Download Chem Periodicity, Reactivity, Redox 2009 Yingxin

Periodicity
Horizontal Rows = Period (numbered 1,2,3,4,5,6,7)
Vertical Columns = Group (numbered I, II, III, IV, V, VI, VII, 0/ VIII)
Groups I, II, III, IV, V, VI, VII, VIII are main group elements
Elements between groups II and III (3-12) are transition elements
Importance of the Periodic Table
@ Arrangement of elements help us understand them better (Groups have similar properties)
@ Easier to remember properties of several elements
@ Make predictions about other elements
@ Arranged by atomic number
Patterns
Electronic Structure
@ Group
@ Represents the number of valence electrons in an atom of the element
@ Similar electronic configurations
@ Leading to similar chemical properties
@ Period
@ Refers to the number of electron shells in an atom of the element
Charges on Ions
Elements on the left hand side (Group I, II)
@ Lose their valence electrons (tendency to lose
electrons)
@ Form positive ions (Cations)
@ Charge on positive ions correspond to group
number
Elements on the right hand side (Group III, IV, V, VI, VII, VIII)
@ Gain electrons to obtain a noble gas structure
@ Form negative ions (anions)
@ Charge on negative ions correspond to the number of
electrons gained = 8-(group number) usually
Bonding
@ Elements in same group form same type and number of bonds because they have the same number of valence
electrons
@ Probably will form compounds with similar formulae
@ Exceptions: Group IV, V, VI, as metallic character varies down the group
Metallic Character
@
@
Metals
Non-metals
- Left hand side of the Periodic Table
- Right hand side of the Periodic Table
- Groups I and II
- Groups III, IV, V, VI, VII, VIII/ 0
- Few valence electrons
- Many valence electrons
- Tendency to lose electrons
- Tendency to gain electrons
- Form cations
- Form anions
Elements change from non-metal to metal across a period
Elements adjacent to dividing line are metalloids (have properties of metals and non-metals)
Periodic Trends
(always opposite to group trends)
Melting Points
@  from Group I to Group IV due to a change in bonding from metallic [strong] (I, II, III) to giant covalent or
macromolecular (IV)
@  from Group IV to Group VIII as the nature of the bonding within the element changes from giant covalent (IV)
to simple covalent [discrete molecules] (V, VI, VII) to monoatomic atoms [weak Van der Waals forces of
attraction] (VIII)
Electrical Conductivity
@ Decreases across a Period as the nature of the elements changes from metallic to metalloid to non-metallic
@ Metallic: good conductors of electricity due to their delocalized electrons
@ Metalloid: semi-conductors
@ Non-metallic: insulators
Acid/Base Nature of Oxides
@ Oxides of metallic elements are basic
@ Oxides of metalloids are amphoteric (Definition: Having the characteristics of an acid and a base and capable
of reacting chemically either as an acid or a base.)
@ Oxides of non-metallic elements are acidic
Oxidising and Reducing Power
@ Metallic elements are good reducing agents—they readily lose electrons to other chemicals
@ Non-metallic elements are good oxidizing agents—they readily remove electrons from other chemicals.
Electronegativity 
Tendency to attract electrons to itself
Atomic Radius
- Measured by the distance between the nuclei of 2 touching atoms / 2
Type of Bond
Radii
Metallic Bond
Metallic radii
Covalent bonds (non-metals) Covalent radii
No Bond (e.g. Noble gas)
Van der Waals radii
Down the Group
Atomic radius increases down the group
Across the Period
Atomic radius decreases across the period
Reason: Down the group, an extra electron shell is
added, increasing the size of atoms.
Reason: Across the period, the number of protons increases,
increasing effective nuclear charge. The increasing effective
nuclear charge exerts a greater force of attraction on the
electrons, pulling the valence electrons closer to the nucleus.
Greater screening effect, decreasing forces of
attraction
Effect: Valence Electrons further away from
nucleus, thus the force exerted by the nucleus on
the valence electrons is reduced.
Ionic Radius
Positive Ions (Cations)
Greater nuclear force of attraction, stronger bonds, so it’s
smaller
Same amount of screening across period
Negative Ions (Anions)
Cations have a SMALLER ionic radius than their atoms.
Negative ions have a BIGGER ionic radius than their
atoms
Reason: Cations lost all their valence electrons. Cations
have less electrons than their atoms. Thus, the same
number of protons exert the same amount of force on
less electrons, thus, the electrons of the cations will be
pulled in even nearer to the nucleus.
Reason: There is a greater number of electrons
attracted by the same number of protons. Extra
repulsion produced by incoming electron causes ion to
expand.
Decrease in ionic radius across the period from Group I
to III
Decrease in ionic size across the period from Group V to
VII
Reason: Electropositive metals -lose electrons
- Number of protons increases => Nuclear
Charge increases
- Number of electrons stays the same (all
attained noble gas configuration) => Shielding
Effect stays constant
Reason: Electronegative elements -gain electrons
- Number of protons increases => Nuclear
Charge increases
- Number of electrons stays the same (all
attained noble gas configuration) => Shielding
Effect stays constant
Ionisation Energy
Definition:
First Ionisation Energy (I.E) is the energy required to remove the MOST LOOSELY held electron from 1 mole of
atom to form 1 mole of ion with charge +1.
@ Can only occur in the gaseous state, so MUST put state symbols (g)
@ Unit: kJ mol-1
@ All elements have I.E. (despite metal/non-metal)
Factors Affecting I.E
1. Charge on nucleus (Proton number)
The more the protons there are in the nucleus, the stronger electrons are attached to the nucleus. =>> Greater I.E.
2. Distance of electron from nucleus (atomic radius)
The smaller the distance, the more strongly attached the electron is to the nucleus =>> Greater I.E.
3. Number of electrons between valence electrons and nucleus (screening effect)
The greater the number of electrons between the valence electrons and the nucleus, the less the net pull from the
nucleus experienced by the valence electron, the greater the screening/ shielding effect, the less attached the
valence electrons are to the nucleus =>> Smaller I.E.
The lower the I.E., the higher the reactivity.
Down a Group
General trend is for I.E. to decrease down a group
Reason:
The proton number greatly increases down the group, greatly
increasing nuclear charge.
Across a Period
General trend is for I.E. to increase
across a period
Reason:
The screening effect of valence
electrons is the same across the period.
But, there are also much more electrons between valence electrons
and the nucleus down the group. Down the group, an extra shell and
layer of electrons is added, increasing the screening effect and
However, across the period, the number
of protons in the nucleus increases. This
increases the nuclear charge, and hence
decreasing the net pull from the centre experience by the valence
electron.
Effect of extra protons is compensated for by the effect of extra
screening electrons.
Down the group, the distance between valence electrons and nucleus
greatly increases (atomic radius increase). Thus, the force of attraction
on the valence electrons is reduced- decreasing I.E.
Easier to lose electrons – more screening, weaker attraction
the valence electrons are more strongly
attached to the nucleus- increasing the
I.E.
The increased nuclear charge also pulls
the valence electrons nearer to the
nucleus, reducing the distance between
them- increasing I.E
Greater force of attraction, harder to
lose electrons
Group Properties
Group I elements: Alkali Metals
@
@
@
@
@
@
@
@
@
@
@
@
Most reactive metals
1 valence electron only
@ Which is lost when they react, forming a cation with +1 charge
Soft, can be cut easily (malleable)
Shiny, silvery solids
Salts formed are white in colour (eg NaCl)
Low melting points and boiling points
@  down the group
@ Atomic radii of atoms increase down the group due to addition of shell, thus the forces of attraction
between the positive ions and the negative valence electrons are weaker – lower charge density
(same charge over bigger area)
Low densities
@  down the group
@ Density = mass/volume
@ Though both mass and volume increase, the increase in mass is greater than the increase in volume
@ Lithium, Sodium, Potassium float on water
Metals tarnish easily in air (react with air)
@ Must be kept under oil to prevent contact with air
React with water to give alkaline solutions (soluble salt + hydrogen gas)
@ Alkaline – turn red litmus paper blue
@ React vigorously with cold water to produce hydrogen gas and a solution of metal hydroxide
@ Catch fire and may even cause explosions
@ Due to their high reactivity, they are powerful reducing agents
Similarities
Differences
Li, Na, K float on water
Li: no flame observed
Li, Na, K move about on surface of water
Na: hydrogen burns with yellow flame
Faster reaction than Li
Effervescence is observed
K: hydrogen burns with lilac flame
Faster reaction than Na
Hissing sound is heard
Alkaline solution produced
Reacts with acid to form salt and hydrogen gas
All ionic, cations of charge +1 (similar formulae)
Reactivity
@ Increases down the group
@ Number of shells increase down the group (according to period number)
@
@
Increasing the shielding/screening effect down the group (which outweighs the increased nuclear
charge)
@ Electrons are lost easier as the nuclear force of attraction on these valence electrons is decreased
@ Valence electrons bond with other electrons and form compounds easier
@ Lithium least reactive, Caesium most reactive
@ Greater reactivity, more readily give out electrons = easily oxidized = strong Reducing Agent
Conduct electricity
Group II elements
@
@
@
@
Readily lose both valence electrons to form cation with a 2+ charge.
Down the group, the reactivity of Group II metals increases
Reacts with acids to form salt and hydrogen gas
Displacement Reaction: When a more reactive metal reacts with the compound of another less reactive metal,
the metal will be displaced from the compound.
Group VII elements: Halogens
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@
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@
@
@
Exists as diatomic molecules in natural state
Reactivity
@ Most reactive non-metals
@ Fluorine is the most reactive element in the periodic table
@ Hence, they are powerful oxidizing agents
@ React vigorously with most metals to form ionic salts
@ Reactivity decreases down group
@ Decreased oxidizing power, less ability to gain electrons due to more shells = easily reduced = strong
Oxidising Agent
@ Halogens have more shells down the group, increasing the shielding effect down the group. This
makes it harder for the halogen to gain electrons as the forces of attraction between the nucleus and
the electrons are not very high.
Displacement reaction: More reactive halogen will displace less reactive one from an aqueous solution of its
ions
@ Eg Cl2 + 2KBr  2KCl + Br2
@ Solution will turn the colour of the displaced halide (in this case, reddish-brown)
Reaction with silver nitrate
@ When Ag(NO3)2 is added to Cl or I, a precipitate will be formed.
@ Silver chloride/silver iodide is insoluble in water
7 valence electrons
@ Tendency to gain 1 electron to form an ion of charge -1
Do not conduct electricity
Low melting points and boiling points
@ Increase down group
@ As molecules get larger, Van der Waals forces of attraction are stronger
Gas > Solid down group
Colour intensity increases down the group
Element Colour
Fluorine Pale Yellow Gas
Chlorine Greenish yellow gas
Bromine Reddish brown liquid
Iodine
Purple black solid *purple gas – can sublime!
Uses
Halogen
Application
Fluorine
@
Chlorine
@
@
@
@
@
@
@
Bromine
Iodine
@
@
@
Compounds of fluorine (eg KF) are used in toothpaste and drinking water, to strengthen tooth
enamel and therefore reduce tooth decay
Non-stick polymer Teflon
Chlorofluorocarbons (CFCs), used in refrigerators and aircon units
Anaesthetic Halothane
Bleach, disinfectants, antiseptics
Polymer PVC
Chlorofluorocarbons (CFCs), used in refrigerators and aircon units
Silver bromine (AgBr) is sensitive to light (photosensitive) – used in manufacture of
photographic films. When light strikes the film, the silver ions are reduced to form small dark
crystals of silver metal. This produces the negative from which the actual photograph is
developed.
Pesticides and tear gas
Required by human body to make the hormone thyroxine
Mild antiseptic
Transition Metals
@
@
Typical metals
@ High tensile strength
@ High melting and boiling points
@ Depends on strength of metallic bonds – number of valence electrons
@ MP very high – in the thousands
@ High density
@ Sonorous
@ Ductile
@ Malleable
@ Good conductor of heat
@ Good conductor of electricity
@ Shiny
Form coloured compounds
Compound
Colour
Iron (II) Sulphate Fe2+
Green
3+
Iron (III) Chloride Fe
Brownish-yellow
Potassium Manganate (VII) KMnO4
Purple
Potassium Dichromate (VI) K2Cr2O7
Orange
Copper (I) oxide Cu+
Reddish Brown
Copper (II) oxide Cu2+
Black
Copper Sulphate CuSO4
@ Reactions = colour changes
Variable Oxidation States
@ Can form ions with different oxidation states (eg Fe2+ and Fe3+)
@ Each atom can lose/gain different number of electrons based on conditions of reaction
@ Not as reactive as alkali metals/halogens
@ Catalysts
@ Increase rate of chemical reactions
@ Eg Fe3+ is used in the Haber Process – produce ammonia
@ Eg Vanadium Oxide is used in the Contact Process – produce sulphuric acid
Uses
@ Titanium
@ High MP, low density, won’t rust
@
@
@
@
@
@ Aircraft
Chromium
@ Hardens steel
@ Manufacture stainless steel, forming protective layer (resistant to corrosion)
Iron
@ Form alloys (eg steel)
Copper
@ Good conductor of electricity
@ Easily bent into shape (malleable)
@ Does not react with H2O
@ Electricity cables, water pipes
Zinc
@ Form alloys (brass)
@ Galvanise metals, prevents corrosion
Noble gases
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@
Least reactive elements in the Periodic Table
Full valence shells – stable electronic configuration
@ Except for Helium (2), they all have 8 electrons in valence shell
@ Chemically unreactive
@ Do not need to gain/lose electrons (bond with other atoms) to obtain stable configuration
@ Colourless
@ Exist as monoatomic gases
@ Very low melting and boiling points
@ Increase down the group
@ Van der Waals’ forces of attraction increases down the group, but still quite low as they are single
atoms (do not form compounds)
@ Density increases down the group
@ Atomic mass increases
@ So density increases too
@ He, Ne: lighter than air. Ar, Kr, Xe: heavier than air.
@ Used because of lack of reactivity
Noble Gas
Property
Application
Helium
Very light, not combustible
Filling balloons and airships
Neon
Glows brightly when a current passes
Used for advertising signs
through the gas
Neon lights (disco!)
Argon
Unreactive
Provides an inert atmosphere inside light bulbs
Krypton and Xenon
Used in fluorescent bulbs, flash bulbs and lasers
Radon
Radioactive
Treat cancer and detect leaks in gas pipelines
Valence
Electrons
Lose/ Gain
Charge on
Ion
Metal/Non-
Alkali Metals
1
Halogens
7
Lose
+1
Gain
-1
Noble Gases
2 (Helium), usually
8
Neither
Neutral
Metal
Non-metal
Non-metal
Transition Metals
Varies
Varies
1 element may
form ions with
different oxidation
states (lose/ gain
different number
of electrons)
Metal
metal
State of
Solids
Matter
Atomic
Structure
Melting
Low melting points
Point/
 Generally decrease
Boiling Point
down the group
 Atomic radii of atoms
increase due to the
addition of shell, thus
the forces of attraction
are weaker
Density
Colour
Electrical
Conductivity
Reactivity
Low densities
 Generally increase down
the group
 Density= mass/ volume
 Though both mass and
volume increase, the
increase in mass is
greater than the
increase in volume
Silvery
Gas > Solid down the group
Gases
Diatomic
Monoatomic
Low melting points and
Very low melting
boiling points
and boiling points
 Increase down the group  Increase down
 Van Der Waals forces of
the group
attraction are stronger
 Van Der Waals’
down the group
forces of
attraction
increase; but still
quite low as they
are single atoms
Density increases
down the group
 Atomic mass
increases down
the group
High melting
points
Elements become darker
down the group
Colourless gases
Varies (Coloured)
Characteristic
colours
DOES NOT REACT
INERT
Not as reactive as
alkali metals or
halogens
High density
No
Most reactive metals in
the periodic table
Reactivity increases
down the group
 Number of shells
increase down,
increasing the
shielding/ screening
effect down the group
 Electrons are lost
easier
Very Reactive (fluorine is
the most reactive element)
React vigorously with most
metals to form ionic salts
Reactivity decreases down
the group
 more shells, increasing
the shielding effect
 Harder for the halogen
to gain electrons
Reactivity
Physical properties of metals
Atoms in a metal are held together by strong metallic bonds
In a metal lattice, a “sea” of mobile electrons surrounds a lattice of positive metal ions
Attribute
Appearance
Melting & Boiling Point
Metals
Shiny surface
Usually high
Physical state
Density
Electrical conductivity
Solid (except mercury)
Usually high
Good conductor
Thermal conductivity
Good conductor
Strength
Ease of shaping
Strong
Ductile and malleable
Reason
Metallic bonding is strong, hence a lot of energy is required to
break the bonds and separate the atoms
Presence of mobile delocalized electrons through the metal
that carry the electrical current
When one end of a metal is heated, the delocalized electrons
get heated and heat energy is converted into kinetic energy.
They move faster and collide with neighbouring particles,
transferring heat energy through collisions.
Layers of atoms in a metal can slide over each other easily
when a force is applied. The metal does not break easily as the
“sea” of electrons holds the atoms in the metal together.
Chemical properties
Metals
React with oxygen to form basic oxides or amphoteric oxides
Form positive ions by losing electrons
Form ionic chlorides and oxides
Most metals react with dilute acids to give hydrogen and a salt
Non-metals
React with oxygen to form acidic oxides or
neutral oxides
Form negative ions by gaining electrons
Form covalent chlorides and oxides
No reaction with dilute acids
Reactivity of Metals
@ With water/steam
@ Some metals react with water/steam, some don’t
@ If it does, it produces hydrogen gas and an oxide/hydroxide of the metal
@
@
M [metal] (s) + H20 (l)  H2 (g) + MO or MOH
With Dilute Hydrochloric Acid
@ Many metals react with dilute hydrochloric acid to produce a metal chloride and hydrogen gas
@ M [metal] (s) + H20 (l)  H2 (g) + MCl (aq)
Metals
Potassium
Sodium
Calcium
Magnseium
*Aluminium
Zinc
Iron
H2O
Can react with cold water
K: vigorously, heat, fire, explosion
Na: vigorous, may catch fire/explode
Ca: readily, bubbles produced
Eg: Ca + 2H2O  Ca(OH)2 + H2
Needs to react with steam
Mg: vigorously (very slow in cold water)
Zn: less vigorously
Will NOT get hydroxide! Only metal oxides
HCl
Can react with cold acids
K, Na: explodes (dangerous!)
Ca: very fast, bubbles produced
Eg: 2K + 2HCl  2KCl + H2
Needs to react with warm acid
Mg: fast
Zn: moderately fast
Fe: slowly
Tin
Lead
(steam is a dry gas)
Eg: Zn + H2O  ZnO + H2
Copper
Mercury
Silver
Gold
Pb: very slowly with warm HCl
No reaction
Less reactive than hydrogen!
Alloys
@ Pure metals are not widely used as they are too soft and have a low resistance to corrosion
@ Mixture of 2 or more metals (mix molten elements in right proportions and allow them to solidify)
@ More useful physical properties than pure metals
@ Harder and stronger (less malleable because different sized atoms prevent slipping)
@ Appearance can be enhanced
@ Improve resistance to corrosion
@ Lower melting point
Reactivity series
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@
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Arrangement of metals in order of reactivity
Order determined by experimental observations in lab
Reactive metals are good RAs (they want to get oxidized, lose electrons)

Please
Stop
Calling
Me
Element
Potassium
Sodium
Calcium
Magnesium
A
Aluminium
Crazy
Zebra
I
Truly
Love
Hamburgers
Carbon
Zinc
Iron
Tin
Lead
Hydrogen
@
Very reactive (reacts vigorously, less
activation energy used)
@ Requires electrolysis to extract pure
metal from ore
@ Readily gives up electrons to form
positive ions
@ Corrodes easily
@ Tendency to form compounds
Special!
@ It will not react with dilute acids to
produce hydrogen because it reacts
readily with oxygen in the air to form
aluminium oxide Al2O3
@ Always covered with a layer of oxide,
preventing reaction with acid
Can be reduced by carbon
Important benchmark (ref. point)
@ Metals less reactive than hydrogen do
not react with acids (to produce
hydrogen gas)
@ Metals more reactive than hydrogen
react with acids/water/steam to
Reactivity
@ Reacts
vigorously
(fast) with
chemicals
@ Readily gives
up electrons
in reactions
to form
positive ions
@ Corrodes
easily
Symbol
K
Na
Ca
Mg
Valency
1
1
2
2
Al
3
C
Zn
Fe
Sn
Pb
H
4
2
2, 3
2, 4
2, 4
1
Containing
M
S
G
Copper
Mercury
Silver
Gold
@
@
@
@
produce hydrogen gas
Can be reduced by carbon
But usually are unreactive, so they
don’t even form compounds
Does not react vigorously with
chemicals
Tendency not to form compounds
Cu
Hg
Ag
Au
1, 2
1
1, 3
Uses
@ Predict chemical reactions that can occur
@ Helps in extraction of metals and prevention of rusting
Displacement Reactions
A metal can displace another metal from a compound
Displacement of Metals from Solutions
A more reactive metal will displace the ions of a less
reactive metal in the reactivity series from solution
Displacement of Metals from Metal Oxides
A more reactive metal will take the oxygen from the
oxide of a less reactive metal
@
@
@
The more reactive metal gives up electrons more
readily to form positive ions
(atom -> ion)
Electrons are transferred from the more reactive
metal to the less reactive ions
(ions -> atom)
@
@
e.g. Zn (s) + CuSO4 (aq)  Cu (s) + ZnSO4 (aq)
Zn (s)  Zn2+ + 2eCu 2+ (aq) + 2e-  Cu (s)
The more reactive metal gives up electrons to the
less reactive metal ion
Such that the more reactive metal becomes an ion
and the less reactive metal becomes an atom
Thermit reaction (great amount of heat produced)
@ Used for on-site welding of large meal objects
such as railway lines
e.g. Mg (s) + CuO (s)  MgO (s) + Cu (s)
Mg (s)  Mg2+ (s) + 2eCu2+ + 2e-  Cu (s)
The more reactive metal displaces the less reactive metal
@ A more reactive metal has a higher tendency to lose its valence electrons to form positive ions
@ A more reactive metal is more ready to form compounds
Reaction of Metal Oxides with Carbon
Carbon can remove the oxygen from metal oxides
(only the oxides of metals that are below carbon in the
reactivity series)
Carbon acts as the reducing agent, and reduces the
oxides of less reactive metals.
The higher the metal (oxide) is on the reactivity series,
the stronger the heat required for carbon to “steal” the
oxygen
@ The more reactive a metal is, the more difficult it is
to split up its oxide (because it has a higher
tendency to form compounds)
2CuO (s) + C (s)  2Cu (s) + CO2 (g)
Reaction of Metal Oxides with Hydrogen
Hydrogen can remove the oxygen from metallic oxides
to produce the metal and water.
Hydrogen acts as a reducing agent, and reduces the
oxides of less reactive metals.
Hydrogen can only “steal” oxygen from less reactive
metal(lic oxides). The reaction is easier when the oxide
is of a less reactive metal.
PbO (s) + H2 (g)  Pb (s) + H2O (l)
Oxides that cannot be reduced by carbon (can only be
reduced by electrolysis)
Potassium
Sodium
Calcium
Magnesium
Oxides that can be reduced by carbon (below carbon)
Zinc
Iron
Lead
Copper
Silver
@
@
@
Oxides that cannot be reduced by hydrogen (can only
be reduced by electrolysis)
Potassium
Sodium
Calcium
Magnesium
Zinc
Oxides that can be reduced by hydrogen
Iron
Lead
Copper
Silver
This method is often used to extract metals from
their ores
The lower the position of a metal in the reactivity series, the easier it is for hydrogen/ carbon to remove
oxygen from the metal oxide
Hydrogen/ carbon acts as the reducing agent, to reduce the oxides of the less reactive metals.
Electrolysis
@ Passing an electric current through molten compound
@ Extract more reactive metals like potassium, soldium, magnesium, alumnium
Ionic equations
@ Eliminate the spectator ions
@
Original: Mg(s) + CuSO4 (aq)  MgSO4 (aq) + Cu (s)
@
Write in ions: Mg (s) + Cu2+ (aq) + SO42- (aq)  Mg2+ (aq) + SO42- (aq) + Cu (s)
@
Take out the spectators: Mg (s) + Cu2+ (aq)  Mg2+ (aq) + Cu (s)
Decomposition of Metal Carbonates
When heated, Metal Carbonates  Metal oxide + Carbon dioxide gas
The more reactive the metal, the more stable the carbonate is, the harder to decompose
The less reactive the metal, the more easily the metal carbonate decomposes when heated
Carbonate
Potassium
Sodium
Calcium
Magnesium
Aluminium
Zinc
Iron (II)
Lead
Copper (II)
Silver
Observation
Unaffected by heat
Does not decompose
Decomposes into metal oxide and carbon
dioxide
CuCO3  CuO + CO2
Reason
Too reactive
Stable (eg Na2O3)
Decomposes into silver and carbon
dioxide
Silver oxide that is produced is thermally unstable,
hence it further decomposes to form silver
Rusting
@
@
@
@
@
@
Exothermic process
Slow oxidation of ions to form Hydrated Iron (III) Oxide
Corrosion of most metals
Rust
@ Reddish brown substance formed on the surface of metals (Hydrated Iron (III) Oxide)
@ Brittle and flaky
When a metal corrodes, the rusted surface will flake away, producing a new surface to corrode
Eventually, all the metal will rust and flake away
Conditions
@ Air and water required
@ Other factors that speed up rusting
@ Dissolved salt (NaCl)
@ Iron objects near the sea rust faster because of salt
@ Other acidic pollutants (eg sulphur dioxide, carbon dioxide)
Prevention of Rust
@ Using a protective layer (surface protection)
@ Coat metal with a layer of substance (eg paint/grease/plastic)
@ Layer prevents water and air from coming into contact with the metal surface
@ Sacrificial metals
@ Attach metals that are more reactive to the metal object, so the more reactive metal will corrode
instead of the less reactive metal object
@ Opposite: less reactive metals will increase the rate of rusting
@ Using alloys
@ Best known rust-resistant alloy of iron: Stainless Steel
@ Upon exposure to air and water, a hard coating of chromium oxide forms on the surface of stainless
steel, protecting it from further corrosion
@ Alloys: size of added element differs from main metal atoms
Examples of Application
Method
Painting
Oiling/ Greasing
Plastic Layer
Zinc-plating
Tin-plating
Chrome-plating
More reactive metal block
Stainless Steel
Application
Large objects e.g. vehicles, bridges
Machinery and tools
Small iron/ steel objects e.g. wires
Buckets, dustbins, kitchen sinks
Food cans
Taps, kettles, bicycles
Underground pipes, ships
Cutlery, surgical instruments
Remarks
Metal will rust if the paint is scratched
Also helps in lubrication
Metal will rust if plastic layer is scratched
Metal doesn’t rust even if zinc layer is damaged
Metal rusts if coating is scratched
Also makes it shiny
More reactive metal will corrode instead of iron
Form chromium oxide layer on the metal
Redox
A reaction in which one substance is oxidized while another is reduced at the same time is called a Redox reaction.
Oxidation
Reduction
OXYGEN
Gain of Oxygen by a substance
Loss of Oxygen by a substance
e.g. 2Ca (s) + O2 (g)  2CaO (s)
e.g. Zn (s) + CuO (s)  ZnO (s) + Cu (s)
Calcium gained oxygen, thus, it has been oxidised
Copper loses oxygen, thus it has been reduced
HYDROGEN
Loss of hydrogen by a substance
Gain of hydrogen by a substance
e.g. 2NH3 (g) + 3CuO (s)  N2 (g) + 3Cu (s) + 3H2O (l)
e.g H2S (g) + Cl2 (g)  2HCl (g) + S (s)
Ammonia has lost hydrogen, thus it has been oxidised
Chlorine gains hydrogen, thus it is reduced
ELECTRONS
Loss of electrons from a substance
Gain of electrons by a substance
Put in the charges (even for compounds)
e.g. 2 Na (s) + Cl2 (g)  2Na+ Cl- (s)
+ e.g. 2 Na (s) + Cl2 (g)  2Na Cl (s)
Gain of electrons = negative charge
Loss of electrons = positive charge
Chlorine gains electrons, thus it is reduced
Sodium loses electrons, thus it has been oxidised
*Rewrite the equation in the half equations (including addition of electrons where appropriate) to see if the
substance has gained or lost electrons clearly
1.
2.
3.
Write the before and after of the substance
Where should electrons be added such that the reaction makes sense?
Must include STATE SYMBOLS (molten/melted = liquid (l), solution = aqueous (aq)
Na (s)  Na+ (s) + _______
To equate the reaction, the ___ must be a eSodium has lost electrons, thus it has been oxidised.
Cl2 (g) + 2e-  2Cl- (s)
Electrons have been gained
Thus, chlorine has been reduced
Some Redox Reactions involving Transfer of Electrons
@ Metal and Dilute Acid
@ Displacement reactions
OXIDATION STATE
An increase in oxidation state during a reaction
A decrease in oxidation state during a reaction
Oxidation state (number) is the charge an atom would have if it existed as an ion
1. Work out oxidation state of EVERY element
2. Compare before and after a reaction
3. Did the oxidation state of the element increase or decrease?
How to Work out the Oxidation State?
A few rules:
- Every atom in a PURE element has an oxidation number of 0
- The oxidation number of a simple ion (pure element) (monoatomic ions) is the charge of the ion
- The sum of oxidation numbers in a neutral compound is 0
-
The sum of the oxidation numbers of elements in an ion is equal to the charge of the ion
-
Some fixed oxidation numbers:
H : +1 when bonded to non- metals; -1 when bonded to metals
O: -2 (except in peroxides when it is -1) eg H2O2
F: -1
Cl, Br and I: -1 (unless combined with O or F)
Group I elements: + 1
Group II elements: +2
Group III elements: +3
Some examples of redox reactions involving changes in oxidation states
1. Reaction of metals with dilute acid
2. Halide displacement
3. Extraction of metals (Carbon oxidising the other substance)
Elements with variable O.S.
@ Manganese: +2 (MgCl2), +4 (MnO2), +7 (KMnO4)
@ Chromium: +2 (CrCl2), +3 (CrCl3), +6 (K2Cr2O7)
@ Iron: +2 (FeCl2), +3 (FeCl3)
@ Sulphur: -2 (FeS), +4 (SO2), +6 (H2SO4)
@ Carbon: +2 (CO), +4 (CaCO3)
Oxidizing agents and Reducing agents
OA
RA
Causes another substance to get oxidized
Causes another substance to get reduced
Gets reduced
Gets oxidized
Oxidizing Agents
@ Gain electrons readily (eg Fluorine)
@ Chlorine
@ Oxygen
@ Acidified potassium manganate (VII) (potassium permanganate) KMNO4
@ Acidified potassium chromate (VI) (potassium dichromate) K2Cr2O7
Reducing Agents
@ Lose electrons readily (eg Zinc)
@ Metals (especially those high in the reactivity series)
@ Carbon monoxide
@ Hydrogen
@ Carbon
@ Potassium Iodide (KI)
Test for Oxidising Agent
@ Aqueous potassium iodide (KI) is a common RA.
@ When it reacts with an OA, it is oxidised.
@ It will turn from colourless to brown.
2I- (aq)  I2 (aq) + 2e- (half reaction for oxidation)
Colourless  Brown
Reason: Iodide ion is colourless. However, when it reacts with an OA, it is oxidised and loses electrons. It is oxidised
to an iodine atom, which is brown in colour.
KI is used to test for the presence of OA, and will turn from colourless to brown in the presence of OA.
Starch Iodide Paper
@ Can also be used to test for presence of OA. When it is oxidised by the OA, it will turn from white to blue.
@ Reason: When the iodide ion is oxidised, it will lose electrons to form iodine atom. The iodine will react with
starch to give blue colour.
Tests for Reducing Agent
@ Acidified potassium dichromate solution (aqueous potassium dichromate + dilute sulphur acid) is orange.
@ When it reacts with a RA, it is reduced.
@ Its colour will change from orange to green because it will be reduced to chromium ions.
Cr2O72- (aq)  Cr2+ (aq)
Orange  Green
K2Cr2O7 is used to test for the presence of RA, and will turn from orange to green when it is reduced by the RA.
@
@
@
Acidified potassium manganate (VII) (potassium permanganate) is purple.
When it reacts with a RA, it is reduced.
The solution turns from purple to colourless.
MnO4  Mn2+
Purple  Colourless
KMNO4 is used to test for the presence of RA, and will turn from purple to colourless when it is reduced by the RA.
Reactions
that are NOT redox reactions
1.
Decomposition of metal carbonates
@ Metal carbonate  Metal Oxide + Carbon dioxide
@ Oxidation state of each element remains unchanged
2.
Neutralisation Reactions
@ Acid + Alkaline
@ Oxidation state remains unchanged
3.
Precipitation Reactions
@ AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3
@ Aqueous + aqueous form precipitate in SOLID.
Recognising Redox Reactions
1. Work out oxidation state of EVERY element (remember the rules)
2. Compare before and after a reaction
3. Did the oxidation state of the element increase or decrease?
@
Best method, as some reactions may not involve oxygen/hydrogen/transfer of electrons
@
But the oxidation state method applies to all reactions
Balancing Redox Reactions
@ Mass and charge must be balanced
@ Electrons lost and gained in both sides of the equation must be equal
Step-by-step:
1.
2.
3.
4.
5.
6.
7.
8.
Calculate the oxidation states of each element before and after the reaction
Identify which element undergoes oxidation (reducing agent) and which undergoes reduction (oxidising agent)
Write the half equations for oxidation and reduction
Balance out the mass (same number of each element on each side)
@
ACIDIC CONDITIONS: Use H2O or H+
@
BASIC CONDITIONS: Use H2O or OH@
Make sure the non-O and non-H atoms are balanced first (If they aren’t balanced, add numbers in front to
make them balanced)
@
Balance out the OXYGEN first, (1 side has x oxygen atoms, add H 2O to make the other side have x oxygen
atoms too)
@
Then, add H+ to balance out the extra hydrogen
Balance the charges
@ Add electrons to make the charge on LHS = RHS
@ The number IN FRONT of the ion must be multiplied by the ion charge
@ E.g. the charge of 5HSO4 - is NOT -1, it is 5 X -1 = -5
Equate the number of electrons for both half reactions
Equate the 2 half-reactions
Cancel all substances that appear on both sides
Disproportionation Reaction
2H2O2  2H2O + O2
O.S. of O in H2O2 = -1
O.S. of O in H2O = -2
O.S. of O in O2 = 0
O.S. of O increases from -1 in H2O2 to 0 in O2 so O is OXIDISED
O.S. of O decreases from -1 in H2O2 to -2 in H2O so O is REDUCED
O has been both oxidized and reduced, so it underwent a disproportionation reaction.