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A-level chemistry > Module 2 > Flashcards

Module 2 Flashcards

1
Q

Relative isotopic mass

A

Mass of an isotope relative to 1/12 mass of an atom of carbon 12

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2
Q

Relative atomic mass

A

Weighted mean mass of an atom relative to 1/12 of an atom of carbon-12

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3
Q

Relative molecular mass

A

simple molecules

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4
Q

Relative formula mass

A

Giant ionic compounds

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5
Q

Isotopes

A

Atom of same element with different number of neutrons and different mass.

Isoelectronic

Different physical properties

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6
Q

Avogadros constant

A

6.02 x 10 ^23 mol-1

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7
Q

Isoelectronic

A

Same electron structure

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8
Q

Molar mass

A

Mass in g of one mole of a substance

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9
Q

mass=

A

Mr moles

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10
Q

Number of particles =

A

moles x avagadros constant

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11
Q

Empirical formula

A

Simplest whole number ration of atoms of each element present in a compound

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12
Q

Water of crystallisation

A

Water that is chemically bonded into a crystalline structure

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13
Q

Calculating the formula of a hydrated salt

A

1) moles of anhydrous and H2O
2) calculate ratio of amounts and formulas

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13
Q

Check water of crystallisation has been removed by …..

A

Heating to a constant mass

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14
Q

Concentration

A

The amount (in mol) of a dissolved substance in 1 dm3 of solution

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14
Q

Moles =

A

Concentration x volume

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15
Q

Molar gas volume increases as

A

Temperature increases

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16
Q

Molar gas volume decreases as

A

Pressure increases

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17
Q

Ideal gas equation

A

pV= nRT

T in Kelvin (+273)
V (m3)
R = 8.314 Jmol-1K-1

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17
Q

Calculating moles of a gas (RTP)

A

n = vol/24 (dm3)

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18
Q

Ideal gas equation Conversions

A

kPa-> Pa X1000
cm3-> m3 X10^-6

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19
Q

Stoichiometry

A

Ratio of moles in a chemical reaction

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20
Q

Why is a high atom economy good?

A

-efficient
-produce little waste
-less raw materials used
-sustainable

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20
Q

Atom economy=

A

(Sum of Mr of desired products/ sum of molar mass of all products) x100

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21
Q

Percentage yield=

A

(Actual/ theoretical) x100

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22
Q

acids

A

Release H+ ions when dissolved in water
Proton donor

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22
Q

Strong acids

A

Fully dissociate when dissolved in water

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23
Q

Weak acids

A

Partially dissociate when dissolved in water

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24
Q

Bases

A

Accept H+ ions

Proton acceptors

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25
Q

Examples of bases

A

-metal oxides
-metal hydroxides
-metal carbonates
-alkalis

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26
Q

alkalis

A

Dissolve in water releasing OH- ions

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27
Q

Neutralisation

A

Reaction between acid and base to produce a salt

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28
Q

Neutralisation by carbonate

A

Acid + carbonate -> salt + water + carbon dioxide

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29
Q

Neutralisation by metal oxide

A

Acid + metal oxide -> salt + water

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30
Q

Neutralisation by alkali

A

Acid + alkali -> salt + water

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30
Q

Acid + ammonia

A

Ammonium salt

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31
Q

Standard solution

A

Solution of known concentration

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32
Q

Preparing standard solution of NaOH

A

-mass (e.g. 1.00 g) weighed out and added to beaker
-dissolve NaOH with distilled water use stirring rod
-pour into 250 cm3 volumetric flask
-rinse beaker with distilled water, wash rinsings into flask
-add distilled water until bottom of meniscus on graduation line
-stopper, invert, thoroughly mix

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33
Q

How to find mass needed to prepare standard solution

A

Find moles (n = cv)
Find mass (m = Mn)

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33
Q

Oxidation number ions

A

Monatomic- oxidation no= charge

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33
Q

Carrying out acid-base titration
(Finding conc H2SO4 reacting with 25cm3 NaOH)

A

-Pipette, add 25.0 cm3 NaOH into conical flask
-white tile, few drops of indicator (phenolphthalein or methyl orange)
-burette, sulfuric acid, initial reading, nearest 0.05 cm3
- add sulfuric acid, swirl
-colour change, final reading
- final- Initial = trial titre
-repeat, dropwise near end point until concordant results (within 0.1)
-mean using concordant results
-calculations

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34
Q

Ethanoic acid

A

CH3COOH

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35
Q

Oxidation number elements

A

0

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35
Q

Titration finding concentration

A

Balance equation
Mol known
Mol unknown
Conc unknown

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36
Q

Oxidation no Compounds

A

Halides -1
H +1
O -2

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37
Q

Oxidation no Exceptions

A

Metal hydride H-1
F2O O +2
Peroxide (H2O2) O -1

38
Q

Oxidation no compounds + polyatomic ions

A

Sum of oxidation numbers in compound = 0

Sum of oxidation no in polyatomic ion = overall charge

39
Q

How are oxidation numbers shown in elements with variable oxidation numbers?

A

Roman numerals
E.g. sodium chlorate (I) cl +1

40
Q

Oil rig

A

Oxidation is loss of electrons
Reduction is gain of electrons

41
Q

Redox in terms of hydrogen

A

Oxidation is loss of hydrogen

42
Q

Redox in terms of oxygen

A

Oxidation is gain of oxygen

43
Q

Redox in terms of oxidation no

A

Oxidation is increase in oxidation number

44
Q

What are the 4 types of orbitals?

A

S,p,d,f

45
Q

Shape of s orbital

A

Spherical

46
Q

Shape of p orbital

A

Dumbell
(3D axis)

47
Q

What is an orbital?

A

Region around Nucleus where an electron is most likely to be containing 2 electrons with opposite spins

47
Q

Number of electrons in first four shells

A

2,8,18,32

48
Q

Sub shells

A

Orbitals within energy level grouped together

49
Q

Metals form ——- ions

A

Positive (cations)

49
Q

Exceptions- electron structure

A

Cr and Cu
Less energy to fill 3d and have 1 4s
Or have 5 3d (all occupied singularly) and 1 4s

50
Q

Electron configuration

A

Electrons fill orbitals in increasing energy
4s lower energy than 3d
Maximum of 2 electrons with opposite spin represents by up and down arrow
Occupied singularly before pairing

51
Q

Non-metals form ——— ions

A

Negative (anions)

52
Q

Ionic bonding definition

A

Electrostatic attraction between oppositely charged ions

53
Q

Structure of ionic compounds

A

Giant ionic lattices, ions fixed in place

54
Q

Properties of ionic compounds

A

High mp and bp (lots of energy needed to overcome strong ionic bonds)
Soluble in polar solvents (water)
Insoluble in non-polar solvents (hydrocarbons)

55
Q

Strength of ionic bonds

A

Smaller ions form stronger bonds
Highly charged ions form stronger bonds

56
Q

Conduct because

A

Moving charged particle (electron, ion)

57
Q

Covalent bond definition

A

Attraction between positive nuclei and shared pair of electrons

58
Q

Dative covalent bond

A

Both electrons contributed by one atom

59
Q

Lone pairs

A

Pair of valence electrons nor bonded to anther atom

Able to form dative covalent bonds with atoms that have vacant orbitals

60
Q

Valence

A

Outer shell

61
Q

Average bond Enthalpy

A

Measure of covalent bond strength

Larger= stronger

62
Q

Salt

A

H+ ion replaced by metal ion or NH4+ ion

63
Q

Displayed formula

A

Bonds represented by lines

64
Q

Covalent bonding happens between ….

A

Non-metals

65
Q

Dative covalent bond displayed formula

A

Arrow

66
Q

Electron deficient

A

Less than 8 electrons

67
Q

Expanded octet

A

More than 8 electrons

68
Q

valence shell electron pair repulsion theory

A

e- repel as far apart as possible

69
Q

How much does each lone pair reduce the angle by??

A

≈2.5 °

70
Q

LP> LP-BP> BP

A

repulsion

71
Q

2 bonding regions
2bp, 0lp

A

linear 180 °

72
Q

3 bonding regions
3bp, olp

A

trigonal planar 180 °

73
Q

3 bonding regions
2bp, 1lp

A

non-linear 112.5 °

74
Q

4 electron pairs
4bp, 0lp

A

tetrahedral 109.5°

75
Q

4 electrons pairs
3bp, 1lp

A

pyramidal 107°

76
Q

4 electron pairs
2bp, slp

A

non-linear 104.5°

e.g. H2O

77
Q

5 bonding regions
5bp 0lp

A

trigonal bipyramidal
120°
90°

77
Q

5 bonding regions
4bp, 1lp

A

pyramidal 119° 89°
see-saw 119° 89°

78
Q

5 bonding regions
3bp, 2lp

A

trigonal planar 120 °
T-shaped 89°

79
Q

6 electron pairs
6bp, 0lp

A

octahedral 90°

80
Q

VESPA

A

valance shell electron pair repulsion theory

81
Q

6 electron pairs
4bp, 2lp

A

square planar 90°

81
Q

6 electron pairs
5bp, 1lp

A

square pyramid 89°

82
Q

method for determining shapes of elements

A

eg BrF3
central atom Br
outer e- 7
e- gained from bonds 3
e- gained from change 0
total e- 10
bp 3
lp 2

83
Q

shapes of ions; SO4 2-

A

2 x S=O
2 x S–O

two atoms with single bond O have extra e- in outer shell resulting in 2- charge

tetrahedral shape 109.5

84
Q

what can be said about the bonding affects of single bonds to double bonds?

A

they are similar

85
Q

covalent bond H-H

A

attract electrons equally
electron equally shared

86
Q

electronegativity

A

the ability of an atom to attract the bonding electrons in a covalent bond

87
Q

trends in electronegativity

A

F most electronegative
decreases down group
increases across period

88
Q

polar bond

A

pair of e- shared unequally
δ- slightly negative
δ+ slightly positive

89
Q

dipole

A

separation of partial charges in a molecule

90
Q

polar molecules

A

polar bonds (regions with different e- densities)
non-symmetrical
dipoles do not cancel

91
Q

non-polar molecules

A

if polar bonds: symmetrical, dipoles cancel

92
Q

types of intermolecular forces

A

induced dipole-dipole interactions (London forces)
permanent dipole-dipole interactions
hydrogen bonds

93
Q

london forces

A

*e- move randomly, at any instant in time, distribution of e- may be uneven
*this creates weak instantaneous dipoles δδ-
*induces dipole in neighbouring molecules

94
Q

permanent dipole-dipole interactions

A

2 polar molecules with permanent dipoles attract

95
Q

hydrogen bonds

A

strong permanent dipole-dipole interactions
H and O/F/N

96
Q

hydrogen bond diagrams

A

dipoles labelled
h in one molecule bonded to lone pair in another
hydrogen bond represented by dotted line
straight line

97
Q

ice less dense than water

A

regular structure, open lattice
when ice melts, the open lattice structure collapses allowing H2O molecules to move closer together increasing density

98
Q

water expands when it freezes

A

tetrahedral structure maximising hydrogen bonding
when ice melts, crystal structure melts as water molecules fall into empty spaces

99
Q

H2O higher melting point than expected

A

hydrogen bonds

100
Q

Aufbau principle

A

e- enter lowest energy level available

101
Q

Hund’s rule

A

e- prefer to occupy orbitals on their own and can only pair up when no orbital of the same energy is available (bus analogy)

occupy singularly before pairing

A-level chemistry (7 decks)

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