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periodicity Flashcards

1
Q

atomic radius trends

A

INCREASE DOWN the group

DECREASE ACROSS a period

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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]

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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

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4
Q

ionic radius: cations

predict whether Na or Na+ is larger in size

A

Na

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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)

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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

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5
Q

ionic radius trends

A

INCREASE DOWN the group

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6
Q

ionic radius: cations

predict whether S or S2- is larger in size

A

S2-

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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

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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

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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

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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

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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-

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12
Q

Is (1st) IE endothermic or exothermic

A

endothermic, bc IE is positive

bond breaking

bond between valence electron and nucleus is broken

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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

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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]

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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

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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

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17
Q

is electron affinity (1st) endothermic or exothermic?

A

exothermic

bond forming

when a metal forms an anion

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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)

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18
Q

electron affinity trends

A

DECREASES DOWN a group

INCREASES ACROSS a period

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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

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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

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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

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22
does Li or Be have a more negative EA?
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
23
is 1st EA positive or negative?
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
24
is 2nd EA positive or negative?
positive there is repulsive force between the anion and the negatively charged electron
25
electronegativity
ability of an atom to attract shared pairs of electrons (do not accept singular) in a covalent bond
26
what does high electronegativity mean?
pulls the electron pair towards it
27
electronegativity trends
DECREASES DOWN the group
28
electronegativity trends: in a group
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
29
electronegativity trends: across period
INCREASES across period going across the period: increasing number of protons/increased nuclear charge similar shielding effect from inner electrons stronger attraction electronegativity increases
30
which of the following properties of the halogens increase from F to I? a. atomic radius b. melting point (& boiling point) c. electronegativity
A and B (mp and bp increase)
31
* state and explain the trends in the atomic radius and ionization energy for the period 3 elements Na to Cl [4]
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]
32
melting point trends: metal
DECREASE DOWN group atomic/ionic radius increases smaller charge density OR force of attraction between metal ions and delocalised electrons decreases weaker metallic bonds
33
melting point trends: non-metal
INCREASE DOWN group molecular mass increase (van der waal force increase) **except group 14 (carbon is giant covalent structure)
34
metal trends: ionic radius across period
metal ionic radius decrease across period metallic bond more strong smaller ionic radius more delocalised electron
35
non-metal trends: ionic radius across period
no obv trend depends on structure & size of molecule structure: some form polyatomic instead of diatomic size: bigger more van der waal force
36
metallic/non metallic behaviour: why does metal conduct electricity?
low ionization energy / more electronegativity valence electron can become delocalised
37
metallic/non metallic behaviour: why metal forms cation?
low ionization energy
37
metallic/non metallic behaviour: why non-metals form anions?
negative electron affinity
38
exothermic reaction
energy released in bond forming is more than energy needed in bond breaking
39
reactivity of alkali metals what ions do they form when they react?
form 1+ ions reactivity increases down the group IE decrease down group valence electrons further from nucleus easier to remove valence electrons
39
displacement of halides
a solution of more reactive halogen, X2, reacts with a solution of halide ions, X-, formed by a less reactive halogen
40
how can the color from displacement of halides be differentiated?
by adding in organic solvent
41
across period 3: bonding w oxygen
bonding: ionic to covalent across difference in electronegativity decreases
42
metal oxides are either
basic or amphoteric
43
non-metal oxides are either
acidic or neutral
44
give than Na2O can react with H2SO4 (aq), what can be deduced?
it is basic need to do 2 experiments: carry out another experiment to show it cannot react with a base
45
across period 3: acidity of oxides
group 1 & 2: hydroxides considered strong aq hydrogen involved with acidity
46
is SiO2 soluble in water?
no, insoluble acidic, reacts with alkali SiO2 + 2NaOH --> Na2SiO3 + H2O
47
is Al2O3 soluble in water?
no, insoluble
48
how to test amphoteric substance?
react with acid and base + suggest reagent (SPECIFIC)
49
amphoteric
behaves as an acid or a base depending upon the reaction it is involved with react with acids AND bases
50
is aluminium oxide amphoteric?
yes
51
give the equation of the reaction between aluminium oxide and hydrochloric acid
Al2O3 (s) + 6HCl (aq) --> 2AlCl3 (aq) + 3H2O (l) reacts with a strong acid to make a salt with water
52
give the equation of the reaction between aluminium oxide and sodium hydroxide
Al2O3 (s) + 2NaOH (aq) + 3H2O (l) --> 2NaAl(OH)4 (aq) reacts with a strong base to form sodium aluminate
53
what is the equation that represents the first ionization energy of fluorine?
F (g) --> F+ (g) + e-
54
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
D greater than 7 = basic A. SiO2 = acidic B. SiCl4 = acidic C. NaCl = neutral D. Na2O = BASIC!
55
which oxides produce an acidic solution when added to water? A. SiO2 B. P4O6 C. SO2
B & C P4O6 & SO2 SiO2 is insoluble in water
56
is this statement correct? the melting points decrease from Li --> Cs for the alkali metals
YES
57
is this statement correct? the melting points increase from F --> I for the halogens
YES
58
is this statement correct? the melting points decrease from Na --> Ar for the period 3 elements
NO it is hard to predict due to diff molecule size and shape
59
acid deposition
acidic particles, gases and precipitation that fall to the earth
60
wet deposition examples
acid rain, fog, snow
61
2 types of acid deposition
wet & dry
62
dry deposition examples
acidic gases and particles such as SO2 and salts (when acidic gas reacts with an alkali)
63
effect of acid deposition on humans
irritation of mucus membranes and lung tissue when breathing in fine droplets of acid rain increase risk of respiratory illness --> asthma, bronchitis
64
effect of acid deposition on buildings
corrosion of materials such as marble and dolomite CaCO3, MgCO3 faster corrosion of iron and steel structures in bridges (e.g. Fe)
64
give the equation of the reaction between calcium carbonate and sulphuric acid (effect of wet deposition on buildings)
CaCO (s) + H2SO4(aq) --> CaSO4 (s) + H2O (l) + CO2 (g)
65
effect of acid deposition on aquatic life
fish, algae, insect larvae even plankton cannot survive in water below a certain pH
66
why is rainwater naturally acidic?
it contains dissolved CO2
67
give the equations to do with rainwater
H2O (l) + CO2 (g) ⇌ H2CO3 (aq) H2CO3 (aq) ⇌ H+ (aq) + HCO3- (aq)
68
what is the pH of acid rain and what is the cause of this pH
caused by oxides of nitrogen and sulphur
69
(sulphur oxides are a source of acid deposition) give the equations for the formation of atmospheric sulphuric acid and sulforous acid
SO3 (g) + H2O (l) --> H2SO4 (aq) SO2(g) + H2O (l) --> H2O3 (aq)
70
what is the cause for the creation of nitrogen oxides NOx
result of high temp in internal combustion engines (i.e. cars and jet engines)
71
give the equation for the production of nitrogen oxides
N2 (g) + O2 (g) --> 2NO (g) 2NO (g) + O2 (g) --> 2NO2 (g)
72
what is the color of NO2?
brown/orange
73
what is the source of sulphur oxides
burning of coal which contains sulphur smelting plants to extract metals from sulphur containing ore
74
give the equation for the extraction of metal from sulphur containing ore
2ZnS + 3O2 --> 2ZnO + 2SO2
75
give equations for the formation of sulphur oxides
S (g) + O2 (g) --> SO2 (g) 2SO2 (g) + O2 (g) --> 2SO3 (g)
76
how to reduce acid deposition for sulphur dioxide
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
what are the post combustion methods for sulphur dioxide
(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
post combustion methods for nitrogen oxides NOx
(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
oxidation
oxidation = oxidation number increase, loss of electrons e on RHS Fe2+ --> Fe3+ + e- +2 to +3
79
reduction
reduction = oxidation number decrease, gain of electrons e on LHS 2H+ +2e- --> H2 +1 to 0
80
reducing agent
reduces another substance during redox reaction itself oxidised
80
oxidising agent
oxidises another substitute during redox reaction itself reduced
80
[FeCl2(NH3)4]Cl what is the oxidation number for iron?
+3
80
rules for oxidation number
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
what is the oxidation number of hydrogen ion and the exception
+1 except when bonded with metal e.g. NaH oxidation number = -1, due to hydrogen having higher electronegativity
82
[FeCl2(NH3)4] what is the oxidation number for iron?
+2
83
what does the oxidation number of non-metals depend on? give an example to exception due to this reason
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
2MnO4- + 5SO2 + 2H2O --> 2Mn2+ + 5SO42- + 4H+ which element is oxidised and which is reduced?
oxidised: sulphur, b/c increase in oxidation number reduced; Mn, b/c decrease in oxidation number
85
what is Cr's electron configuration? (chromium)
1s2 2s2 2p6 3s2 3p6 4s1 3d5
85
what is Cu's electron configuration? (chromium)
1s2 2s2 2p6 3s2 3p6 4s1 3d10
86
what is special about the electron configuration of Cu and Cr?
4s1
86
when transition metals form ions, electron are removed from what first?
4s BEFORE 3d
87
transition metal definition (2016)
form ions (2+) wih incomplete d-orbital
88
which d-block element is not a transition metal?
Zn zinc (3d10 bc its 4s2) Ca? (no d orbital, d0)
89
why is Zn not a transition metal?
do not formed coloured ions ions do not have partially filled d orbitals
90
main properties of transition metals
harder (FYI) higher melting point higher density (FYI) variable oxidation state colored complexes with ligands have magnetic and catalytic properties
91
successive IE of aluminium and the relationship between the oxidation state
large difference in energy between 3rd and 4th ionization oxidations state +3
91
explain maximum oxidation state
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
chemical properties of transition metal
variable oxidation number form complex ions with ligands act as catalyst have magnetic property colored compound
92
successive IE of transition metal
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
why do transition metals have higher melting point than group 1 and 2 metals?
shorter ionic radius = stronger attraction 1) distance 2) charge
93
metallic bonding
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
define catalyst
increase rate of reaction by providing an alternative reaction pathway with a lower activation energy
93
homogenous catalysis
catalyst same state/phase as reactant
94
heterogenous catalysis
catalyst in diff state/phase as reactant
94
why do transition metals show catalytic activity?
due to partly filled d-orbitals which can be used to form bonds with adsorbed reactants which helps reactions take place more easily
94
examples of transition metals that are catalysts
iron: haber process (manufacture of ammonia) nickel: hydrogenation reactions (margarine manufacture) rhodium: catalytic converters vanadium (V) oxide: contact process (manufacture of sulphuric acid)
94
equation for contact process
Contact Process: Production of H2SO4 V2O5: 2SO2 (g) + O2 (g) 2SO3 (g)
94
equation for haber process involving its catalyst
Haber Process: Production of ammonium Fe: N2(g) + 3H2 (g) 2NH3(g)
94
equation for catalytic converter
Pd & Pt: 2CO (g) + 2NO (g)2CO2 (g) + N2 (g)
94
equation for decomposition of hydrogen peroxide
MnO2: 2H2O2 (aq)2H2O (l) + O2 (g)
95
illustrate an example of homogenous catalyst
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
give and explain example of homogenous catalyst
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
each spinning electron acts as...?
magnetic
98
ferromagnetic and give examples
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
paramagnetic
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
diamagnetic and give an example
contain only paired electrons (all paired up) not attracted, even weakly repelled by external magnetic field example: zinc
99
what type of magnetic is oxygen?
paramagnetic
100
what type of magnetic is nitrogen?
diamagnetic
101
formation of complex ions is due to?
- high charge density - small atomic size - unfilled orbitals to accept lone electron pairs
102
ligand
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
monodentate
form 1 coordinate bond
104
polydentate
form 2 or more coordinate bonds
105
number of coordinate bond = ?
coordinate number
106
coordination number, give the electron domain geometry
3: trigonal planar 4: tetrahedral 5: trigonal bipyramidal 6: octahedral
107
which electrons are lost by an atom of iron when it forms the Fe3+ ion?
two s orbital electrons and one d orbital electron
108
which properties are typical of d-block elements A. complex ion formation B. catalytic behaviour C. colorless compounds
A and B
109
Outline why transition metals form coloured compounds [4]
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
(NH4)5[CuCl6] for the Cu ion, what magnetic property does the Cu complex have?
diamagnetic
111
what color is Fe3+
yellow
112
what color is Fe2+
green
113
what happens if you put ligand put close to electron
repulsion (b/c both negative) make orbital energetically less stable
114
what is electron transition
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
color of complexes (describe d-d splitting)
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
what gives the color of transition metals?
energy gap between the d orbitals after d-d splitting the bigger the gap, absorb higher energy - higher frequency - shorter wavelength
117
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chem hl (2 decks)

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