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Advanced Higher Chemistry > Last Minute Cram - Inorganic > Flashcards

Last Minute Cram - Inorganic Flashcards

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
Hund's rule
for degenerate orbitals, electrons fill each orbital singly before pairing starts
26
pauli exclusion principle
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)
27
ionisation energy across a period
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***
28
isoelectronic
particles which have the same electronic configuration
29
unusual electronic configurations
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
30
order of orbital filling | *note: always write in order of principal quantum number with accompanying s.p.d.f*
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p
31
ground state
electrons in lowest possible energy level
32
how is the periodic table divided into 4 blocks
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
33
which subshell empties first when forming ions (d block)
4s before 3d
34
when will a covalent bond form?
when atoms rearrange their electrons (by sharing) to produce an arrangement of lower energy
35
when will an ionic bond form?
when atoms can rearrange their electrons (transfer) to produce an arrangement of lower energy
36
confirmation of bonding type
electronegativity values are a useful guide but it is ultimately the properties which provide confirmation
37
lewis diagrams
explain properties associated with covalent bonding | shows bonding and non-bonding pairs (lone pair)
38
how a dative covalent bond forms
atoms use a lone pair as a bonding pair eg: NH4+
39
how to calculate no. of electron pairs
add number of outer electrons in the central atom to the number of bonded electrons from the outer atoms and divide by 2
40
how to determine shape of molecules and ions
if +ve, subtract electron from total for each positive charge and then divide by 2 (for -ve, same but add)
41
how to predict molecular shape
whichever 3D geometry minimises electron repulsion
42
repulsive forces
bonding electrons do not repel as much as non-bonding electrons because they are attracted by two nuclei
43
lone pairs and electron pair repulsion
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
order of strength of repulsions
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
rule for bond angle in tetrahedral molecules
each lone pair decreases bond angle by 2.5
46
transition metals
found in the d-block of the table | have an incomplete d subshell (excluding Cu and Zn)
47
transition metal ions
transition metals also have at least one ion with an incomplete d subshell (apart from Sc)
48
3 main characteristics of a transition metal
1. can produce ions with different valencies 2. produce coloured compounds 3. act as catalysts **Sc and Zn are exceptions**
49
transition metal varying oxidation states
have variable oxidation states of differing stability | can lose 4s and some or all of 3d electrons
50
transition metal ion examples
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
oxidation numbers rules
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
ox and red with oxidation numbers
oxidation involves an increase in oxidation number | reduction involves a decrease in oxidation number
53
oxidation numbers and ox and red agents
compounds containing metals with high oxidation states tend to be oxidising agents low oxidation state - reducing agent
54
oxidation state of vanadium and colours
+5 yellow +4 blue +3 green +2 violet *mnemonic: You Better Get Vanadium*
55
transition metal complexes
metal (atom or ion) surrounded by ligands, with the ligands bonded to the metal through dative covalent bonds
56
ligands
ion, atom or molecule which may be negative ions or have lone pairs electron donors
57
monodentate ligands
Cl- CN NH3 H20
58
bidentate ligands
Oxalate ion (C2H4) ethylenediamine Oxalic acid
59
hexadentate ligands
EDTA
60
coordination number
total number of dative covalent bonds to the metal atom or ion
61
monodentate meaning
ligands that form one coordinate bond to a metal atom or ion
62
naming complexes - chemical formula
1. enclosed in square brackets 2. metal symbol written first 3. negative ligands next 4. then neutral ligands
63
naming complexes - systematic name
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
negative ligand names
``` chloride - chlorido oxide - oxo cyanide - cyanido oxalate - oxolato ammonia - ammine water - aqua carbon monoxide - carbonyl hydroxide - hydroxido ```
65
negative complex ion name examples
iron - ferrate copper - cuprate lead - plumbate
66
colour observed transition metal
observed colour is the complimentary colour to that absorbed by the compound
67
what determines colour of transition metal?
interactions between ligands and electrons occupying d-orbitals create the circumstances which leads to the absorption of some light from the visible spectrum
68
splitting of d orbitals
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
order of splitting
(increase in splitting) | I- Br- Cl- F- H20 NH3 CN-
70
size of energy gap between two sets of d orbitals depends on...
the transition metal ion the oxidation state of the transition metal the type of ligand
71
complexes more likely to absorb in the visible region
weak field ligands | eg water
72
what can be used to determine concentration of the ions?
colorimetry | a filter of the complementary colour od the solution being tested at a defined wavelength is used in the colorimeter
73
complexes more likely to absorb in the UV region
strong field ligands eg cyanide ion compounds will be colourless so UV spectroscopy is used
74
which ions can have no d-d transitions
ions that have no d electrons or ions that have a complete d subshell
75
heterogeneous catalysts
different physical state to the reactants | eg iron in haber process
76
homogeneous catalysts
same physical state as the reactants | eg enzymes catalysing reaction in the body
77
what allows intermediates to form with reacting molecules?
unpaired d electrons or empty d orbitals these provide alternative reaction pathways with a lower activation energy.
78
what can Uv/visible spectroscopy be used for?
a quantitative method of analysis to determine the concentration and hence the mass of a transition metal in a compound or alloy
79
why can transition metals act as catalysts?
due to the variable oxidation states, they have d orbitals to form intermediates on catalyst active sites
80
example of transition metal catalyst
Co2+ pink --> Co3+ green --> Co2+ pink reactants intermediate products regenerated at end

Advanced Higher Chemistry (12 decks)

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