Download 03 Periodicity

03 Periodicity
3.1 THE PERIODIC TABLE
The Periodic Table
• Groups - Vertical columns from I to VIII;
number of outershell/valence electrons; down
a group – gradual change in physical
properties and similar chem. properties
• Period – Horizontal rows from 1 to 7; number
of occupied (electron) shells;
Definitions: Periodicity and The
Periodic Table
• The Periodic Table: Arrangements of elements
in the Periodic Table is in order of increasing
atomic number
• Periodicity: repeating pattern of (physical and
chemical) properties across a period
3.2 PHYSICAL PROPERTIES
Effective Nuclear Charge
• Descending a group: remains the same increase in nuclear charge is offset by the
increase in shielding (# of inner electrons)
• Across a period: increases - the nuclear charge
increases and there is no increase in shielding
(or increase in # of inner electrons as
electrons are added to the same energy shell)
Atomic Radius
Atomic Radius: Measured as half the distance
between two bonded atoms (hence noble gas
are given no value as they do not bond with
each other atoms)
Atomic Radius Trends
• Down a group: increases - # of occupied
electron shells (given by period #) increases
• Across a period: decreases – increased
effective nuclear charge, increased forces of
attraction draws the electrons closer to the
Ionic Radius
Ionic Radius: Distance between nucleus and
outer most electrons of a positive metal cation
or a negative non-metal anion
• Metal cations tend to be smaller than their
atom – loss of outermost electrons and shell
• Non-metal anions tend to be larger than their
atom - gained electrons into their outer
energy levels, increased electron repulsion
between the electrons
Comparing Ionic Radius
• Na: 11 p, 11 (2.8.1 – 3 filled energy levels) and
Na+: 11 p, 10 (2.8 – 2 filled energy levels)
Na+ has greater net positive charge (same
number of protons
pulling smaller number of electrons);
• Si4+: 10 e– (2.8) in 2 filled energy levels;
P3–: 18 e– (2.8.8) in 3 filled energy levels
therefore larger
Ionic Radius Trends
• Across a Group: Increases – increased # of
occupied shells
• Across a Period: cations: decrease - increased
nuclear charge; anions: decrease – increased
nuclear charge
First Ionisation Energy - RECAP
• Define: energy (in kJ mol⁻¹ per mole) required
to remove one mole of electrons from one
mole of gaseous atoms
𝐻(𝑔) → 𝐻(𝑔) + + 𝑒 −
REMEMBER STATE SYMBOLS AND + AND -
First Ionisation Energy Trends
• Down a Group: decrease – atomic radius and
nuclear charge increases, yet level of shielding
also increases and hence the effective nuclear
charge decreases
• Across a Period: generally increases – as there
is an increase in effective nuclear charge as
the number of protons increase with no
increase in shielding
Electronegativity Trends
• Define: ability of an atom to attract bonding electrons in a
covalent bond ^ derived indirectly from experimental bond
energy data
• Noble Gases do not have a value as it doesn’t form bonds
easily and cannot attract more electrons to its full energy
level
• Down a Group: decreases – atomic radius and nuclear
charge increases, yet level of shielding also increases and
hence the effective nuclear charge decreases
• Across a Period: increases – increased effective nuclear
charge w/ similar shielding and decreased atomic radius
Background Knowledge
• Valance electrons are delocalised and are
attracted to the now positively charged ions
• As we go across a period there are more
delocalised electrons and more positively
charged ions hence a higher charge to mass
ratio making the bonds stronger and hence
increasing mp
Melting Point Group Trends
• Down a metal group (1): decreases - metallic
bonding gets weaker due to smaller charge to
volume ratio (charge density) as atomic radii
gets bigger and delocalised valence electrons
are shielded by inner electrons
• Down a non-metal group (7): increases –
stronger VDW forces due to increase in # of
electrons and higher molar mass of halogen
molecules
Melting Point Period Trends
• Na, Mg, Al: steady increase across metals – stronger
metallic bonding due to greater charge density (larger
charge, smaller atomic radii) leading to increase in
attraction of delocalised electrons
• Si: massive increase Si - giant macromolecular
structure; requires a lot of energy to break;
• P: large decrease - 𝑃4 molecules
• S: small increase - 𝑆8 crown shaped molecules (higher
Mr (larger molecule or greater # of electrons), greater
VDW forces )
• Cl and Ar: small decrease - Cl2 molecules and simple
monoatomic Ar atoms (only weak VDW)
3.3 CHEMICAL PROPERTIES
Group 0: Noble Gases
• Colorless gases
• Monoatomic – exist as single atoms
• Very unreactive – inability of their atoms to
lose or gain electrons; don’t form cations as
they have the highest ionisation energies;
don’t form anions as extra electrons would
have to be added to an empty outer shell
• Stable octet (He – stable with 2 electrons)
Group 1: The Alkali Metals
Physical Properties
• Grey shiny surfaces yet rapidly reacts with oxygen in the air, giving it
a dark oxide coat
• soft, can be easily cut with a knife
• low densities
• conduct electricity and heat – delocalised mobile electrons
Chemical Properties
• Very reactive – increases in reactivity down the group as electron
lost is further from nucleus so it is less tightly held; Rb electron is in
5th energy level and (Na less reactive) as electron lost in 3rd energy
level
• Form ionic compounds with non-metals
Alkali Metals and their violent
reactions with water
•
•
•
•
float and move across the surface of the water
Fizz, effervescences, bubbles - produce gas
Metal Decreases in Size
Li releases Hydrogen but heat produced is not
sufficient to melt the unreacted metal
• Na sometimes catches fire (orange/ yellow flame)
• K catches fire vigorously (lilac flame)
the solution has blue colour when universal
indicator is added = alkali (NOT AN OBSERVATION!)
Equations: Alkali Metals and their
reactions with water
• 2Li (s) + 2H2O (g)  2LiOH (aq) + H2 (g)
• 2Na (s) + 2H2O (g)  2NaOH (aq) + H2 (g)
• 2K (s) + 2H2O (g)  2KOH (aq) + H2 (g)
2𝐾 𝑠 + 𝐻2 𝑂 𝑙 → 2𝐾 + 𝑎𝑞 + 2𝑂𝐻 − 𝑎𝑞 + 𝐻2 (𝑔)
KOH is an ionic compound and dissociates in water
Group 7: Halogens
• Diatomic molecules 𝑋2
• Oxidising agents – react by gaining an electron
to form a negative anion
Chemical Properties:
• Reactivity decreases – as atomic radii
increases and attraction for electrons
decreases
Halogens Physical Properties
• Colored elements
• Gradual changes in color and state
• Density increases
F2
Cl2
Br2
I2
Colour
Yellow
Green
Red/brown
Grey
State (at RTP)
GAS
GAS
LIQUID
SOLID
Reaction with Group 1 metals
• Forms Ionic Halides – redox
• Most vigorous reactions occur between
elements with large EN∆
• Transfer of electrons to form electrostatic
forces of attraction between two ions
2𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 2𝑁𝑎𝐶𝑙 (𝑠)
Needed? Displacement Reaction
• Gas is bubbled into solution of Potassium Halide –
more reactive halogen (up) displaces the ions of the
less reactive halogen
• Final solution can be shaken with hydrogen solvent as
halogens are more soluble in non-polar solvents
• Redox reaction
• Solution becomes yellow/orange/brown/darker;
chlorine is more reactive than iodine (and displaces it
from solution)
Allow correct equation (2KI + Cl2 → KCl + I2) for second
mark or stating that iodine/I2 is formed.
Displacement reactions
2𝐾𝐵𝑟(𝑎𝑞) + 𝐶𝑙2(𝑎𝑞) → 2𝐾𝐶𝑙(𝑎𝑞) + 𝑩𝒓𝟐(𝒂𝒒)
2𝐾𝐼(𝑎𝑞) + 𝐶𝑙2(𝑎𝑞) → 2𝐾𝐶𝑙(𝑎𝑞) + 𝑰𝟐(𝒂𝒒)
2𝐾𝐼(𝑎𝑞) + 𝐵𝑟2(𝑎𝑞) → 2𝐾𝐵𝑟(𝑎𝑞) + 𝑰𝟐(𝒂𝒒)
Cl-(aq)
Br-(aq)
Solution
turns
from
colorless
to
orangeColorless / no reddish brown; as Cl is
Cl2
reaction
more reactive than Br
and displaces it forming
Br₂ - not spontaneous
Br
no reaction
no reaction
I2 no reaction
no reaction
2
I-(aq)
Solutions
turns
from
colorless to brown; as
Cl/ Br is more reactive
than I and displaces I
forming I₂ (w/ black
percipitate?)
no reaction
Silver Halides
• Form insoluble salts with silver
𝐴𝑔𝑁𝑂3 𝑎𝑞 + 𝑁𝑎𝐶𝑙 𝑠 → 𝐴𝑔𝐶𝑙 𝑎𝑞 + 𝑁𝑎𝑁𝑂3 (𝑎𝑞)
𝐴𝑔+ 𝑎𝑞 + 𝑋 − 𝑎𝑞 → 𝐴𝑔𝑋 (𝑠)
Cl-(aq)
Br-(aq)
Pale cream
white) ppt
Ag+ white ppt
reaso insoluble
n formed
AgCl insoluble
formed
equa
Ag+ + Cl- AgCl
tion
Ag+ + Br- AgCl
I-(aq)
(offAgBr
yellow ppt
insoluble AgI formed
Ag+ + I- AgI
13.1 AND 3.3.2 PERIOD 3 ELEMENTS
needed? Physical Properties Period 3
Na
Mg
Al
Si
P
S
Physical
State and
Appearance
at RT
Silver
Solid
Silver
Solid
Silver
Solid
Silver
Solid
White
Solid
Yellow Green
Powder Gas
Gas
Conductivity
in molten
Good
Good
Good
Semi
(Poor)
Poor
Poor
Poor
Structure
Giant
Metallic
Giant
Metallic
Giant
Metallic
Giant
Simple Simple Simple Simple
Molecular Molec Molecu Molecu Atomic
ular
lar
lar
Type
Metal
Metal
Metal
Metalloid
NonMetal
NonMetal
Cl
Poor
NonMetal
Ar
NonMetal
Period 3 Oxides
• React by direct combination on heating
• Each Oxide increases by 1/2 an oxygen
Formula of
Oxide 𝑷𝟒 𝑶𝟏𝟎 (s) 𝑺𝑶𝟐 (𝒈) 𝑪𝒍𝟐 𝑶 (𝒈)/
Physical 𝑵𝒂𝟐 𝑶 (𝒔) 𝑴𝒈𝑶 (𝒔) 𝑨𝒍𝟐 𝑶𝟑 (𝒔) 𝑺𝒊𝑶𝟐 (𝒔)
/ 𝑷𝟒 𝑶𝟔 (𝒔) / 𝑺𝑶𝟑 (𝒍) 𝑪𝒍𝟐 𝑶𝟕 (𝒍)
State under
STP
Oxidation #
+1
+2
+3
Electrical
Conductivity
in molten
State
Good and High – Conductors
Structure
Giant Ionic
Acid – Base
Character
Has ions as freely moving
charged particles
Basic Oxide
Amphoteric
Oxide
+4
+5/ +3
Very
Low
+6/ +4
+7/ +1
None Insulators
Does not have charged particles
Giant
Covalent
Molecular Covalent
Acidic Oxide
Na2O
Adding
H2 O
Adding
HCl
Fully
dissolves
to form
strong
basic sol.
MgO
Semidissolves
to form
weak
basic
solution
𝑁𝑎2 𝑂(𝑠)
𝑀𝑔𝑂 𝑠
+ 2𝐻𝐶𝑙 𝑎𝑞 + 2𝐻𝐶𝑙 𝑎𝑞
→ 2𝐿𝑖𝐶𝑙 𝑎𝑞 → 𝑀𝑔𝐶𝑙2 𝑎𝑞
+ 𝐻2 𝑂 (𝑙) + 𝐻2 𝑂 (𝑙)
Al2O3
Insoluble
SiO2
P4O10 (or
P4O6)
Strong/
Insoluble medium
acid
Al2O3 +
6H+ ->
2Al3+ +
3H2O
No
reaction
Al2O3 +
2OH- +
3H2O ->
2Al(OH)4
SiO2 +
2OH- ->
SiO32- +
H2 O
No
reaction
SO3 (or SO2)
Cl2O7
Strong
acid
Cl2O7 +
H2O ->
HClO4
No
reaction
No
reaction
𝑂2− + 𝐻2 𝑂 → 2𝑂𝐻 −
Neutralization – forms
salt and water
Adding
NaOH
No
reaction
No
reaction
P4O10 +
Cl2O7 +
SO3 + OH12OH ->
OH- ->
2-> SO4
4PO43- +
2ClO4- +
+ H2O
6H2O
H2 O
Basic Oxides in Water
𝑁𝑎2 𝑂(𝑠) + 𝐻2 𝑂 𝑙 → 2𝑁𝑎𝑂𝐻 𝑎𝑞 - asked in pp
• Fully dissolves to form strong basic sol
𝑀𝑔𝑂 𝑠 + 𝐻2 𝑂 𝑙 → 𝑀𝑔(𝑂𝐻)2 (𝑎𝑞)
• semi-dissolves to form weak basic sol
Basic Oxides in Acid
• Solid Na₂O, MgO react with an acid to form a
salt and water
• Oxide Ion combines with two 𝐻+ ions
𝑂2− 𝑠 + 2𝐻+ 𝑎𝑞 → 𝐻2 𝑂 (𝑙)
𝑁𝑎2 𝑂(𝑠) + 2𝐻𝐶𝑙 𝑎𝑞 → 2𝐿𝑖𝐶𝑙 𝑎𝑞 + 𝐻2 𝑂 (𝑙)
𝑀𝑔𝑂 𝑠 + 2𝐻𝐶𝑙 𝑎𝑞 → 𝑀𝑔𝐶𝑙2 𝑎𝑞 + 𝐻2 𝑂 (𝑙)
Phosphorus Acidic Oxides in Water
• Non-metallic oxides react readily with water
to produce strong acidic solutions
• 𝑃4 𝑂10 𝑠 + 6 𝐻2 𝑂 𝑙 → 4𝐻3 𝑃𝑂4 (aq) –
Phosphrous (v) oxide to Phosphoric (v) Acid
• 𝑃4 𝑂6 𝑠 + 6 𝐻2 𝑂 𝑙 → 4𝐻3 𝑃𝑂3 (aq) –
Phosphorous (III) oxide to Phosphoric (III)
Acid (asked in pp)
Suplhur Acidic Oxides in Water
• Non-metallic oxides react readily with water
to produce strong acidic solutions
• S𝑂3 𝑙 + 𝐻2 𝑂 𝑙 → 𝐻2 𝑆𝑂4 (aq) – Sulfur
Trioxide to Sulfuric (VI) Acid – asked in pp
• S𝑂2 𝑔 + 𝐻2 𝑂 𝑙 → 𝐻2 𝑆𝑂3 (aq) – Sulfur
Dioxide to Sulfuric (IV) Acid - asked in pp
Chlorine Acidic Oxides in Water
• Non-metallic oxides react readily with water
to produce acidic solutions
• 𝐶𝑙2 𝑂7 𝑙 + 𝐻2 𝑂 𝑙 → 2𝐻𝐶𝑙𝑂4 (𝑎𝑞) –
Dichlorine Heptoxide to Chloric (VII) acid
• 𝐶𝑙2 𝑂 𝑙 + 𝐻2 𝑂 𝑙 → 2𝐻𝐶𝑙𝑂 (𝑎𝑞) –
Dichlorine Monoxide to Chloric (I) acid
Silicon Oxide with Water
• Silicon Dioxide – insoluble in water – weakly
acidic oxide
• Generally: the bonding changes from ionic to
covalent: solutions in water change from
alkaline to acidic – reflected in the way oxide
bonds to water molecules – affecting the
strength of the O-H bond
Silicon Acidic Oxides with Acid
• Does not react with water, but reacts with
concentrated alkalis to form silicates
• 𝑆𝑖𝑂2 𝑠 + 2𝑂𝐻− 𝑎𝑞 → 𝑆𝑖𝑂2− 3 𝑎𝑞 +
𝐻2 𝑂 (𝑙) – Silicon Dioxide to silicates
Amphoteric (has basic and acidic
propeties) Oxides
• 𝐴𝑙2 𝑂3 – 3+ ion really high lattice enthalpy – insoluble in
water
As a Base
• 𝐴𝑙2 𝑂3 𝑠 + 6𝐻 + → 2𝐴𝑙 3+ 𝑎𝑞 + 3𝐻2 𝑂 𝑙
• 𝐴𝑙2 𝑂3 𝑠 + 3𝐻2 𝑆𝑂4 (𝑎𝑞) → 𝐴𝑙2 (𝑆𝑂4 )3 𝑎𝑞 +
3𝐻2 𝑂 𝑙
As an Acid
𝐴𝑙2 𝑂3 𝑠 + 3𝐻2 𝑂 𝑙 + 2𝑂𝐻 − (𝑎𝑞) → 2𝐴𝑙(𝑂𝐻)4 − 𝑎𝑞
Period 3 Chlorides
Formula of
Chloride and
𝑨𝒍𝑪𝒍𝟑 𝒔 /
𝑵𝒂𝑪𝒍(𝒔) 𝑴𝒈𝑪𝒍𝟐 (𝒔)
𝑺𝒊𝑪𝒍𝟒 (𝒍)
Physical
𝑨𝒍𝟐 𝑪𝒍𝟔 (𝒔)
State at RT
Structure
Electrical
Conductivity
in molten
State
Add H2O
Ionic
Ionic
Ionic
Covalent
𝐏𝑪𝒍𝟓 (𝒔)
𝑪𝒍𝟐 (𝒈
Simple Molecular Covalent
Very
Low
Good and High – Conductors
Dissolve Dissolve Hydrolysi
s to give s to give s to give
free
free
[Al(H2O)6
ions
ions
]3+ and
Cl- ions
𝑷𝑪𝒍𝟑 (𝒍)
Reacts
to
produc
e HCl
and
Si(OH)
4
None - Insulators
Reacts to
produce
H3PO3 and
HCl
Reacts t
o
produce
H3PO4 a
nd HCl
Dissociate
s to give
HOCl and
HCl
13.2 FIRST ROW D-ELEMENTS
Physical Properties of Transition
Elements (less tested on – presumed)
• Stronger metallic bonds than other metals as electrons in
4s and 3d are close in energy and hence are both involved
in bonding and form part of the delocalised sea of electrons
– this results in a high charge density
1. High electrical and thermal conductivity due to large # of
delocalised electrons
2. High melting point ^
3. High tensile strength (can hold large loads) ^
4. Malleable (shaped) and ductile (drawn into wires) – layers
of ions can easily slide across each other – easier as
smaller atomic radii
5. High density – smaller atomic radii compared to s-block
metals
Chemical Properties of Transition
Elements
Define: transitional Elements form more than one
ion with partially filled d-orbitals
1. Variable Oxidation Numbers – electron
movement in d-orbits
2. Complex Ion Formation – eg ligands (can donate
lone pair of electrons to central to central
transition element to form datives)
3. Colored Compounds: not black, white or
colorless – electron movement in d-orbits
4. Have catalytic properties
Sc and Zn not Transition Metals?
• 𝑆𝑐: 𝐴𝑟 4𝑠 2 3𝑑1
• When 𝑆𝑐 3+ : Ar - no electrons in d-orbitals
• 𝑍𝑛: 𝐴𝑟 4𝑠 2 3𝑑10
• When 𝑍𝑛2+ : 𝐴𝑟
3𝑑10 - full d-orbitals
• They both don’t have ions with partially filled
d-orbitals (don’t fit the definition)
Variable Oxidation Numbers in
Transition Metals
• Reason: 3d and 4s electrons have similar energy
levels and the large jump in successive ionization
energy only happens between 4s and 3p hence
still stable at higher OS
All transition metals have +2 (and +3) oxidation
numbers bar
• Cr also has +3 (CrCl3) and +6 (Cr2O72-)
• Mn also has +4 (MnO2) and +7 (MnO4-)
• Fe also has +3 (Fe2O3)
• Cu also has +1 (Cu2O)
Oxidation State Rules (OS)
•
•
•
•
Oxidation states above +3 show covalent character
Compounds with higher OS tend to be oxidizing agents
Compounds with lower OS tend to be reducing agents
Higher oxidation states become less stable compared
to lower ones as you move from left to right across the
series
• Relative Stability of +2 and +3 increases across the
period - +3 more stable initially until Cr with +2 highly
reducing; while +2 more stable after Cr and +3 highly
oxidising
• INSERT IMAGE OF OS
Ligands
• Ligands: a species (molecule or ion) that
donates (hence must have lone pairs) an
electron pair to a metal atom or ion to form to
form a coordination complex
• Ligands can be neutral or ions (positive,
negative)
• Eg. 𝐻2 𝑂 𝑤𝑎𝑡𝑒𝑟 CN¯ cyanide ion Cl¯ chloride
ion NH₃ ammonia
Forming Complexes
• Ligands are lewis bases - donates a lone pair
of electrons
• Metal Ions are lewis acids – accepts lone pair
of electrons
• Each ligand donates an electron pair to the
metal ion forming a dative covalent bond
Co-odrination Number (CN)
• # of dative bonds with a single metal ion is
called the coordination number which
determines the geometric shape of the
complex
Co-ordination #
Shape
2
Linear
4
Square Planar
4
Tetrahedral
6
Octahedral
Complex Ions Examples
• 𝐹𝑒 3+ + (𝐻2 𝑂)6 → [𝐹𝑒(𝐻2 𝑂)6 ]3+
ℎ𝑒𝑥𝑎𝑎𝑞𝑢𝑎𝑖𝑟𝑜𝑛 𝐼𝐼𝐼 𝑖𝑜𝑛
𝑦𝑒𝑙𝑙𝑜𝑤 − 𝑜𝑟𝑎𝑛𝑔𝑒
• 𝐹𝑒 3+ + 6𝐶𝑁 − → [𝐹𝑒(𝐶𝑁)6 ]3−
ℎ𝑒𝑥𝑎𝑐𝑦𝑛𝑖𝑑𝑒𝑟𝑖𝑟𝑜𝑛 𝐼𝐼𝐼 𝑖𝑜𝑛
−𝑟𝑒𝑑𝑖𝑠ℎ − 𝑏𝑟𝑜𝑤𝑛 :
Complex Ions Examples
• 𝐴𝑔+ + 2𝑁𝐻3 → [𝐴𝑔(𝑁𝑂3 )2 ]+
• 𝑑𝑖𝑎𝑚𝑖𝑛𝑒𝑠𝑖𝑙𝑣𝑒𝑟 𝐼 𝑖𝑜𝑛 - colorless
• [𝐶𝑢𝐶𝑙4 ]2− ℎ𝑒𝑥𝑎𝑎𝑞𝑢𝑎𝑐𝑜𝑝𝑝𝑒𝑟 𝐼𝐼 𝑖𝑜𝑛 − 𝑦𝑒𝑙𝑙𝑜𝑤
• (only required to know about monodentate
complexes – only occupies one site in coordinate
sphere?)
Complex Ligand Exchange
• [𝐶𝑢(𝐻2 𝑂)6 ]2+ + 4𝐶𝑙 − → [𝐶𝑢𝐶𝑙4 ]2− + 6𝐻2 𝑂
• Sky Blue – Cu surrounded by 6 𝐻2 𝑂
• Green – partial replacement of 𝐻2 𝑂 by 𝐶𝑙 −
• Yellow – Cu surround by 4 𝐶𝑙 −
• 𝐶𝑢2+ : 𝐴𝑟
3𝑑9
Complex Ions Examples
• [𝐶𝑢(𝐻2 𝑂)6 ]2+ + 4𝑁𝐻3 →
[𝐶𝑢(𝑁𝐻3 )4 (𝐻2 𝑂)2 ]2+ + 4𝐻2 𝑂
• Sky Blue – absorbs red light
• Dark Blue – absorbs yellow light
Colored TM ion complexes
• 5 d-orbitals are degenerate (orbitals of same
energy)
• In presence of ligands, d-obritals split –
because of electric field produced by ligand’s
lone pair of electrons
Colored TM ion complexes
• TM appears coloured due to transitions
(movement of electrons) between the d
orbitals
• TM absorbs certain frequencies of visible light
are absorbed transmitting its complementary
frequency (colour)
• Light is absorbed as energy to promote an
electron from lower to the higher split dorbital
Energy separation
between split d-orbitals
• Color absorbed (hence transmitted) depends on
the energy separation between the orbitals is
∆𝐸 and depends on
1. Identity of TM ion – nuclear charge
2. Oxidation State - # of d electrons present
3. Identity of ligand – higher charge density = more
splitting of the d orbital
4. Co-odination Number – shape of complex ion as
it determines the electric field of ligand’s lone
pair
Catalysts
• Define: 1) speeds up the rate of chemical
reactions by providing an alternative pathway
of 2) lower activation energy 3)unused
Heterogeneous Catalysis
• Where the catalyst is in a different state from
the reactants
• Useful for their ability to allow reactants to
absorb onto their surface where they can
come into closer contact
• Preferred in industry as it can be easily
removed via filtration after use
• makes processes more economical and
efficient
Heterogeneous Catalytic Examples +
Economical Significance of Catalysts
• 𝑉2 𝑂5 𝑖𝑛 𝑡ℎ𝑒 𝐶𝑜𝑛𝑡𝑎𝑐𝑡 𝑝𝑟𝑜𝑐𝑒𝑠𝑠
2𝑆𝑂2 + 𝑂2 ↔ 3𝑆𝑂3
economic significance in 𝑆𝑂3 reacting with 𝐻2 𝑂 to
form 𝐻2 𝑆𝑂4 (world’s most important chemical)
• 𝐹𝑒 𝑖𝑛 𝑡ℎ𝑒 𝐻𝑎𝑏𝑒𝑟 𝑃𝑟𝑜𝑐𝑒𝑠𝑠
2𝑁2 + 3𝐻2 → 4𝑁𝐻3
economic significance in its use as toilet cleaners ,
raw material for Ammonium Nitrate fertiliser and
explosives and other chem. for drugs, plastics
MORE EFFICIENT AND ECONOMIC
Heterogeneous Catalysts
• Ni in the conversion of alkenes to alkanes
eg unsaturated oil to margarine
1
2
4
3
Heterogeneous Catalytic Examples
• 𝑀𝑛𝑂2 𝑖𝑛 𝑡ℎ𝑒 𝑑𝑒𝑐𝑜𝑚𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛 𝑜𝑓 𝐻𝑦𝑑𝑟𝑜𝑔𝑒𝑛 𝑃𝑒𝑟𝑖𝑜𝑥𝑖𝑑𝑒 (𝐻2 𝑂2 )
1
2
𝐻2 𝑂2 → 𝐻2 𝑂 + 𝑂2
used as rocket fuel
• Palladium (Pd) and Platinum (Pt)
2𝐶𝑂(𝑔) + 2𝑁𝑂(𝑔) → 2𝐶𝑂2 𝑔 + 𝑁2 𝑔
Catalytic converters on cars to remove harmful primary pollutants
from a car’s exhaust gases to reduce the greenhouse effect
Homogeneous Catalysts
•
•
•
•
in the same state as the reactants
Useful for their variable oxidation states
Often in redox reactions
They mix effectively with the reactants hence
work under mild conditions of the human
body
• Many enzyme-catalysed cell reactions in the
body involve TM as homogeneous catalysis
and are of fundament biological importance
Homogeneous Catalysts
• Fe²⁺ in heme.oxygen – transported through the
bloodstream by forming a weak bond with the
heme group of hemoglobin (central Fe²⁺ ion
surrounded by 4 Nitrogen atoms) O₂ - Fe²⁺ bond
is easily broken when needed
• Co³⁺ octahedral complex in vitamin B₁₂ w/ 5 sites
occupied by nitrogen atoms, leaving the 6th site
available for biological activity. vitamin B₁₂ is
needed for the production of red blood cells and
a healthy nervous system