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a level chemistry > module 2: electronic configuration > Flashcards

module 2: electronic configuration Flashcards

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1
Q

what are the rules for assigning oxidation numbers

A
  • O is +2 with F
  • O is -1 in a peroxide
  • H is -1 with metals
  • aluminium = +3 always
  • g1 = +1 always
  • g2 = +2 always
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2
Q

what is an acid

A

a species that is a proton donor

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3
Q

what is a base

A

a species that is a proton acceptor

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4
Q

what is an alkali

A

a water soluble base

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5
Q

what is a strong acid

A

an acid that fully dissociates into a solution of H+ ions

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6
Q

what is disproportionation

A

when one element in a reaction is both oxidised and reduced

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7
Q

in a redox reaction between an acid and a metal what is oxidised/reduced

A

metal is oxidised => positive metal ion
hydrogen from acid is reduced => hydrogen as an element

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8
Q

acid + metal –>

A

acid + metal –> metal salt + hydrogen

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9
Q

acid + alkali –>

A

acid + alkali –> metal salt + water

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10
Q

acid + metal oxide –>

A

acid + metal oxide –> metal salt + water

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11
Q

acid + metal hydroxide –>

A

acid + metal hydroxide –> metal salt + water

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12
Q

acid + metal carbonate –>

A

acid + metal carbonate –> metal salt + water + carbon dioxide

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13
Q

what is an oxidation number

A

a measure of the number of electrons that an atom uses to bond with the atoms of another element

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14
Q

what does a reduction/increase in oxidation state mean

A

reduction in os = species was reduced
increase in os = species was oxidised

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15
Q

what are the 2 disproportionation reactions that you need to know?

A

Cl2(aq) + H2O –> HClO(aq) + HCl(aq)
chlorine is added to water to kill bacteria to make water safe to drink, hydrochloric and chloric acid are formed

Cl2(aq) + 2NaOH(aq) –> NaCl(aq) + NaClO(aq) + H2O(l)
bleach is formed when cold, dilute, aqueous sodium hydroxide and chlorine react at room temperature

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16
Q

what is first ionisation energy

A

the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ atoms

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17
Q

what are the 3 factors that affect ionisation energy

A
  1. nuclear charge
  2. distance from nucelus
  3. electron shielding
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18
Q

how does ionisation energy change as you go down group 2

A

it decreases
the number of filled shells increases as you go down a group
so shielding and distance from nucleus increases, making it easier to remove the outermost electron

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19
Q

what is an orbital

A

a region within an atom that can hold up to two electrons, with opposite spin

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20
Q

outline an s orbital

A

its spherical
each shell will contain an s orbital
the greater value of n, the greater the s-orbital radius

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21
Q

outline p-orbitals

A

3D dumb-bell shape
from the 2nd shell onwards, each shell contains 3 p-orbitals
the greater the value of n, the further the p-orbital from the nucleus

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22
Q

how does ionic bonding occur

A

a type of bonding where one atom loses electrons and one atom gains them

it occurs between a metal and a non-metal, atoms become charged and are attracted to each other

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23
Q

what is covalent bonding

A

a type of bonding where atoms share electrons with each other

it usually occurs between non-metal atoms, diatomic molecules are formed through covalent bonding

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24
Q

what is metallic bonding

A

a type of bonding where a sea of free electrons is created

it occurs between metal atoms, electrons become delocalised and more free to move

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25
Q

what are properties of ionic compounds

A
  • solid at room temperature
  • high melting and boiling points; strong electrostatic forces of attraction
  • soluble
  • conducts when molten/in solution
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26
Q

do ionic compounds conduct electricity

A

only when molten or in solution
they are made of charged particles but electricity is the movement of these, they can move when the substance is a liquid

the ionic lattice dissolves in a polar substance like water

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27
Q

what is a dative covalent bond

A

a shared electron pair that has been provided by only one of the bonding atoms; the shared pair was originally a lone pair in one of the bonded atoms

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28
Q

why does dative covalent bonding occur

A

it can be energetically favourable for one species to donate a lone pair to another atom/ion, both species are left with full subshells

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29
Q

properties of simple covalent substances

A
  • low melting and boiling point; weak intermolecular forces don’t need a lot of energy to be broken even though they have strong covalent bonds
  • don’t conduct electricity; lack of charged particles
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30
Q

properties of metals

A
  • electrical conductors

charge is carried by delocalised electrons

  • high melting and boiling points; strong forces of electrostatic attraction between oppositely charged ions need to be overcome
  • insoluble
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31
Q

what is electron pair repulsion theory + what is relative repulsion

A

bonded electron pairs exert a repulsion onto other repulsion pairs
this causes electron pairs to move as far apart as possible
determining the 3 dimensional shape of the molecule

lonepair-lonepair > lonepair-bondpair > bondpair-bondpair

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32
Q

what are the bond angles in CH4

A

109.5°

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33
Q

what are the bond angles in NH3

A

107°

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34
Q

what are the bond angles in H2O

A

104.5°
there are two bonding pairs and two lone pairs

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35
Q

what is the shape name and bond angle with 4 bonding pairs and no lone pairs

A

tetrahedral, 109.5°

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36
Q

what is the shape and bond angle with 2 bonding pairs and no lone pairs

A

linear, 180°

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37
Q

what is the shape name and bond angle with 3 bonding pairs and no lone pairs

A

trigonal planar, 120°

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38
Q

what is the shape name and bond angles with 5 bonding pairs and no lone pairs

A

trigonal bipyramidal, 120° and 90°

39
Q

what is the shape and bond angles with 6 bonding pairs and no lone pairs

A

octahedral, 90°

40
Q

what is the shape name and bond angle with 3 bonded pairs and one lone pair

A

pyramidal, 107°

41
Q

what is the shape name and bond angle with 2 bonding pairs and 1 or more lone pairs

A

non-linear, bent, 104.5°

42
Q

define electronegativity

A

a measure of an atom’s tendency to attract a bonding pair of electrons in a covalent bond

43
Q

what are the factors affecting electronegativity

A

atomic charge
atomic radius
electron shielding

44
Q

how does electronegativity change across a period and why

A

nuclear charge increases
atomic radius decreases
electron shielding stays constant

overall increase

45
Q

how does electronegativity change down group 2

A

overall decrease

the bonding pair of electrons is increasingly distant from the attraction of the nucleus

46
Q

what is used to assign values for electronegativity

A

the Pauling Scale

47
Q

what makes a molecule polar vs non-polar

A

it has to be an asymmetrical molecule, the dipoles shouldn’t cancel out each other

symmetrical molecules are non-polar because the dipoles cancel out even though they contain polar bonds

it isn’t dependent on whether or not a molecule contains polar bonds

48
Q

what are the three most common intermolecular forces, in order of strongest to weakest

A

hydrogen bonds
permanent dipole-dipole interactions
induced dipole-dipole interactions

49
Q

what is an induced dipole-dipole interaction

A

where an instantaneous dipole induces a dipole in another molecule, these attractions can occur between any molecule (even non polar molecules)

because electrons move at very high speeds in orbitals and can lie on one side of an atom at any given point, causing an instantaneous dipole that produces an induced dipole in a neighbouring molecule/atom that is attracted

50
Q

how does increasing the no. of electrons in a molecule impact the induced dipole-dipole interactions

A

larger instantaneous and induced dipole –> greater induced dipole interactions and stronger attractive forces between molecules, will require more energy to overcome the forces

51
Q

permanent dipole-dipole interaction

A

a weak attractive force between permanent dipoles in neighbouring polar molecules

polar molecules can also have induced dipole-dipole interactions at the same time, leading to even higher boiling points

52
Q

what is hydrogen bonding

A

a strong dipole-dipole attraction between an electron deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom (O, N and sometimes F) on a different molecule

53
Q

what is an oxidising agent vs a reducing agent

A

a reducing agent gets oxidised
an oxidising agent gets reduced

54
Q

what are 3 definitions of reduction

A

gain of electrons
loss of oxygen
gain of hydrogen

55
Q

how many electrons can each of the first 3 shells hold

A

shell 1 - 2
shell 2 - 8
shell 3 - 18

56
Q

why are groups 1 and 2 known as s-block elements

A

because the highest energy electrons are found in the s subshell

56
Q

what is a covalent bond

A

a shared pair of electrons and a strong electrostatic attraction

57
Q

first ionisation energy equation with an acid

A

H2SO4 (aq) —> HSO4- + H+

remove a H+

58
Q

what makes a branched hydrocarbon have a lower boiling point than a straight chain one?

A

branched has weaker induced and permanent dipole-dipole forces between molecules because the molecules aren’t packed as closely together meaning that there are less surface interactions

59
Q

how is dative covalent bonding shown in diagrams

A

using an arrow that points away from the ‘donor atom’

60
Q

what is the electronic configurations of chromium

A

Cr is: [Ar] 3d5 4s1 not [Ar] 3d4 4s2

61
Q

what equation is used when trying to calculate heat energy and units

A

Q = mc∆T
Heat energy (J) = Mass(of the thing changing temp)(g) × SHC × change in temperature (C)

62
Q

what determines the order of elements in the periodic table

A

number of protons

63
Q

what are the giant covalent structures

A

boron
carbon
silicon

64
Q

are intermolecular forces or intramolecular forces stronger?

A

intramolecular forces are stronger
(a H bond is like a tenth the strength of a covalent bond)

65
Q

how is percentage uncertainty calculated

A

% uncertainty = (uncertainty x no. times read) ÷ measured value x 100

66
Q

what is volumetric analysis

A

a process that uses the volume and concentration of one chemical reactant (standard solution) to determine the concentration of another, unknown solution
most commonly done via titration

67
Q

how is a standard solution made

A

they are made as accurately and precise as possible by using 3 d.p balances and volumetric flasks to minimise measurement uncertainties

  1. precisely weigh out solid
  2. add to a small volume of water and pre-dissolve the solid
  3. transfer to a volumetric flask with a funnel
  4. rinse the beaker and funnel with distilled water, adding the rinsing to the flask
  5. make up to the scratch mark with more distilled water, add the slobber and mix the contents
68
Q

how is a volumetric pipette used in titrations

A

to measure a known volume of a solution which is added to the conical flask

69
Q

how is a burette used in titrations

A

it holds one of the solutions and the tap is opened to release some slowly

70
Q

what are concordant results

A

within 0.1cm3 of each other

71
Q

what is the electronic configuration of copper

A

Cu is: [Ar] 3d10 4s1 not [Ar] 3d9 4s2

72
Q

what are ionic compounds made from

A

a metal and a non-metal

73
Q

what are covalent compounds made from

A

2 non-metals

74
Q

what ions are deflected the most in mass spectrometry

A

ions with a smaller m/z ratio

75
Q

what can increase the size of an induced dipole

A

number of electrons

76
Q

are double, single or triple bonds harder to break

A

single are easiest and it gets harder usually!

77
Q

what are the relative sizes of the cation and anion in an ionic lattice

A

cation is smaller because it’s lost the electron
anion is bigger because it gains electron

78
Q

what does a lewis structure of an ion always need

A

THE OVERALL CHAGRGW!!!!£ 😝😝😝😝😜😜🫵🫵🫵🫵🫵

79
Q

when going from Q(joules) to delta H neutralisation (kJmol-1) what molar value do you use

A

moles of water formed

80
Q

what should you talk about in a question on melting points

A

intermolecular forces!!
ie.
permanent dipole-dipole
hydrogen bonding
induced dipole-dipole

81
Q

why is 2nd ionisation energy usually higher

A

because the proton:e- ratio in a positive ion is greater than before

82
Q

what can you do when struggling to determine what element has a higher x ionisation energy

A

write out the electron configurations

83
Q

what are the units for temp in Q=mcdeltaT

A

celsius

84
Q

how do you calculate deltaH from Q

A

divide Q by the moles of the reactant that isn’t in excess

85
Q

what will the increase in temp be if the same concentrations of reactant are used, but the volumes are doubled (10deg before) and why

A

the same! because double the energy will be spread over double the volume

86
Q

how is a redox equation balanced

A

use oxidation numbers
1. determine oxidation numbers of every species
2. write half reactions with the oxidation numbers and e- to make up for charge differences
3. make the no. of e- equal and balance
4. use to balance actual

87
Q

what is the structure and bonding in a metal

A

giant lattice

88
Q

what is the structure and bonding in a covalent substance

A

simple molecular

89
Q

what is the structure and bonding in an ionic substance

A

giant lattice

90
Q

what states are used in ionisation energy equations

A

gas!!!!

91
Q

what is the type of compound when a question says that a gas is formed

A

carbonate and CO2 is being formed

92
Q

what to discuss when comparing ionisation energies

A
  • sub-shells (state them)
  • energy levels eg. 2p is a higher energy level than 2s
  • shielding
  • nuclear charge
  • attraction

a level chemistry (15 decks)

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