periodicity Flashcards
atomic radius trends
INCREASE DOWN the group
DECREASE ACROSS a period
atomic radius trend in a group & explain
radius INCREASE down the group
electrons occupied another energy level
valence electrons further from nucleus
[make specific]
atomic radius trend across the period & explain
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
ionic radius: cations
predict whether Na or Na+ is larger in size
Na
what is effective nuclear charge
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)
ionic radius trend in group & explain (both anion and cation)
INCREASE down the group
(has more occupied electron shells)
valence electrons are further from nucleus
ionic radius trends
INCREASE DOWN the group
ionic radius: cations
predict whether S or S2- is larger in size
S2-
ionic radius trends: size of cations & reason
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
ionic radius trends: size of anions & reason
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
ionic radius trends: across the period
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
is cation smaller or larger than anions & why?
SMALLLER
cation: 1 occupied electron shell
anion: 2 occupied electron shell
difference in size, bc cation have one less occupied electron shell
ionization energy (1st IE)
energy required to remove one mole of electron from one mole of gaseous atoms
X (g) → X+ (g) + e-
Is (1st) IE endothermic or exothermic
endothermic, bc IE is positive
bond breaking
bond between valence electron and nucleus is broken
ionization energy trends: in a group & explain
DECREASE DOWN a group
increase in the number of occupied electron shell
valence electron further away from nucleus
less attraction to the nucleus
ionization energy trends: across period & explain
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]
why is there a drop in ionization energy from Be to B?
**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
why is there a drop in ionization energy from N to O?
**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
is electron affinity (1st) endothermic or exothermic?
exothermic
bond forming
when a metal forms an anion
electron affinity
energy change when one mole of electron is attached to one mole of natural atoms in gaseous state
X (g) + e- → X- (g)
electron affinity trends
DECREASES DOWN a group
INCREASES ACROSS a period
electron affinity trends: in a group & explain
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
electron affinity trends: across period & explain
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
does Si or P have a more negative EA?
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
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
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
is 2nd EA positive or negative?
positive
there is repulsive force between the anion and the negatively charged electron
electronegativity
ability of an atom to attract shared pairs of electrons (do not accept singular) in a covalent bond
what does high electronegativity mean?
pulls the electron pair towards it
electronegativity trends
DECREASES DOWN the group
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
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
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)
- 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]
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
melting point trends: non-metal
INCREASE DOWN group
molecular mass increase (van der waal force increase)
**except group 14 (carbon is giant covalent structure)
metal trends: ionic radius across period
metal ionic radius decrease across period
metallic bond more strong
smaller ionic radius
more delocalised electron
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
metallic/non metallic behaviour:
why does metal conduct electricity?
low ionization energy / more electronegativity
valence electron can become delocalised
metallic/non metallic behaviour:
why metal forms cation?
low ionization energy
metallic/non metallic behaviour:
why non-metals form anions?
negative electron affinity
exothermic reaction
energy released in bond forming is more than energy needed in bond breaking
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
displacement of halides
a solution of more reactive halogen, X2, reacts with a solution of halide ions, X-, formed by a less reactive halogen
how can the color from displacement of halides be differentiated?
by adding in organic solvent
across period 3: bonding w oxygen
bonding: ionic to covalent across
difference in electronegativity decreases
metal oxides are either
basic or amphoteric
non-metal oxides are either
acidic or neutral
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
across period 3: acidity of oxides
group 1 & 2: hydroxides considered strong
aq hydrogen involved with acidity
is SiO2 soluble in water?
no, insoluble
acidic, reacts with alkali
SiO2 + 2NaOH –> Na2SiO3 + H2O
is Al2O3 soluble in water?
no, insoluble
how to test amphoteric substance?
react with acid and base + suggest reagent (SPECIFIC)
amphoteric
behaves as an acid or a base depending upon the reaction it is involved with
react with acids AND bases
is aluminium oxide amphoteric?
yes
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
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
what is the equation that represents the first ionization energy of fluorine?
F (g) –> F+ (g) + e-
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!
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
is this statement correct?
the melting points decrease from Li –> Cs for the alkali metals
YES
is this statement correct?
the melting points increase from F –> I for the halogens
YES
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
acid deposition
acidic particles, gases and precipitation that fall to the earth
wet deposition examples
acid rain, fog, snow
2 types of acid deposition
wet & dry
dry deposition examples
acidic gases and particles such as SO2 and salts (when acidic gas reacts with an alkali)
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
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)
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)
effect of acid deposition on aquatic life
fish, algae, insect larvae even plankton cannot survive in water below a certain pH
why is rainwater naturally acidic?
it contains dissolved CO2
give the equations to do with rainwater
H2O (l) + CO2 (g) ⇌ H2CO3 (aq)
H2CO3 (aq) ⇌ H+ (aq) + HCO3- (aq)
what is the pH of acid rain and what is the cause of this pH
caused by oxides of nitrogen and sulphur
(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)
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)
give the equation for the production of nitrogen oxides
N2 (g) + O2 (g) –> 2NO (g)
2NO (g) + O2 (g) –> 2NO2 (g)
what is the color of NO2?
brown/orange
what is the source of sulphur oxides
burning of coal which contains sulphur
smelting plants to extract metals from sulphur containing ore
give the equation for the extraction of metal from sulphur containing ore
2ZnS + 3O2 –> 2ZnO + 2SO2
give equations for the formation of sulphur oxides
S (g) + O2 (g) –> SO2 (g)
2SO2 (g) + O2 (g) –> 2SO3 (g)
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
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
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
oxidation
oxidation = oxidation number increase, loss of electrons
e on RHS
Fe2+ –> Fe3+ + e-
+2 to +3
reduction
reduction = oxidation number decrease, gain of electrons
e on LHS
2H+ +2e- –> H2
+1 to 0
reducing agent
reduces another substance during redox reaction
itself oxidised
oxidising agent
oxidises another substitute during redox reaction
itself reduced
[FeCl2(NH3)4]Cl what is the oxidation number for iron?
+3
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
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
[FeCl2(NH3)4] what is the oxidation number for iron?
+2
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
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
what is Cr’s electron configuration? (chromium)
1s2 2s2 2p6 3s2 3p6 4s1 3d5
what is Cu’s electron configuration? (chromium)
1s2 2s2 2p6 3s2 3p6 4s1 3d10
what is special about the electron configuration of Cu and Cr?
4s1
when transition metals form ions, electron are removed from what first?
4s BEFORE 3d
transition metal definition (2016)
form ions (2+) wih incomplete d-orbital
which d-block element is not a transition metal?
Zn zinc (3d10 bc its 4s2)
Ca? (no d orbital, d0)
why is Zn not a transition metal?
do not formed coloured ions
ions do not have partially filled d orbitals
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
successive IE of aluminium and the relationship between the oxidation state
large difference in energy between 3rd and 4th ionization
oxidations state +3
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
chemical properties of transition metal
variable oxidation number
form complex ions with ligands
act as catalyst
have magnetic property
colored compound
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
why do transition metals have higher melting point than group 1 and 2 metals?
shorter ionic radius = stronger attraction
1) distance
2) charge
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
define catalyst
increase rate of reaction by providing an alternative reaction pathway with a lower activation energy
homogenous catalysis
catalyst same state/phase as reactant
heterogenous catalysis
catalyst in diff state/phase as reactant
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
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)
equation for contact process
Contact Process:
Production of H2SO4
V2O5: 2SO2 (g) + O2 (g) 2SO3 (g)
equation for haber process involving its catalyst
Haber Process:
Production of ammonium
Fe: N2(g) + 3H2 (g) 2NH3(g)
equation for catalytic converter
Pd & Pt: 2CO (g) + 2NO (g)2CO2 (g) + N2 (g)
equation for decomposition of hydrogen peroxide
MnO2: 2H2O2 (aq)2H2O (l) + O2 (g)
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
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
each spinning electron acts as…?
magnetic
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
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
diamagnetic and give an example
contain only paired electrons (all paired up)
not attracted, even weakly repelled by external magnetic field
example: zinc
what type of magnetic is oxygen?
paramagnetic
what type of magnetic is nitrogen?
diamagnetic
formation of complex ions is due to?
- high charge density
- small atomic size
- unfilled orbitals to accept lone electron pairs
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
monodentate
form 1 coordinate bond
polydentate
form 2 or more coordinate bonds
number of coordinate bond = ?
coordinate number
coordination number, give the electron domain geometry
3: trigonal planar
4: tetrahedral
5: trigonal bipyramidal
6: octahedral
which electrons are lost by an atom of iron when it forms the Fe3+ ion?
two s orbital electrons and one d orbital electron
which properties are typical of d-block elements
A. complex ion formation
B. catalytic behaviour
C. colorless compounds
A and B
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
(NH4)5[CuCl6]
for the Cu ion, what magnetic property does the Cu complex have?
diamagnetic
what color is Fe3+
yellow
what color is Fe2+
green
what happens if you put ligand put close to electron
repulsion (b/c both negative)
make orbital energetically less stable
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
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
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
