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OCR A Level Chemistry > Paper 3B Part 2 > Flashcards

Paper 3B Part 2 Flashcards

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1
Q

Molecular formula of sodium carbonate

A

Na2CO3

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2
Q

Molecular formula of sodium hydroxide

A

NaOH

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3
Q

Molecular formula of pottasium hydroxide

A

KOH

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4
Q

Molecular formula of phosphoric acid

A

H3PO4

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5
Q

Define hydrogen ion

A

An ion with no electrons and 1 proton

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6
Q

Define acid

A

In water an acid releases hydrogen ions into the solution

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7
Q

Define strong acid

A

Releases all of its hydrogen ions into the solution and completely disassociates, this is irreversible

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8
Q

Define weak acid

A

Only releases a small number of its hydrogen ions into the solution and partially disassociates, this is reversible

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9
Q

How is sulphuric acid both a strong and weak acid?

A

Sulphuric acid is both strong and weak, it can form 2H+ and SO42- or H+ and HSO4-

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10
Q

Define a base

A

A compound that neutralises an acid, by accepting a hydrogen ion to form a salt – for example ammonias, hydroxides, carbonates and metal oxides

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11
Q

Define an alkali

A

A type of base that dissolves in water to form hydroxide ions

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12
Q

Define a salt

A

The product of a reaction in which the H+ ions from the acid are replaced by the metal or ammonium ions

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13
Q

Describe and explain the conductivity of sodium Na, chlorine Cl2, and sodium chloride NaCl (6 marks)

A

.Sodium is a metal and conducts as both a solid and liquid.Free moving delocalised electrons to carry the charge.Chlorine is a small simple molecule that does not conduct.No free moving delocalised electrons to carry the charge.Ionic bond between sodium and chlorine to form sodium chloride, conducts as molten or aqueous but not when solid.Free moving ions when molten and aqueous so can conduct as they carry the charge

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14
Q

Why is the shape of PCl5 special?

A

Trigonal bi-pyramidBond angles of 120’ and 90’

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15
Q

Describe and explain the shape of a NH3 molecule

A

3 bonding pairs + 1 lone pair = 4 electron pairsSo repels to corners of tetrahedronDon’t see the lone pair so therefore the shape is pyramidal

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16
Q

Bond angle of NH3

A

107’

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17
Q

Why do lone pairs of electrons have greater repulsion

A

not attracted by nuclei unlike bonding pairs

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18
Q

What is the order of repulsion?

A

LP:LP > LP:BP > BP:BP

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19
Q

By how much does one LP in a tetrahedron reduce the bond angle by?

A

2.5’

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20
Q

Bond angle of H2O and why?

A

109.5 – 2.5 – 2.5 = 104.5’2 x LP’s so (– 2.5’) 2 times to the bond angle on a tetrahedron (109.5)

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21
Q

Why is CO2 linear?

A

C has 2 electron regions (2 BP’s) so it is linear

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22
Q

Define electronegativity

A

a measure of the tendency of an atom to attract a bonding pair of electrons within a covalent bond.

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23
Q

Where are the electrons positioned when both atoms in a bond are identical?

A

When the 2 atoms are identical, the electrons are shared equally and stay in the middle

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24
Q

What did Linus Pauling invent and when?

A

Linus Pauling – invented the Pauling scale in 1932 which measures electronegativity on an atom

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25
Q

Why does it get more reactive as you go down group 1?

A

More reactive as you go down the groupLess nuclear attraction to the outer electronEasier to loseThe nucleus is being shielded by the electrons in-betweenSo the electron feels less nuclear attraction

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26
Q

Across a period what happens to the electronegativity?

A

increases

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27
Q

Down a group what happens to the electronegativity?

A

decreases

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28
Q

Why does electonegativity increase across a period?

A

.Size roughly equal across period with same shielding by 2, 8.Nuclear charge increases across period therefore greater nuclear attraction so electronegativity increases

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29
Q

Why does electonegativity decrease down a group?

A

.Down group size of atom increases.Shielding of outer shell increases.This reduces the nuclear attraction and electronegativity

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30
Q

In a bond, what happens if one atom is more electronegative than the other? Example.

A

.The more electronegative atom will have a greater share of the electrons.Take H and Cl as an example.H = 2,1, Cl = 3.0.Cl more electronegative so the electrons are attracted more towards it

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31
Q

What does delta +/- mean?

A

means it has a slight positive or negative charge now

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32
Q

When H reacts with Cl what happens to its electron density?

A

.H has lost a small amount of electron density – becomes s+

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33
Q

What type of molecule is H-Cl?

A

Polar

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34
Q

Why is H2O a polar molecule?

A

the dipoles act in the same direction

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35
Q

Why is Carbon Dioxide a non-polar molecule?

A

Symmetrical dipoles will cancel each other out, so CO2 is non-polar

36
Q

What atom gets the delta- charge?

A

The more electronegative atom will take the s- charge

37
Q

Can you predict bonds from eletronegativity?

A

Yes

38
Q

If there is no difference between electronegativity there is what sort of bond?

A

non-polar covalent bond

39
Q

If there is a small difference (0 to 1.8) between electronegativity there is what sort of bond?

A

polar covalent bond

40
Q

If there is a large difference (>1.8) between electronegativity there is what sort of bond?

A

ionic

41
Q

metal + hydrochloric acid –>

A

metal chloride + hydrogen

42
Q

Metal + sulphuric acid –>

A

metal sulphate + hydrogen

43
Q

Metal + nitric acid –>

A

Metal nitrate + hydrogen

44
Q

Metal + phosphoric acid –>

A

Metal phosphate + hydrogen

45
Q

Metal oxide + acid –>

A

Salt + water

46
Q

Metal hydroxide + acid –>

A

salt + water

47
Q

Metal carbonate + acid –>

A

salt + carbon dioxide + water

48
Q

Ammonia + acid –>

A

ammonium salt

49
Q

Where to observe the meniscu

A

eye level

50
Q

What is weighing by difference?

A

With weighting boat – used weighing boat = mass

51
Q

What is an intramolecular force?

A

strong bonds inside molecules

52
Q

What is an intermolecular force?

A

act between different molecules, weak compared to covalent bonds

53
Q

What are intermolecular forces caused by?

A

Intermolecular forces are caused by weak attractive forces between very small dipoles in different molecules

54
Q

3 types of intermolecular forces and their relative strength if ionic and covalent are 1000

A

Hydrogen Bonds 50Dipole-dipole forces 10London forces 1

55
Q

What is a dipole dipole force?

A

permanentA weak attractive force between permanent dipoles in neighbouring polar molecules

56
Q

Explain HCl as an example of dipole dipole intermolecular force

A

Take HCL, H is s+ and Cl s-, these charges weakly attract opposite charges in neighbouring molecules

57
Q

In London forces as the number of electrons increase …

A

so does the strength of the London Forces due to larger instantaneous dipoles created.

58
Q

In terms of intermolecular forces what can and cant dissolve in water

A

If no permanent dipoles it can’t dissolve in water, London forces don’t count as they are temporary dipoles

59
Q

How do London forces work?

A

.Movement of electrons produces a changing dipole in a molecule.At any instant, an instantaneous dipole will exist, but its position is constantly shifting.The instantaneous dipole induces a dipole on a neighbouring molecule.This induced dipole induces further dipoles on neighbouring molecules, which then attract one another

60
Q

How to find uncertainty?

A

Half the smallest possible unit you can make

61
Q

Percentage uncertainty =

A

(uncertainty of equipment/size of measurement) x 100

62
Q

What anomalous properties does water have?

A

Elevated melting and boiling points Ice floats (solid is less dense than the liquid)relatively high surface tension and viscosity

63
Q

Why does water have an elevated melting and boiling point?

A

due to strong H bonds between molecules

64
Q

Why does ice float?

A

as liquid water freezes the molecules arrange themselves in a more open, regular lattice structure

65
Q

What is a hydrogen bond?

A

A strong dipole-dipole attraction between .An electron-deficient hydrogen atom on one molecule (O-HS+ or N-HS+)and .A lone-pair of electrons on a highly electronegative atom on a different molecule (H-OS- or H-NS-)

66
Q

Why does water have relatively high surface tension and viscosity?

A

The hydrogen bonds

67
Q

In what compounds are hydrogen bonds useful?

A

organic compounds containing O-H and N-H bonds (alcohols, carboxylic acid etc)

68
Q

What are hydrogen bonds responsible for the shape of?

A

They are responsible for shape of proteins and even DNA

69
Q

New concentration equation

A

New conc = (original volume / new volume) * original conc

70
Q

Define oxidation number

A

A measure of the number of electrons that an atom uses t bond with an atom of another element. Oxidation numbers are derived from a set of rules.Also known as an element’s oxidation state

71
Q

What is the oxidation number of a neutral element?

A

zero

72
Q

What is the oxidation number of a monatomic ion?

A

The same as the charge of the ion

73
Q

What is the sum of all oxidation numbers in a neutral compound?

A

zero

74
Q

What is the sum of all oxidation numbers in a polyatomic ion equal to?

A

the charge on the ion

75
Q

What is the oxidation number of groups 1-3 in compounds

A

In compounds the elements of:Group 1 have an oxidation number of +1Group 2 have an oxidation number of +2Group 3 have an oxidation number of +3

76
Q

What is the oxidation state of hydrogen in a compound? What if it is in a binary metal hydride? What is a binary metal hydride?

A

The oxidation state of hydrogen in a compound is usually +1. If the hydrogen ion is part of a binary metal hydride (compound of metal and some hydrogen), then the oxidation state is -1.

77
Q

What is fluorines oxidation number in compounds?

A

Fluorine has an oxidation number of -1 in compounds.

78
Q

What is the oxidation state of oxygen usually? What is its oxidation state in a peroxide? What is its oxidation state when bonded to fluorine?

A

The oxidation number of oxygen in a compound is usually -2. If, however, the oxygen is in a class of compounds called peroxides (like hydrogen peroxide), then the oxygen has n oxidation number of -1. If the oxygen is bonded to fluorine, the number is +2.

79
Q

What happens with transition metals oxidation numbers?

A

In transition metals, the oxidation number can vary.

80
Q

What are chlorine, bromine, and iodines oxidation numbers? What is the exception?

A

Chlorine, Bromine and iodine usually have numbers of -1, except when it is in a compound with oxygen.

81
Q

In compound names with roman numerals, what does the roman numeral refer to?

A

The oxidation number of the element that is joined in name to the oxygen, but not the oxygen (for example in element(1) element(2)ate (Roman Numeral) then the roman numeral refers to the element (2))

82
Q

In potassium manganate(VII), what does the VII refer to?

A

The oxidation number of manganese

83
Q

In equations, if an elements oxidation state increases what has it undergone?

A

Oxidation

84
Q

In equations, if an elements oxidation state decreases what has it undergone?

A

Reduction

85
Q

What is a disproportionation reaction?

A

when the same element is both oxidised and reduced at the same time in a reaction

86
Q

Rules for writing ionic equations

A

1) Re-write out the equation showing the separate ions present in each substance.2) Look for ions that change state and/or form non-ionic compounds and underline them3) Cross out spectator ions 4) Rewrite the equation leaving out the spectator ions. 5) Make sure your equation still balances

87
Q

How does developing chemical processes with high atom economy benefit sustainability?

A

Less environmental waste as a high percentage of atoms are being used in useful products

OCR A Level Chemistry (17 decks)

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