Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

chem hl > periodicity > Flashcards

periodicity Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

atomic radius trends

A

INCREASE DOWN the group

DECREASE ACROSS a period

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

atomic radius trend in a group & explain

A

radius INCREASE down the group

electrons occupied another energy level

valence electrons further from nucleus

[make specific]

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

atomic radius trend across the period & explain

A

radius DECREASE across the period

nuclear charge increase (with increasing number of proton)

all valence electrons in the same energy level

similar screening/shielding effect (same number of inner shell electrons)

valence electron experience stronger attractive force from the nucleus and gets pulled in

decrease atomic size

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

ionic radius: cations

predict whether Na or Na+ is larger in size

A

Na

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

what is effective nuclear charge

A

inner electron SHIELDS outer electrons from the nucleus

outer electron does not feel the full nuclear charge

attractions reduced by electron-electron repulsion

net nuclear charge felt → effective nuclear charge (Z eff)

(nuclear charge - number of inner shell electrons)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

ionic radius trend in group & explain (both anion and cation)

A

INCREASE down the group

(has more occupied electron shells)

valence electrons are further from nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

ionic radius trends

A

INCREASE DOWN the group

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

ionic radius: cations

predict whether S or S2- is larger in size

A

S2-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

ionic radius trends: size of cations & reason

A

cations are smaller than original neutral atom

main reason [1]: cation has one less occupied electron shell

has greater net positive charge/same number of protons pulling smaller number of electrons

smaller shielding effect

electron held more tightly

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

ionic radius trends: size of anions & reason

A

anions are bigger than original neutral atom

has greater net negative charge/same number of protons

electron held less tightly

TREND: increase down the group; has more occupied electron shells

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

ionic radius trends: across the period

A

CATIONS
decrease from group 1-14

isoelectronic (identical configuration), but increase in nuclear charge

ANIONS
decreased from group 14-17

isoelectronic, but increase in nuclear charge

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

is cation smaller or larger than anions & why?

A

SMALLLER

cation: 1 occupied electron shell

anion: 2 occupied electron shell

difference in size, bc cation have one less occupied electron shell

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

ionization energy (1st IE)

A

energy required to remove one mole of electron from one mole of gaseous atoms

X (g) → X+ (g) + e-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Is (1st) IE endothermic or exothermic

A

endothermic, bc IE is positive

bond breaking

bond between valence electron and nucleus is broken

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

ionization energy trends: in a group & explain

A

DECREASE DOWN a group

increase in the number of occupied electron shell

valence electron further away from nucleus

less attraction to the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

ionization energy trends: across period & explain

A

INCREASE ACROSS period

increasing number of protons (nuclear charge)

electrons are added to the same shell

SIMILAR SHIELDING effect from inner electons

stronger attraction

[shielding = repulsion]

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

why is there a drop in ionization energy from Be to B?

A

**GROUP 2 & 13

p orbital (2p?) is a higher orbital so electron is further from nucleus

weaker attraction

less energy needed to break the interaction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

why is there a drop in ionization energy from N to O?

A

**GROUP 15 & 16

N: electrons fill up p sub-levels and are unpaired

O: paired electrons in p sub-level

more electron-electron repulsion

easier to remove valence

lower IE

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

is electron affinity (1st) endothermic or exothermic?

A

exothermic

bond forming

when a metal forms an anion

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

electron affinity

A

energy change when one mole of electron is attached to one mole of natural atoms in gaseous state

X (g) + e- → X- (g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

electron affinity trends

A

DECREASES DOWN a group

INCREASES ACROSS a period

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

electron affinity trends: in a group & explain

A

DECREASES DOWN a group

increase in the number of occupied electron shell

valence electron further away from nucleus

less attraction to the nucleus

[most 1st EA values are negative]
become less negative down the group

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

electron affinity trends: across period & explain

A

INCREASE across period [values become more negative]

increasing umber of protons/increasing nuclear charge

electrons are added to the same shell

similar shielding effect from inner electrons

stronger attraction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

does Si or P have a more negative EA?

A

P (group 15): less negative than expected

[Ne] 3s2 3p3

electron gained will be paired up in 3p sub level, more electron-electron repulsion

Si
electron config

if add 1 electron to P, it becomes 3p4, have paired electrons = more repulsion

in Si – remain 3p3 singlely filled

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

does Li or Be have a more negative EA?

A

Be (group 2) less negative than expected

[He] 2s2

electron gained will be put into 2p sub level

2p sub level further away from nucleus

reduced attraction to nucleus with increased shielding

???

supposedly Be

(beryllium has a lower electron affinity compared to lithium)

Be: add into new sublevel?, further from nucleus

Li: 1s2 2s1
Be: 1s2 2s2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

is 1st EA positive or negative?

A

negative

electron is added to the outermost shell

the attraction between the nucleus and the electron leads to a decrease in energy of the system

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

is 2nd EA positive or negative?

A

positive

there is repulsive force between the anion and the negatively charged electron

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

electronegativity

A

ability of an atom to attract shared pairs of electrons (do not accept singular) in a covalent bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

what does high electronegativity mean?

A

pulls the electron pair towards it

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

electronegativity trends

A

DECREASES DOWN the group

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

electronegativity trends: in a group

A

DECREASES down the group

going down the group:

increasing atomic radii (down the group)

shared pair of electrons (must specify which electron) is further away from nucleus

increase shielding effect

less attraction for the shared pair of electrons

electronegativity decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

electronegativity trends: across period

A

INCREASES across period

going across the period:

increasing number of protons/increased nuclear charge

similar shielding effect from inner electrons

stronger attraction

electronegativity increases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

which of the following properties of the halogens increase from F to I?

a. atomic radius
b. melting point (& boiling point)
c. electronegativity

A

A and B (mp and bp increase)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q
  • state and explain the trends in the atomic radius and ionization energy

for the period 3 elements Na to Cl [4]

A

atomic radius decreases [1]

because nuclear charge increases AND electrons are added to the same main (outer) energy level [1]

ionization energy increases [1]

bc nuclear charge increases AND the electron removed is closer to the nucleus/is in the same energy level [1]

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

melting point trends: metal

A

DECREASE DOWN group

atomic/ionic radius increases

smaller charge density OR force of attraction between metal ions and delocalised electrons decreases

weaker metallic bonds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

melting point trends: non-metal

A

INCREASE DOWN group

molecular mass increase (van der waal force increase)

**except group 14 (carbon is giant covalent structure)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

metal trends: ionic radius across period

A

metal ionic radius decrease across period

metallic bond more strong

smaller ionic radius

more delocalised electron

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

non-metal trends: ionic radius across period

A

no obv trend

depends on structure & size of molecule

structure: some form polyatomic instead of diatomic

size: bigger more van der waal force

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

metallic/non metallic behaviour:

why does metal conduct electricity?

A

low ionization energy / more electronegativity

valence electron can become delocalised

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

metallic/non metallic behaviour:

why metal forms cation?

A

low ionization energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

metallic/non metallic behaviour:

why non-metals form anions?

A

negative electron affinity

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
38
Q

exothermic reaction

A

energy released in bond forming is more than energy needed in bond breaking

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

reactivity of alkali metals

what ions do they form when they react?

A

form 1+ ions

reactivity increases down the group

IE decrease down group

valence electrons further from nucleus

easier to remove valence electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

displacement of halides

A

a solution of more reactive halogen, X2, reacts with a solution of halide ions, X-, formed by a less reactive halogen

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
40
Q

how can the color from displacement of halides be differentiated?

A

by adding in organic solvent

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
41
Q

across period 3: bonding w oxygen

A

bonding: ionic to covalent across

difference in electronegativity decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
42
Q

metal oxides are either

A

basic or amphoteric

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
43
Q

non-metal oxides are either

A

acidic or neutral

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
44
Q

give than Na2O can react with H2SO4 (aq), what can be deduced?

A

it is basic

need to do 2 experiments: carry out another experiment to show it cannot react with a base

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
45
Q

across period 3: acidity of oxides

A

group 1 & 2: hydroxides considered strong

aq hydrogen involved with acidity

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
46
Q

is SiO2 soluble in water?

A

no, insoluble

acidic, reacts with alkali

SiO2 + 2NaOH –> Na2SiO3 + H2O

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
47
Q

is Al2O3 soluble in water?

A

no, insoluble

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
48
Q

how to test amphoteric substance?

A

react with acid and base + suggest reagent (SPECIFIC)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
49
Q

amphoteric

A

behaves as an acid or a base depending upon the reaction it is involved with

react with acids AND bases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
50
Q

is aluminium oxide amphoteric?

A

yes

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
51
Q

give the equation of the reaction between aluminium oxide and hydrochloric acid

A

Al2O3 (s) + 6HCl (aq) –> 2AlCl3 (aq) + 3H2O (l)

reacts with a strong acid to make a salt with water

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
52
Q

give the equation of the reaction between aluminium oxide and sodium hydroxide

A

Al2O3 (s) + 2NaOH (aq) + 3H2O (l) –> 2NaAl(OH)4 (aq)

reacts with a strong base to form sodium aluminate

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
53
Q

what is the equation that represents the first ionization energy of fluorine?

A

F (g) –> F+ (g) + e-

54
Q

which compound of an element in period 3 reacts with water to form a solution with a pH greater than 7?

A. SiO2
B. SiCl4
C. NaCl
D. Na2O

A

D

greater than 7 = basic

A. SiO2 = acidic
B. SiCl4 = acidic
C. NaCl = neutral
D. Na2O = BASIC!

55
Q

which oxides produce an acidic solution when added to water?

A. SiO2
B. P4O6
C. SO2

A

B & C
P4O6 & SO2

SiO2 is insoluble in water

56
Q

is this statement correct?

the melting points decrease from Li –> Cs for the alkali metals

A

YES

57
Q

is this statement correct?

the melting points increase from F –> I for the halogens

A

YES

58
Q

is this statement correct?

the melting points decrease from Na –> Ar for the period 3 elements

A

NO

it is hard to predict due to diff molecule size and shape

59
Q

acid deposition

A

acidic particles, gases and precipitation that fall to the earth

60
Q

wet deposition examples

A

acid rain, fog, snow

61
Q

2 types of acid deposition

A

wet & dry

62
Q

dry deposition examples

A

acidic gases and particles such as SO2 and salts (when acidic gas reacts with an alkali)

63
Q

effect of acid deposition on humans

A

irritation of mucus membranes and lung tissue when breathing in fine droplets of acid rain

increase risk of respiratory illness –> asthma, bronchitis

64
Q

effect of acid deposition on buildings

A

corrosion of materials such as marble and dolomite

CaCO3, MgCO3

faster corrosion of iron and steel structures in bridges (e.g. Fe)

64
Q

give the equation of the reaction between calcium carbonate and sulphuric acid (effect of wet deposition on buildings)

A

CaCO (s) + H2SO4(aq) –> CaSO4 (s) + H2O (l) + CO2 (g)

65
Q

effect of acid deposition on aquatic life

A

fish, algae, insect larvae even plankton cannot survive in water below a certain pH

66
Q

why is rainwater naturally acidic?

A

it contains dissolved CO2

67
Q

give the equations to do with rainwater

A

H2O (l) + CO2 (g) ⇌ H2CO3 (aq)

H2CO3 (aq) ⇌ H+ (aq) + HCO3- (aq)

68
Q

what is the pH of acid rain and what is the cause of this pH

A

caused by oxides of nitrogen and sulphur

69
Q

(sulphur oxides are a source of acid deposition) give the equations for the formation of atmospheric sulphuric acid and sulforous acid

A

SO3 (g) + H2O (l) –> H2SO4 (aq)

SO2(g) + H2O (l) –> H2O3 (aq)

70
Q

what is the cause for the creation of nitrogen oxides NOx

A

result of high temp in internal combustion engines (i.e. cars and jet engines)

71
Q

give the equation for the production of nitrogen oxides

A

N2 (g) + O2 (g) –> 2NO (g)

2NO (g) + O2 (g) –> 2NO2 (g)

72
Q

what is the color of NO2?

A

brown/orange

73
Q

what is the source of sulphur oxides

A

burning of coal which contains sulphur

smelting plants to extract metals from sulphur containing ore

74
Q

give the equation for the extraction of metal from sulphur containing ore

A

2ZnS + 3O2 –> 2ZnO + 2SO2

75
Q

give equations for the formation of sulphur oxides

A

S (g) + O2 (g) –> SO2 (g)

2SO2 (g) + O2 (g) –> 2SO3 (g)

76
Q

how to reduce acid deposition for sulphur dioxide

A

pre-combustion methods

(1) crush the coal and wash with water

metal sulphide with higher density sinks to bottom and separates from clean coal

(2) hydrodesulfurization

catalytic process

react petroleum with hydrogen

sulphur reacts with hydrogen to form hydrogen sulphide H2S

the toxic gas is captured and used to make H2SO4

77
Q

what are the post combustion methods for sulphur dioxide

A

(1) flue-gas desulfurization

acidic gas react with base –> neutralisation occurs

[CaCO3 (s) + SO2 (g) → CaSO3 (s) + CO2 (g)
2CaSO3 (s) + O2 (g) → 2CaSO4 (s)]

CaSO4 can be used to make plasterboard (building material)

wet slurry of CaO and CaCO3 reacts with SO2 to form CaSO4

CaO (s) + SO2 (g) –> CaSO3 (s)

CaSO3 (s) + SO2 (g) –> CaSO3 (s) + CO2 (g)

2CaSO3 (s) + O2 (g) –> 2CaSO4 (s)

CaSO4 can be used to make plasterboard, a building material

78
Q

post combustion methods for nitrogen oxides NOx

A

(1) catalytic converter –> a redox reaction

use Pt and Pd as catalyst

turn toxic gas into harmless product

2NO (g) + 2CO (g) –> 2CO2 (g) + N2 (g)

(2) low temperature combustion = lower operating temperature

recirculate exhaust gas back into engine

reduce formation of NOx

*cannot combust fully bc of lowered temp

79
Q

oxidation

A

oxidation = oxidation number increase, loss of electrons

e on RHS

Fe2+ –> Fe3+ + e-

+2 to +3

79
Q

reduction

A

reduction = oxidation number decrease, gain of electrons

e on LHS

2H+ +2e- –> H2

+1 to 0

80
Q

reducing agent

A

reduces another substance during redox reaction

itself oxidised

80
Q

oxidising agent

A

oxidises another substitute during redox reaction

itself reduced

80
Q

[FeCl2(NH3)4]Cl what is the oxidation number for iron?

A

+3

80
Q

rules for oxidation number

A

1) number for elements = 0

2) same as charge of monoatomic ion

3) main group metal ions follow group number

4) transition metals (variable oxidation state)

charge depends on bonded non-metals

roman numeral: iron lll oxide

81
Q

what is the oxidation number of hydrogen ion and the exception

A

+1

except when bonded with metal e.g. NaH

oxidation number = -1, due to hydrogen having higher electronegativity

82
Q

[FeCl2(NH3)4] what is the oxidation number for iron?

A

+2

83
Q

what does the oxidation number of non-metals depend on? give an example to exception due to this reason

A

electronegativity

example: F2O
F: -1
O: +2

it is +2 instead of -2 because the electronegativity of F > O

others are variable??

___

except H2O2: -1? hydrogen peroxide

84
Q

2MnO4- + 5SO2 + 2H2O –> 2Mn2+ + 5SO42- + 4H+

which element is oxidised and which is reduced?

A

oxidised: sulphur, b/c increase in oxidation number

reduced; Mn, b/c decrease in oxidation number

85
Q

what is Cr’s electron configuration? (chromium)

A

1s2 2s2 2p6 3s2 3p6 4s1 3d5

85
Q

what is Cu’s electron configuration? (chromium)

A

1s2 2s2 2p6 3s2 3p6 4s1 3d10

86
Q

what is special about the electron configuration of Cu and Cr?

A

4s1

86
Q

when transition metals form ions, electron are removed from what first?

A

4s BEFORE 3d

87
Q

transition metal definition (2016)

A

form ions (2+) wih incomplete d-orbital

88
Q

which d-block element is not a transition metal?

A

Zn zinc (3d10 bc its 4s2)

Ca? (no d orbital, d0)

89
Q

why is Zn not a transition metal?

A

do not formed coloured ions

ions do not have partially filled d orbitals

90
Q

main properties of transition metals

A

harder (FYI)
higher melting point
higher density (FYI)
variable oxidation state
colored complexes with ligands
have magnetic and catalytic properties

91
Q

successive IE of aluminium and the relationship between the oxidation state

A

large difference in energy between 3rd and 4th ionization

oxidations state +3

91
Q

explain maximum oxidation state

A

increases in steps of +1 up till Mn

then decreases in steps of +1

nuclear charge increases

ease of losing 3d electron decrease

M3+ is stable for Sc –> Cr

M2+ is stable for Mn –> Cu

91
Q

chemical properties of transition metal

A

variable oxidation number
form complex ions with ligands
act as catalyst
have magnetic property
colored compound

92
Q

successive IE of transition metal

A

energy level between 3d and 4s electrons are close

no sudden increase in successive IE (until all 4s and 3d electron are lost)

variable oxidation states

93
Q

why do transition metals have higher melting point than group 1 and 2 metals?

A

shorter ionic radius = stronger attraction

1) distance
2) charge

93
Q

metallic bonding

A

electrostatic force of attraction between lattice of metal ions and the delocalised sea of electrons

what does this mean?

1) smaller ionic radius
2) more valence electron

93
Q

define catalyst

A

increase rate of reaction by providing an alternative reaction pathway with a lower activation energy

93
Q

homogenous catalysis

A

catalyst same state/phase as reactant

94
Q

heterogenous catalysis

A

catalyst in diff state/phase as reactant

94
Q

why do transition metals show catalytic activity?

A

due to partly filled d-orbitals which can be used to form bonds with adsorbed reactants which helps reactions take place more easily

94
Q

examples of transition metals that are catalysts

A

iron: haber process (manufacture of ammonia)

nickel: hydrogenation reactions (margarine manufacture)

rhodium: catalytic converters

vanadium (V) oxide: contact process (manufacture of sulphuric acid)

94
Q

equation for contact process

A

Contact Process:
Production of H2SO4
V2O5: 2SO2 (g) + O2 (g) 2SO3 (g)

94
Q

equation for haber process involving its catalyst

A

Haber Process:
Production of ammonium
Fe: N2(g) + 3H2 (g) 2NH3(g)

94
Q

equation for catalytic converter

A

Pd & Pt: 2CO (g) + 2NO (g)2CO2 (g) + N2 (g)

94
Q

equation for decomposition of hydrogen peroxide

A

MnO2: 2H2O2 (aq)2H2O (l) + O2 (g)

95
Q

illustrate an example of homogenous catalyst

A

Reaction between acidified hydrogen peroxide and iodide

H2O2 + 2H+ + 2Fe2+ –> 2H2O + 2Fe3+

2I- + 2Fe3+ –> I2 + 2Fe2+

Fe2+ is regenerated

Due to presence of variable oxidation states

96
Q

give and explain example of homogenous catalyst

A

Iron (Fe2+) in heme
- Ligand: porphyrin
- oxygen binds to Fe2+

Cobalt (Co3+) in vitamin B12
- Form a octehedral complex
- Ligand: porphyrin
- One site available for biological activities

97
Q

each spinning electron acts as…?

A

magnetic

98
Q

ferromagnetic and give examples

A

contain unpaired electron (all unpaired)

e.g. iron, cobalt and nickel

strongly attracted by external magnetic field, strongest magnetic property

domains remain align even when external magnetic field is removed – will remain in same position

98
Q

paramagnetic

A

contain unpaired electrons

slightly attracted by external magnetic field

do not retain magnetic property when external magnetic field removed

will randomise with removed external field

98
Q

diamagnetic and give an example

A

contain only paired electrons (all paired up)

not attracted, even weakly repelled by external magnetic field

example: zinc

99
Q

what type of magnetic is oxygen?

A

paramagnetic

100
Q

what type of magnetic is nitrogen?

A

diamagnetic

101
Q

formation of complex ions is due to?

A
  • high charge density
  • small atomic size
  • unfilled orbitals to accept lone electron pairs
102
Q

ligand

A

with lone pair of electrons and form coordinate bond with metal ions

both electron used to form the bond come from the same atom

103
Q

monodentate

A

form 1 coordinate bond

104
Q

polydentate

A

form 2 or more coordinate bonds

105
Q

number of coordinate bond = ?

A

coordinate number

106
Q

coordination number, give the electron domain geometry

A

3: trigonal planar
4: tetrahedral
5: trigonal bipyramidal
6: octahedral

107
Q

which electrons are lost by an atom of iron when it forms the Fe3+ ion?

A

two s orbital electrons and one d orbital electron

108
Q

which properties are typical of d-block elements

A. complex ion formation
B. catalytic behaviour
C. colorless compounds

A

A and B

109
Q

Outline why transition metals form coloured compounds [4]

A

Partially filled d-orbitals

(Ligands cause) d-orbitals (to) split

Light is absorbed as electrons transit to a higher level (in d-d transitions)
OR
Light is absorbed as electrons are promoted

Energy gap corresponds to light in the visible region of the spectrum

Colour observed is the complementary colour

110
Q

(NH4)5[CuCl6]

for the Cu ion, what magnetic property does the Cu complex have?

A

diamagnetic

111
Q

what color is Fe3+

A

yellow

112
Q

what color is Fe2+

A

green

113
Q

what happens if you put ligand put close to electron

A

repulsion (b/c both negative)

make orbital energetically less stable

114
Q

what is electron transition

A

in lower level become excited, go to higher level

show the complementary color

light energy match visible light

1 color light absorbed, 1 color pass through, the (opposite) color shows –> complementary

115
Q

color of complexes (describe d-d splitting)

A

light passes through solution of [Cu(H2O)6]2+

one electron in 3d orbital is excited from lower to higher energy sub-level

a photon of yellow light is absorbed

116
Q

what gives the color of transition metals?

A

energy gap between the d orbitals after d-d splitting

the bigger the gap, absorb higher energy

  • higher frequency
  • shorter wavelength
117
Q
A
118
Q
A

chem hl (2 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions