every positive ion is attracted to every negative ion and vice versa.

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7
Q

What does the formula of an ionic compound show?

A

The ratio of the positive and negative ions required to give overall 0 charge.

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8
Q

What is the swap and drop method when working out the formula of an ionic compound?

A

1) Write the 2 ions
2) Swap their charges
3) Drop their charges to the bottom: keep the number but lose the charge symbol.
4) Simplify to lowest whole number ratio if necessary.

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9
Q

Use the swap and drop method to work out the formula of calcium nitrate.

A

1) Ca 2+, NO3 -
2) Ca -, NO3 2+
3) Ca(NO3)2

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10
Q

Use the swap and drop method to work out the formula of calcium oxide.

A

1) Ca 2+, O 2-
2) Ca 2-, O 2+
3) Ca2O2
4) CaO

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11
Q

What are the 6 properties of ionic compounds?

A

All solids at room temperature due to fixed ionic lattice- do not conduct energy when solid.

Conduct electricity when aqueous or dissolved as the ions are free to move and conduct electricity.

High melting and boiling points as there is a large amount of energy required to reduce the electrostatic attraction between the positive and negative ions.

Low conductivity as a gas.

Brittle as when a force is applied to an ionic compound the ions move and the same charged ions are lined up, causing repulsion.

Usually soluble in water as water molecules have attraction towards both the positive and negative ions- they can provide enough energy to remove an ion from the lattice.

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12
Q

What is metallic bonding?

A

The strong electrostatic attraction between positive metal ions (cations) and delocalised electrons.

Metal has lots of positive metal ions packed closed together in a regular arrangement with their delocalised electrons making a giant metallic structure leading to a crystalline structure.

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13
Q

What are the 4 properties of metals?

A

High melting and boiling points as there is a strong attraction between the + metal ions and the delocalised electrons.

Conduct electricity as the outer electrons are free to move within the structure and carry current.

Strong and malleable- applying a large amount of force (eg-heat) causes the layers of + ions to slide over eachother, the delocalised electrons move with the layers so the strength of metallic bonding remains constant.

Insoluble in water but some metals can react with water.

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14
Q

What is the structure of the metal magnesium?

A

The 2 outer electrons get delocalised, forming an Mg2+ ion.

Each Mg ion is surrounded by other Mg nuclei. The delocalised electrons are shared between them, there is a strong attraction between the positive metal ions and the negative electons.

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15
Q

Which factors affect the melting point/ metallic bonding strength in a metal?

A

1) The positive charge of the ions- if the + charge is higher, there will be more electrons in the ‘cloud’ so there will be a greater attraction between positive ions and free electrons= stronger bond/ higher melting point.

2) Size of the metal ions- increase in protons (across a period) decreases the atomic radius as the stronger nuclear charge attracts the electrons more strongly, the electrons are closer to the nuclei= greater density of electron sea = stronger electrostatic attractions= greater current.

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16
Q

What is a covalent bond?

A

When 2 atoms share a pair of electrons.

Occurs mostly between non-metal atoms.

Strongest type of chemical bond.

No charge is produced so there is no conductivity.

Can also be represented by a line.

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17
Q

What is a coordinate/ dative covalent bond?

A

When 1 atom donates 2 electrons to an atom or ion to form a bond.

This can be represented by an arrow- the direction of the arrow is the direction of the electron transfer from one species to another.

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18
Q

Give an example of a dative covalent/coordinate bond between 2 particles.

A

A H+ ion is electron deficient as it has lost the 1 electron in its outer shell.

Ammonia (NH3) has 1 lone pair of electrons: 2 electrons that are not involved in the bonding.

There can be a mutual attraction between NH3 and H+ forming NH4+, both electrons have been transferred from NH3.

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19
Q

What is a covalent double bond or triple bond?

A

Double- when 2 atoms share 2 pairs of electrons.

Triple- when 2 atoms share 3 pairs of electrons.

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20
Q

What happens to the electron density in a covalent bond between two atoms of the same element?

A

When the atoms covalently bonded together are the same, the shared pair of electrons in the bond are shared equally.

This is because the atoms have the exact same electronegativity.

The covalent bond is non-polar.

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21
Q

What happens to the electron density in a covalent bond between two different atoms?

A

The electrons in the bond are not shared equally e.g. in hydrogen fluoride (HF)

F more electronegative than H so electrons in the covalent bond will be distorted towards the fluorine

F end of molecule has a delta negative charge, and the H end has a delta positive charge.

Bond is polar.

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22
Q

When is a covalent bond more polar?

A

If there is a greater difference in electronegativity between the 2 atoms.

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23
Q

What conditions lead to a polar molecule?

A

A molecule will be polar if:

it contains at least 1 polar bond

it has a net dipole (a permanent separation of charge, i.e. the electron density is not evenly distributed)

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24
Q

What are the properties of simple molecular substances?

A

Low melting and boiling points as in a solid state the simple molecules making up the compound are held together by weak VDW’s forces. Only the atoms within the molecule are held together by strong covalent bonds.

Do not conduct electricity as there are no free-moving ions or electrons.

Brittle as there are weak forces between the molecules which are easily broken.

Soluble in non-polar solvents as VDW’s forces form between solvent and molecule.

Do not dissolve in water.

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25
Q

What are examples of giant covalent structures?

A

Diamond, graphite, graphene, silicon oxide.

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26
Q

Properties of silicon dioxide?

A

Giant covalent structure, each silicon atom bonded to four oxygen atoms in a tetrahedral arrangement.

V high melting/boiling point- Very strong silicon-oxygen covalent bonds have to be broken throughout the structure before melting occurs.

Does not conduct electricity- there are no delocalised electrons.

Insoluble in water and organic solvents, the interaction between SiO2 and water molecules is insufficient to overcome the strong covalent bonds.

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27
Q

Properties of diamond?

A

Tetrahedral shape- each carbon is bonded 4 times to 4 other carbon atoms.

Doesn’t conduct electricity as there are no delocalised electrons within the structure that could carry charge.

The strong covalent bonding of the carbon atoms and the tight arrangement of atoms means that diamond is a good heat conductor.

Unlike graphite, diamond can be cut to make gemstones.

Very high melting point as there are very strong covalent bonds which need to be overcome. Very hard.

Insoluble- covalent bonds are too strong for water to break them.

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28
Q

Properties of graphite?

A

Made up of hexagons- each carbon is bonded 3 times: the 4th electron is delocalised.

Very high melting point as there are lots of strong covalent bonds.

Made up of layers with weak VDW’s forces between them, they can slide over each other easily.

Very soft- no covalent bonding between the layers of C atoms, they are held together by weak VDW’s forces.

Graphite can conduct electricity as there are delocalised electrons between the layers which can carry charge.

Graphite layers are far apart in comparison to a covalent bond length so graphite has quite a low density.

Insoluble: the covalent bonds are too strong to be broken by water.

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29
Q

Properties of graphene?

A

Very high melting+ boiling points as each C atom is covalently bonded to 3 other C atoms, covalent bonds are very strong.

Conducts electricity as each C atom has 1 unbonded electron- these are delocalised and can carry charge.

Very strong as there is a regular arrangement of C atoms joined by covalent bonds.

Insoluble in water as attractions between solvent molecules and C atoms not strong enough to break covalent bonds.

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30
Q

When are non-polar substances insoluble and soluble?

A

Non-polar substances are usually insoluble in polar solvents (WATER) as the water molecules form H bonds to each other and do not form VDW’s forces with the substance. NP substances are usually soluble in non-polar solvents .

Polar substances dissolve in polar solvents, and are usually insoluble in non-polar solvents.

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31
Q

How are 3D shapes of molecules represented?

A

Solid line= the bond lies on the plane of the page.

Solid wedge= the bond is coming out of the plane of the page.

Dotted wedge= the bond is projecting back behind the plane of the page.

32
Q

What is the electron pair repulsion theory?

A

The shape of a molecule is determined by the electron pairs surrounding the central atom, in the outer shell.

Pairs of electrons repel other pairs of electrons, these pairs move as far apart as possible to minimise this repulsion.

33
Q

What is the shape of a BeCl2 molecule?

A

There are 2 electron pairs around the central atom of Be, in order to move the furthest apart from each other they would have to move across a straight line.

BeCl2 has a linear structure: the angle between the 2 covalent bonds is 180°

34
Q

What is the shape of a CO2 molecule?

A

The central atom is carbon, this has 2 double bonds to 2 oxygen atoms.

Double bonds are treated in the same way as single covalent bonds so the 2 bonding areas repel and move as far apart as possible.

Linear shape: 180° bond angle.

35
Q

Requirements for a linear shape?

A

If a central atom forms 2 bonds/ has 2 bonding areas.

Central atom must have no lone pairs.

180° bond angle.

36
Q

What is the shape of a BF3 molecule?

A

Central B atom bonded to 3 F atoms, the electron pairs repel and move as far apart.

The bonds arrange themself at the points of a triangle= trigonal shape.

From the side, the molecule would be flat= planar.

This shape is called a trigonal planar.

37
Q

Requirements for a trigonal planar structure?

A

A central atom with 3 pairs of bonding electrons around it.

Central atom must have no lone pairs of electrons.

120° bond angle.

38
Q

What is the shape of a CH4 molecule?

A

Central C atom with 4 bonds forming with H

2 straight lines, 1 dotted line and 1 solid wedge.

All the angles between the bonds are 109.5°

This shape is called a tetrahedral molecule.

39
Q

Requirements for a tetrahedral structure?

A

4 electron pairs around central atom.

NO LONE PAIRS OF ELECTRONS: all involved in bonding.

109.5° between each bond.

40
Q

What is the shape of a PCl5 molecule?

A

Central P atom with 5 pairs of bonding electrons around it.

2 bonding pairs move to opposite sides of the molecule: using straight lines to represent the bonds.

Other 3 bonding pairs take up central position lying on the same plane: 1 straight line, 1 dotted line and 1 wedge.

Bonds pointing up and down are at 90° to the central atom.

Angle between bonds lying on the plane is 120°

This is called a trigonal bipyramidal shape.

41
Q

Requirements for a trigonal bipyramidal shape?

A

Five atoms surround the core atom in a trigonal bipyramidal arrangement.

3 atoms in a plane with bond angles of 120° with 2 on opposite ends of the molecule with bond angles of 90°

NO LONE PAIRS OF ELECTRONS

42
Q

What is the shape of a SF6 molecule?

A

6 bonding pairs around the central atom S.

1 bonding pair above S, 1 bonding pair below S both shown by straight lines both at 90° to the central plane.

4 bonding pairs on central plane, angle between these bonds is 90°- 2 shown by wedges, 2 shown by dotted lines.

Octahedral shape.

43
Q

Requirements for an octahedral shape?

A

6 electron pairs surrounding the central atom.

2 pairs above and below, 4 in central plane.

All at 90° bond angles.

NO LONE PAIRS OF ELECTRONS.

44
Q

Examples of ions with no lone pairs.

A

Carbonate ion = CO2 3-
1 double C=O bond and 2 single C-O bonds: 3 bonding areas around the central atom so the ion forms a trigonal planar structure with bond angles 120°

Nitrate ion= NO3 -
Central atom surrounded by 1 double bond and 2 single bonds, one of the single bonds is a dative bond but the ion still forms a trigonal planar structure.

Sulfate ion= SO4 2-
Central S atom surrounded by 2 single bonds and 2 double bonds= 4 bonding areas forming a tetrahedral structure with bond angles 109.5°

45
Q

Order of bond repulsions?

A

Lone pair–lone pair repulsion is greater than lone pair–bond pair repulsion, which is greater than bond pair–bond pair repulsion

46
Q

What difference does a lone pair make to the structure of a molecule?

A

Lone pairs repel more strongly than bond pairs, this extra repulsion decreases other bond angles by 2.5°

47
Q

Structure of ammonia which has a lone pair (NH3)?

A

N atom of ammonia has a lone pair, there are 3 bonding pairs.

4 pairs of electrons= tetrahedral structure.

The structure of NH3 is based on the tetrahedral structure. Bond angle is reduced: 109.5-2.5=107° bond angles.

This forms a pyramidal shape.

48
Q

Structure of an ammonium ion?

A

The lone pair of ammonia has formed a dative covalent bond between N and H, transferring both of its outer electrons to H.

As dative covalent bonds have the same level of repulsion as normal covalent bonds, the bond angles are regular tetrahedral bond angles of 109.5°

49
Q

Structure of a water molecule?

A

Oxygen atom has 2 single covalent bonds to H atoms, but it also has 2 lone pairs of electrons.

There are 4 overall bonding areas due to the 2 lone pairs of oxygen, so the shape is based off a tetrahedron.

2 lone pairs = 2 x 2.5 reduction in bond angle

=109.5- 2(2.5)= 104.5° = bent/ non-linear shape.

50
Q

What is the method of working out the shape of a molecule or ion?

A

1) Work out central atom and how many outer electrons it has (group number)
2) Add 1 electron for each bond formed.
3) If it is an ion, remove an electron for each + charge or add an electron for each - charge.
4) Work out the total number of electrons and the number of electron pairs.
5)Number of bond pairs= shape. Eg) 4 bond pairs is a tetrahedral shape.

51
Q

What is the shape with 3 bonding pairs and 1 lone pair?

A

Pyramidal

109.5-2.5= 107° bond angles

52
Q

What is the shape with 2 bonding pairs and 2 lone pairs?

A

Bent/non-linear/ V-shaped

Bond angles of 109.5-5= 104.5° bond angles

53
Q

What is the shape with 3 bonding pairs and 2 lone pairs?

A

T-shape

87.5 bond angle

54
Q

What is the shape with 4 bonding pairs and 2 lone pairs?

A

Square planar
Bond angles stay the same -90° (2 lone pairs repel equally from opposite sides.)

55
Q

What is electronegativity?

A

The power of an atom to attract the pair of electrons towards itself in a covalent bond.

56
Q

Where are most electronegative elements on periodic table? +most electronegative element.

A

The further right and up the periodic table.

F= most electronegative element.

NOT NOBLE GASES

57
Q

Which 3 factors affect electronegativity?

A

1) Nuclear charge: more protons = stronger attraction between the nucleus and the bonding pair of electrons.

2) Atomic radius: closer to nucleus= stronger attraction between nucleus and bonding pair of electrons.

3) Shielding: less shells of electrons between nucleus and outer electrons= less repulsion so there is a stronger attraction between the nucleus and bonding pair of electrons.

58
Q

What is the trend in electronegativity down a group?

A

Electronegativity decreases: the atomic radius increases as the number of electron shells (period number) increases so there is more shielding so there is less attraction between the nucleus and the bonding pair of electrons, this decreases the melting/boiling point.

If the atomic radius is lower (above), then the ions can pack closer together and exhibit a stronger electrostatic attraction between them.

59
Q

What is the trend in electronegativity across a period?

A

Electronegativity increases: the atomic radius decreases as the increased + charge shrinks the electron cloud, the higher nuclear charge means there is a stronger attraction between the nucleus and the bonding pair of electrons.

The number of inner principle energy levels stays the same so shielding remains mostly the same

60
Q

What does delta - and delta + show?

A

delta - is the more electronegative atom, it is electron rich.

delta + is the less electronegative atom, it is electron deficient.

61
Q

Which group is not included in electronegativity trends?

A

The noble gases.

62
Q

Define polar bonds- eg HCl.

A

Covalent bonds can become polar if the atoms attached to it have a difference in electronegativity.

The bigger the difference in electronegativity, the more polar the bond is.

HCl- Cl is more electronegative than H so it has a δ- symbol and attracts the electrons towards itself, H has a δ+ symbol as it is less electronegative.

63
Q

When are bonds not polar?

A

If atoms with the same or similar electronegativities are bonded together as the electrons sit basically in the middle of the 2 atoms.

Electrons are shared equally.

64
Q

What happens when the electronegativity difference between 2 atoms exceeds 1.7?

A

The covalent substance becomes ionic.

65
Q

When is a molecule not polar?

A

If the polar bonds are arranged symmetrically: eg CO2

66
Q

What is a permanent dipole?

A

If 2 atoms that are bonded have different electronegativities, a polar bond forms. The more electronegative atom draws more of t
he negative charge towards itself and away from other atom, producing a ∂- region and a ∂+ region.

67
Q

What are the 3 types of intermolecular forces?

A

Van der Waals, permanent dipole-dipole attraction and Hydrogen bonding.

68
Q

Give the intermolecular forces from weakest to strongest.

A

Van der Waals forces, permanent dipole-dipole, hydrogen bonding.

69
Q

What are Van der Waals/ induced dipole-dipole forces?

A

These exist within any molecules that contain electrons.

Electrons in charge clouds are always moving, at some point they may be distributed more on one side of the molecule.

A dipole can be created when any molecule or atom moves near another molecule or atom. Electrons in the molecule/ atom can move from one end to the other, creating a temporary dipole. The ∂+ on one molecule will be attracted to the ∂- of another. This force of attraction is broken when the atoms/ molecules move apart.

The larger the Mr of a molecule, the stronger the VDW forces.

70
Q

Why does the strength of VDW’s forces increase down Group 7?

A

Atoms within the diatomic molecules get larger and have more electrons, so there will be stronger VDW’s forces due to the more significant imbalance in charges so the boiling+ melting points increase.

F and Cl are gases as their VDW’s are so weak that room temp has enough energy to make the particles vaporise.

Bromine is a liquid at room temp: VDW’s forces are strong enough to keep the molecule in liquid form.

Iodine is solid at room temp: VDW’s are so strong that room temp does not have enough energy to separate the molecules.

71
Q

Why does the strength of VDW’s forces increase down Group 0?

A

The atoms get larger and they have more electrons so the VDW’s forces increase, the boiling points will increase.

All are gases as the VDW’s forces are weak for all.

72
Q

Why do straight chain molecules have stronger VDW’s forces than branched molecules?

A

They can line up and pack closer together, reducing the distance over which the force acts.

Eg- boiling point of straight hydrocarbons is larger than the boiling point of branched hydrocarbons.

73
Q

What are permanent dipole-dipole forces?

A

These are stronger than VDW’s forces.

These are between molecules with a polar bond. The ∂+ and ∂- regions attract each other and hold the molecules together in a lattice-like structure.

Molecules with permanent dipole-dipole forces also have VDW’s forces between the polar molecules.

74
Q

How can you test for polar molecules?

A

Place a charged rod near a steady stream of a polar liquid. The liquid will bend towards the rod as the molecules align to face the oppositely charged rod.

75
Q

What is hydrogen bonding?

A

Strongest intermolecular force.

Hydrogen bonding only exists between hydrogen and the 3 most electronegative elements: nitrogen oxygen or fluorine. The lone pair of electrons of these atoms forms a bond with the hydrogen atom of another molecule. N, O or F can draw the bonding electrons away from the H atom, causing the bond to be very polarised.

Hydrogen bonds are shown with dotted lines.

Molecules that have hydrogen bonding also have VDW’s forces and permanent dipole-dipole forces.

76
Q

Why is ice less dense than water?

A

As liquid water cools to form ice, the molecules make more hydrogen bonds and arrange themselves into a more open, regular 3-D lattice structure.

Since hydrogen bonds are relatively long, the average distance between water molecules is greater in ice than in liquid water - so ice is less dense than liquid water.

77
Q

When are simple molecules soluble?

A

If they can form hydrogen bonds to water as water is a polar molecule.

Polar substances will dissolve in polar solvents (water).

Chemistry A Level AQA (12 decks)

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