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inorganic sem1 > 1st John > Flashcards

1st John Flashcards

1
Q

how is the periodic table set out

A

order of atomic number

number of protons

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2
Q

what atoms in the periodic table are electronically similar

A

atoms in the same group

theyre chemically similar

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3
Q

group 18

A

noble gases

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4
Q

group 1

A

alkali metals

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5
Q

group 2

A

alkali earth metals

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6
Q

main group elements

A

outer shells consist of s and p orbitals

s or p block elements

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7
Q

d metals

A

partially full d orbital metals.
centre of periodic table

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8
Q

f metals consist of

A

lanthanides (4f)
actinides (5f)

bottom two rows

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9
Q

what can the schrodinger equation not do

A

find solutions for multi electron systems

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10
Q

what do we do bc the schrodinger cant find solutions to multi electron systems

A

we apply the hydrogen stuff to multi electron systems

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11
Q

what do we need to know when applying hydrogen to multi electron systems

A

same quantum numbers are used

orbitals have the same shapes(radial and angular wave functions)

radial functions are similar to hydrogen (depend on n and l tho)

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12
Q

pauli exclusion principle

A

electrons must have opposite spins
no 2 electrons in an atom can have the same 4 quantum numbers
n, l, ml, ms

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13
Q

aufbau principle

A

electrons are filled from lowest energy to highest energy

we remove from highest energy to lowest energy

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14
Q

hunds first rule

A

for degenerate orbitals, electrons must be added to different degenerate orbitals with parallel spins.

eg: fill in the 3 2p orbitals with 1 electron each parallel.

parallel spins are more stable as they minimise coloumbic repulsion

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15
Q

what does degenerate mean

A

the same in energy

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16
Q

what does the energy in H depend on

A

the principle quantum number
n

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17
Q

what does the energy in multi electron systems depend on

A

principle quantum number, n
angular quantum number, l

angular quantum number = azimuthal quantum number, l.

seen by radial distribution graph

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18
Q

what is the probability of finding an electron at distance r from nucleus in RDF proportional to

A

4 pi r squared psi squared

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19
Q

in rdf, is 1s max closer to nuc than 2s max

A

yes!!
meaning 1s shields 2s from nuc charge (not very well tho)
meaning 1s is a smaller orbital

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20
Q

in rdf how is 2s seen

A

m shape
smaller first hill
larger than 1s bc more further out
penetrates 1s orbital, e- are closer to nuc bc of this (although 2s dont rlly shield 1s)

1 radial node

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21
Q

what is the max in a rdf curve

A

the most probable area to find an electron

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22
Q

what is the value for psi at a radial node

A

0

means than e- are closer to the nuc

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23
Q

what is a core electron

A

electrons in a filled orbital
dont take part in chemical reactions
they shield valence electrons
low energy
close to nucleus

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24
Q

what is a valence electron

A

highest energy electron
they do chemsitry
get shielded by core electrons

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25
Q

order of orbital shielding strength

A

s
p
d
f

s shields best as it has a sphere shape.

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26
Q

what is shielding dependednt on in multi electron systems

A

principle quantum number
angular quantum number, l

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27
Q

shielding

A

electrons experience a lower nuclear charge (Zeff) due to e- closer to the nucleus repelling them.

depends on quantum number n.
larger n = larger shielding = lower Zeff

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28
Q

penetration

A

refers to the electron density of an electron around the nuc.

2s has more e- density near nuc than 2p so 2s is more penetrating

depends on principle and angular quantum number

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29
Q

electrons in orbitals haveee

A

different wavefunctions

therefore

different rdf

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30
Q

penetrating effect strength when n is the same

A

s
p
d
f

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31
Q

aufbau and why 2s is filled before 2p exp. rdf

A

2s penetrates 1s
meaning more e- density near nuc than 2p (and some of 1s)
mini max closer to nuc meaning lower in energy than 2p.

1s shields 2p more than 2s, 2s penetrates 2p and 1s, mini max closer to nuc, lower in energy

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32
Q

4s filled before 3d exp. rdf + aufbau

A

4s penetrates 3d
some 4s e- closer to nuc than 3d
filling 4s reduces shielding, 4s feels greater Zeff, lower in energy, more stable.

closer to nuc = lower energy = more stable

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33
Q

Zeffective = effective nuclear charge

A

Zeff = Z - S

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34
Q

what does the Zeff mean

A

the nuclear charge felt by a valence e- depends on what shell its on.

more shells between e- and nuc = more shielding = lower Zeff felt.

35
Q

how does Zeff change from left to right

A

increases
e- are added but in the same shell as previous ones (esp bc d orbital can fit 10 e-!!) meaning they dont shield the nuc charge as well as if they were closer to the nuc.

36
Q

what does the increase in Zeff across the period mean

A

increase in IE ( more energy needed to remove an electron )

decrease in atomic size ( nuc charge attracts e- more, stronger electrostatic forces of attraction)

37
Q

what is S in Zeff = Z - S

A

slaters rules: shielding constant
- same shell / same n = 0.35

  • shell below/ n-1 = 0.85
  • 2 shells below / n-2 = 1

if theyre d or f everything shields it by 1.

38
Q

how does Zeff change down a group

A

increases

39
Q

what is ionisation enery

A

minimum energy needed to remove an e- from a mole of gaseous atoms (making sure both ion and e- have 0Ke after)

40
Q

ionisation energies are always

A

positive +
energy is needed for this to happen

41
Q

ionisation energy equation

A

E(g) –> E(g)+. + e-

42
Q

why is the IE energy larger for the 2nd IE

A

removing an electron from a positively charged ion required more energy than removing it from a neutral one

43
Q

what group has the highest IR

A

noble gases

44
Q

which group has the smallest IR

A

alkali metals

45
Q

trend of IE across a period

A

increase IE
due to increase Zeff
e- exp more nuc charge, holding e- stronger to nuc, needs more energy to be removed.

46
Q

borons Ie is lower than Be

A

Be = 2s
B = 2s 2p

2s penetrates more, higher Zeff felt, held tighter to nuc, more energy needed to remove

2p = higher energy, easier to remove. larger l value.

47
Q

O IE less than N

A

N = 2p3
O = 2p4

O e- is paired up, less energy needed to remove e- due to repulsion DAOIER

48
Q

repuslion between 2 paired electrons

A

destabilising effect of interelectronic repulsion

49
Q

exchange energy

A

number of parallel electron pairs.
more pairs = more stable

additional stability in half filled or completely filled shell.

3 e- = 3 parallel pairs.

50
Q

trend of IE down a group

A

IE decreases
less energy needed to remove an e-

increase in n, not enough increase in Zeff

51
Q

Ga has higher IE than Al even with a higher n ?

A

e- are added but to 4p orbital (after 3d10 orbital) .
d orbital shield weakly.
4p penetrates 3d bc s>p>d
Ga 4p1 electron exp larg Zeff than Al 3p1 as s>p>d shielding and bc 4p penetrate 3d.

p cant pen s. s can pen anything.
look at what can pen, look at what can shield

52
Q

electron affinity

electron attachment energy

A

energy change when adding 1e- to a gasueous ion atom

opp of IE

+/- values

53
Q

electron affinity equation

A

E(g) + e- –> E(g) -

54
Q

group 17 electron affinity values

halogens

A
  • XXX
    NEGATIVE

no energy needed to add an e-.
they want the e-.

55
Q

group 18 electron affinity values

noble gases

A

+ XXX

energy is needed to add an e-.
they dont rlly want an e-.

56
Q

atom with largest electron affinity

A

Cl - chlorine
not F bc F is so small it repels the e-.

57
Q

left to right electron affinity

A

increases (more negative)
until noble gases (group 18)
bc Zeffffffff

58
Q

down a group electron affinity

A

decreases (more positive)
aka more energy needed to add an e- to it.

59
Q

why do group 1 have - electron affinity values.

aka no energy is needed to give an e-

A

bc u add into s orbital
low in energy
penetrate close to nuc
high electrostatic attraction

60
Q

why do group 2 have a + electron affinity value

aka energy is needed to give e-

A

bc u would be adding to the p orbital
higher in energy

higher l value.
energy of e- depend on n and l.

61
Q

why do group 15 have lower electron affinities than group 14

its harder to add electrons to group 15 than 14

A

in group 15, adding an e- will cause pairing of electrons in the p orbital.

these e- pairs will experience the destabilising effect of interelectronic repulsions.

in group 14, adding e- does not cause them to pair, theres 1 e- in each p orbital.

62
Q

why does the first period of p block have slightly lower electron affinities

A

their small size due to high Zeff causes the e- added to be repelled.
more energy is needed to add an e- to this row compared to the 2nd p block row.

lone pair lone pair repulsions.

63
Q

relationship between IE and electron affinity

A

generally: easier to add = easier to take away

drop for O and F in EA due to their small size causing repulsion.

64
Q

electronegativity

A

atoms ability in a molecule to attract a bonding pair of electrons in a covalent bond to itself.

65
Q

pauling electronegativity

A

A-A and B-B -> A-B
if enthalpy of AB is higher than AA or BB. – ionic contribution, nonpolar to polar.

big electronegative difference between A and B

x of F is 4.0

66
Q

muliken electronegativity

A

measure electronegativity based on IONISATION ENERGY and ELECTRON AFFINITY

scaled to match paulings scale.

67
Q

alfred rochow electronegativity

A

measure electronegativity based on effective nuclear charge and covalent radius

scaled to match pauling

68
Q

same electronegative atoms in a bond

A

bond is nonpolar
equal distribution of e- across the bond

both attract same amount of e-.

69
Q

one atom more eelctronegative than other in a bond

A

bond is polarised
more electronegative atom attracts more e-.

degree of polarisation is proportional to electronegativity.

70
Q

larger electronegative difference favours

A

making of ions
ionic bonding if they are more than 2 units different

71
Q

more / less electronegative atoms are more likely to form

A

ionic bonds

72
Q

atomic radii

A

difference between 2 nuclei

internuclear distance = bond length = affects a bonds energy

73
Q

diff atomic radi occurs when

A

same atom but different bonds are made.
covalent
ionic
vdw
metallic

74
Q

why is it impossible to find size of atom

A

only distance between nuc can be experimentally measured

probability of finding an e- far away from nuc is never 0. asympnotic wf or rdf??

even if distance from nuc was limited to x, radius of atom when measuring e- distance will change between different compounds.

75
Q

covalent radii - single bonds

A

use experimentally measured bond lengths (nuc to nuc)
then divide by 2 to get atomic radius

half ther interatomic distance in a homonuclear diatomic.

half Br - Br bond length

known radii can be used to find unknown ones if known is bonded to unknown.

76
Q

double and triple bonds are shorter than

A

single bonds

this must be remembered when caluclating atomic radi based on experimentally deduced single bond lengths.

77
Q

covalent radi trend

A

across table: decrease
larger Zeff contracts orbital size

down group: increases in size
higher ‘n’ value.

78
Q

how do Al and Ga have the same atomic radii

A

Zeff
Ga has e- in 4p but 3d shield weakly so 4p experiences strong Zeff.
4p also penetrate 3d
bc p>d

same contraction as Al

D BLOCK CONTRACTION

79
Q

filling of 4f orbitals as u go from In to Ti (n 5 –> 6) causing similar atomic radii

A

LANTHANDIE CONTRACTION

80
Q

metallic radii

A

one half of the internuclear distance in a metal
based on the atoms coordination number (its packing)

standardised to a 12 coordinate geometry

more neighbours = larger radius

81
Q

van der waals radii

A

1 half the internuclear distance between 2 atoms that arent bonded in solid or liquid state (so atoms are in contact)

only known for some elements
values can vary considerably.

82
Q

ionic radii

A

determined from crystal structures.
internuclear distance is the sum of anion + cation radii

we arent sure the precise distance of each. aka where the cation ends and where the anion starts.

anions radii are usually larger: less Zeff holding the extra e-.

83
Q

in ionic radii, what happens as the charge becomes more positive

A

ionic radius decreases
less electrons
stronger nuc charge
holds atom closer together

84
Q

what happens to ionic radii when the oxidation state increases

A

ionic radius decreases
charge of cation increases

inorganic sem1 (13 decks)

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