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Advanced Higher Chemistry > Last Minute Cram - Inorganic > Flashcards

Last Minute Cram - Inorganic Flashcards

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1
Q

what can the em radiation be described as?

A

a wave and particle - dual nature

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2
Q

equations for em spectrum

A

c=fλ
E=hf
E=Lhc/λ

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3
Q

order of em spectrum (low f to high f)

A 
radio
microwave
infrared
visible
ultraviolet
x-ray
gamma
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4
Q

what can emission and absorption spectra be used for?

A

to identify and quantify the element

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5
Q

atomic absorption spectroscopy

A

em radiation directed at an atomised sample

radiation absorbed as electrons are promoted to higher energy levels

absorption spectrum produced by measuring how the intensity of absorbed light varies with wavelength

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6
Q

atomic emission spectroscopy

A

high temperatures used to excite electrons to higher energy levels

as electrons drop to lower energy levels, photons are emitted

an emission spectrum of a sample is produced by measuring the intensity of light emitted at different wavelengths

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7
Q

what is the concentration of an element within a sample related to?

A

the intensity of light emitted or absorbed

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8
Q

lyman series

A

electrons dropping to the ground state

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9
Q

balmer series

A

electrons dropping to n=2

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10
Q

as energy increases…

A

levels get closer together and converge

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11
Q

what does the line of greatest energy represent?

A

electrons returning from the outermost shell to the ground state

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12
Q

if slightly more energy than the line of greatest energy?

A

electron removed

i.e. 1st ionisation energy

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13
Q

how to calculate ionisation energy

A

use the convergence limit in E=lhc/λ

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14
Q

principle quantum number

A

indicates the main energy level for an electron and is related to the size of the orbital
nearest nucleus n=1 and so on…

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15
Q

angular quantum number

A 
determines the shape of the subshell and has values from zero to n-1
l=0 s orbital
l=1 p
l=2 d
l=3 f
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16
Q

magnetic quantum number

A

determines the orientation of the orbital and has values between -l and l

use px,py and pz

3 possible p orbitals: -1,0,+1
5 d
7f

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17
Q

spin magnetic quantum number

A

determines the direction of the spin and can have values of +1/2 or -1/2

clockwise or anti-clockwise

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18
Q

orbital

A

region of space with a 90% probability of finding an electron

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19
Q

s orbital

A

spherical
diameter increases as shell no. increases
the only orbital in shell 1

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20
Q

p orbital

A

dumb-bell shaped
only occur from second shell onwards
all have equal energy - degenerate
each p orbtial can hold 2 electrons (px+py+pz=6)

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21
Q

maximum number of electrons in a single orbital

A

2

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22
Q

d orbitals

A

each shell from the third shell contains 5 d-orbitals

orientation:
between x and y axis (dxy)
between x and z axis (dxz)
between y and z axis (dyz)
along x and y axis (dx^2-y^2)
along z axis (dz^2)
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23
Q

aufban diagram

A

orbitals ranked in terms of energy

the third and fourth shells overlap with electrons occupying the 4s orbital before the 3d one

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24
Q

aufbau principle

A

electrons will fill orbitals in order of increasing energy

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25
Q

Hund’s rule

A

for degenerate orbitals, electrons fill each orbital singly before pairing starts

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26
Q

pauli exclusion principle

A

maximum number of electrons in any atomic orbital is two and if there are two electrons in the same orbital, they must have opposite spins (no 2 electron can have the same quantum numbers)

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27
Q

ionisation energy across a period

A

in general, 1st IE increases across a period

exception: 1st IE of Boron is lower than Berylium

explanation: Berylium has a full sub-shell which is a stable arrangement
Boron has a single 2p electron which is less stable and easier to remove

special stability with half-filled and full subshells

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28
Q

isoelectronic

A

particles which have the same electronic configuration

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29
Q

unusual electronic configurations

A

chromium
2,8,13,1
1s2 2s2 2p6 3s2 3p6 3d5 4s1 instead of 3d44s2
stability of filled and half-filled orbitals
one electron in each d-orbital
symmetry around the nucleus

copper
2,8,18,1
1s2 2s2 2p6 3s2 3p6 3d10 4s1 instead of 3d9 4s2
symmetry around nucleus

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30
Q

order of orbital filling

note: always write in order of principal quantum number with accompanying s.p.d.f

A

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p

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31
Q

ground state

A

electrons in lowest possible energy level

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32
Q

how is the periodic table divided into 4 blocks

A

depending on which subshell the highest energy electrons are found in

s block - group 1 and 2 (and He)
p block - groups 3-7
d block - transition metals
f block - lanthanide and actinide series

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33
Q

which subshell empties first when forming ions (d block)

A

4s before 3d

34
Q

when will a covalent bond form?

A

when atoms rearrange their electrons (by sharing) to produce an arrangement of lower energy

35
Q

when will an ionic bond form?

A

when atoms can rearrange their electrons (transfer) to produce an arrangement of lower energy

36
Q

confirmation of bonding type

A

electronegativity values are a useful guide but it is ultimately the properties which provide confirmation

37
Q

lewis diagrams

A

explain properties associated with covalent bonding

shows bonding and non-bonding pairs (lone pair)

38
Q

how a dative covalent bond forms

A

atoms use a lone pair as a bonding pair eg: NH4+

39
Q

how to calculate no. of electron pairs

A

add number of outer electrons in the central atom to the number of bonded electrons from the outer atoms and divide by 2

40
Q

how to determine shape of molecules and ions

A

if +ve, subtract electron from total for each positive charge and then divide by 2
(for -ve, same but add)

41
Q

how to predict molecular shape

A

whichever 3D geometry minimises electron repulsion

42
Q

repulsive forces

A

bonding electrons do not repel as much as non-bonding electrons because they are attracted by two nuclei

43
Q

lone pairs and electron pair repulsion

A

to minimise electrons pair repulsion, lone pairs are always placed in equatorial positions to maximise distance from bonding pairs
eg: ClF3 would be T-shaped

44
Q

order of strength of repulsions

A

LP-LP > LP-BP > BP-BP

Lone pairs are held closer to the central atom, so they have a greater repulsion than bonding pairs

45
Q

rule for bond angle in tetrahedral molecules

A

each lone pair decreases bond angle by 2.5

46
Q

transition metals

A

found in the d-block of the table

have an incomplete d subshell (excluding Cu and Zn)

47
Q

transition metal ions

A

transition metals also have at least one ion with an incomplete d subshell (apart from Sc)

48
Q

3 main characteristics of a transition metal

A
  1. can produce ions with different valencies
  2. produce coloured compounds
  3. act as catalysts

Sc and Zn are exceptions

49
Q

transition metal varying oxidation states

A

have variable oxidation states of differing stability

can lose 4s and some or all of 3d electrons

50
Q

transition metal ion examples

A

during rusting, Fe3+ ion favoured because of extra stability of d-orbitals being half-filled

Cu+ ion favoured because even greater stability due to ions forming in solution

51
Q

oxidation numbers rules

A
  1. simple ions (Na+ etc.) continue to count as -1 or +1
  2. oxygen always -2
  3. Hydrogen usually +1 unless it forms metal hydride (-1)
  4. overall charge of a compound is 0
  5. in polyatomic ions, the sum of all oxidation numbers is equal to the overall charge of the ion
52
Q

ox and red with oxidation numbers

A

oxidation involves an increase in oxidation number

reduction involves a decrease in oxidation number

53
Q

oxidation numbers and ox and red agents

A

compounds containing metals with high oxidation states tend to be oxidising agents

low oxidation state - reducing agent

54
Q

oxidation state of vanadium and colours

A

+5 yellow
+4 blue
+3 green
+2 violet

mnemonic: You Better Get Vanadium

55
Q

transition metal complexes

A

metal (atom or ion) surrounded by ligands, with the ligands bonded to the metal through dative covalent bonds

56
Q

ligands

A

ion, atom or molecule which may be negative ions or have lone pairs

electron donors

57
Q

monodentate ligands

A

Cl-
CN
NH3
H20

58
Q

bidentate ligands

A

Oxalate ion (C2H4)
ethylenediamine
Oxalic acid

59
Q

hexadentate ligands

A

EDTA

60
Q

coordination number

A

total number of dative covalent bonds to the metal atom or ion

61
Q

monodentate meaning

A

ligands that form one coordinate bond to a metal atom or ion

62
Q

naming complexes - chemical formula

A
  1. enclosed in square brackets
  2. metal symbol written first
  3. negative ligands next
  4. then neutral ligands
63
Q

naming complexes - systematic name

A
  1. ligands named first in alphabetical order
  2. then the name of the metal
  3. if ligand is a negative ion the ending changes from ‘ide’ to ‘ido’
  4. more than one ligand = prefixes i.e. di,tri etc.
  5. if complex ion is overall negative suffix ‘ate’ is added. Latin name is also used
  6. oxidation state given in roman numerals after its name
64
Q

negative ligand names

A 
chloride - chlorido
oxide - oxo
cyanide - cyanido
oxalate - oxolato
ammonia - ammine
water - aqua
carbon monoxide - carbonyl
hydroxide - hydroxido
65
Q

negative complex ion name examples

A

iron - ferrate
copper - cuprate
lead - plumbate

66
Q

colour observed transition metal

A

observed colour is the complimentary colour to that absorbed by the compound

67
Q

what determines colour of transition metal?

A

interactions between ligands and electrons occupying d-orbitals create the circumstances which leads to the absorption of some light from the visible spectrum

68
Q

splitting of d orbitals

A

in transition metals, d orbitals are no longer degenerate
as the complex forms, d orbitals split and the orbitals that lie along the axis are repelled and promoted to a higher energy

69
Q

order of splitting

A

(increase in splitting)

I- Br- Cl- F- H20 NH3 CN-

70
Q

size of energy gap between two sets of d orbitals depends on…

A

the transition metal ion
the oxidation state of the transition metal
the type of ligand

71
Q

complexes more likely to absorb in the visible region

A

weak field ligands

eg water

72
Q

what can be used to determine concentration of the ions?

A

colorimetry

a filter of the complementary colour od the solution being tested at a defined wavelength is used in the colorimeter

73
Q

complexes more likely to absorb in the UV region

A

strong field ligands
eg cyanide ion

compounds will be colourless so UV spectroscopy is used

74
Q

which ions can have no d-d transitions

A

ions that have no d electrons or ions that have a complete d subshell

75
Q

heterogeneous catalysts

A

different physical state to the reactants

eg iron in haber process

76
Q

homogeneous catalysts

A

same physical state as the reactants

eg enzymes catalysing reaction in the body

77
Q

what allows intermediates to form with reacting molecules?

A

unpaired d electrons or empty d orbitals

these provide alternative reaction pathways with a lower activation energy.

78
Q

what can Uv/visible spectroscopy be used for?

A

a quantitative method of analysis to determine the concentration and hence the mass of a transition metal in a compound or alloy

79
Q

why can transition metals act as catalysts?

A

due to the variable oxidation states, they have d orbitals to form intermediates on catalyst active sites

80
Q

example of transition metal catalyst

A

Co2+ pink –> Co3+ green –> Co2+ pink
reactants intermediate products

regenerated at end

Advanced Higher Chemistry (12 decks)

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