PeriodicityHorizontal 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 IonsElements on the left hand side (Group I, II) Elements on the right hand side (Group III, IV, V, VI, VII, VIII) @ Lose their valence electrons (tendency to lose

electrons) @ Form positive ions (Cations) @ Charge on positive ions correspond to group

number

@ 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

M etallic Character Metals Non-metals

- Left hand side of the Periodic Table - Groups I and II - Few valence electrons- Tendency to lose electrons - Form cations

- Right hand side of the Periodic Table - Groups III, IV, V, VI, VII, VIII/ 0 - Many valence electrons- Tendency to gain electrons- 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)

M elting 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.

E lectronegativity Tendency to attract electrons to itself

Atomic Radius - Measured by the distance between the nuclei of 2 touching atoms / 2

Type of Bond RadiiMetallic Bond Metallic radii Covalent bonds (non-metals) Covalent radii No Bond (e.g. Noble gas) Van der Waals radii

Down the Group Across the Period Atomic radius increases down the group

Reason: Down the group, an extra electron shell is added, increasing the size of atoms.

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.

Atomic radius decreases across the period

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 nuclear force of attraction, stronger bonds, so it’s smaller

Same amount of screening across period

Ionic Radius Positive Ions (Cations) Negative Ions (Anions)

Cations have a SMALLER 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.

Negative ions have a BIGGER ionic radius than their atoms

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

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

Decrease in ionic size across the period from Group V to VII

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.E1. 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 Across a Period General trend is for I.E. to decrease down a group General trend is for I.E. to increase

across a periodReason: The proton number greatly increases down the group, greatly increasing nuclear charge.

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

Reason: The screening effect of valence electrons is the same across the period.

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 observedLi, 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 NaHissing sound is heardAlkaline 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 @ 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 ColourFluorine Pale Yellow Gas Chlorine Greenish yellow gasBromine Reddish brown liquid Iodine Purple black solid *purple gas – can sublime!

U ses Halogen Application

Fluorine @ 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

Chlorine @ Bleach, disinfectants, antiseptics@ Polymer PVC@ Chlorofluorocarbons (CFCs), used in refrigerators and aircon units

Bromine @ 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 Iodine @ 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 compoundsCompound Colour Iron (II) Sulphate Fe2+ GreenIron (III) Chloride Fe3+ Brownish-yellowPotassium Manganate (VII) KMnO4 PurplePotassium Dichromate (VI) K2Cr2O7 OrangeCopper (I) oxide Cu+ Reddish Brown Copper (II) oxide Cu2+ BlackCopper 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@ 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 reactivityNoble Gas Property ApplicationHelium Very light, not combustible Filling balloons and airshipsNeon Glows brightly when a current passes

through the gasUsed for advertising signsNeon lights (disco!)

Argon Unreactive Provides an inert atmosphere inside light bulbsKrypton and Xenon Used in fluorescent bulbs, flash bulbs and lasersRadon Radioactive Treat cancer and detect leaks in gas pipelines

Alkali Metals Halogens Noble Gases Transition Metals Valence Electrons

1 7 2 (Helium), usually 8

Varies

Lose/ Gain Lose Gain Neither Varies 1 element may form ions with different oxidation states (lose/ gain different number of electrons)

Charge on Ion

+1 -1 Neutral

Metal/Non- Metal Non-metal Non-metal Metal

metalState of Matter

Solids Gas > Solid down the group Gases

Atomic Structure

Diatomic Monoatomic

Melting Point/ Boiling Point

Low melting pointsGenerally decrease

down the group Atomic radii of atoms

increase due to the addition of shell, thus the forces of attraction are weaker

Low melting points and boiling points Increase down the group Van Der Waals forces of

attraction are stronger down the group

Very low melting and boiling points Increase down

the group Van Der Waals’

forces of attraction increase; but still quite low as they are single atoms

High melting points

Density 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

Density increases down the groupAtomic mass

increases down the group

High density

Colour Silvery Elements become darker down the group

Colourless gases Varies (Coloured)Characteristic colours

Electrical Conductivity

No

Reactivity 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 groupmore shells, increasing

the shielding effect Harder for the halogen

to gain electrons

DOES NOT REACT INERT

Not as reactive as alkali metals or halogens

ReactivityP hysical properties of metals

Atoms in a metal are held together by strong metallic bondsIn a metal lattice, a “sea” of mobile electrons surrounds a lattice of positive metal ions

Attribute Metals Reason Appearance Shiny surfaceMelting & Boiling Point Usually high Metallic bonding is strong, hence a lot of energy is required to

break the bonds and separate the atoms Physical state Solid (except mercury) Density Usually highElectrical conductivity Good conductor Presence of mobile delocalized electrons through the metal

that carry the electrical current Thermal conductivity Good conductor 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.

Strength Strong Ease of shaping Ductile and malleable 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.

C hemical properties Metals Non-metals React with oxygen to form basic oxides or amphoteric oxides React with oxygen to form acidic oxides or

neutral oxidesForm positive ions by losing electrons Form negative ions by gaining electrons Form ionic chlorides and oxides Form covalent chlorides and oxides Most metals react with dilute acids to give hydrogen and a salt No reaction with dilute acids

R eactivity 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 H2O HClPotassiumSodiumCalcium

Can react with cold waterK: vigorously, heat, fire, explosionNa: vigorous, may catch fire/explodeCa: readily, bubbles producedEg: Ca + 2H2O Ca(OH)2 + H2

Can react with cold acidsK, Na: explodes (dangerous!)Ca: very fast, bubbles producedEg: 2K + 2HCl 2KCl + H2

Magnseium*AluminiumZincIron

Needs to react with steamMg: vigorously (very slow in cold water)Zn: less vigorouslyWill NOT get hydroxide! Only metal oxides

Needs to react with warm acidMg: fastZn: moderately fastFe: slowly

TinLead

(steam is a dry gas) Eg: Zn + H2O ZnO + H2

Pb: very slowly with warm HCl

CopperMercurySilverGold

No reactionLess reactive than hydrogen!

A lloys @ 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@ 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)

Element Reactivity Symbol ValencyPlease Potassium @ 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

@ Reacts vigorously (fast) with chemicals

@ Readily gives up electrons in reactions to form positive ions

@ Corrodes easily

K 1Stop Sodium Na 1Calling Calcium Ca 2Me Magnesium Mg 2

A Aluminium 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

Al 3

Crazy Carbon C 4Zebra Zinc Can be reduced by carbon Zn 2I Iron Fe 2, 3Truly Tin Sn 2, 4Love Lead Pb 2, 4Hamburgers Hydrogen 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

H 1

produce hydrogen gas Containing Copper @ 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 1, 2M Mercury Hg S Silver Ag 1G Gold Au 1, 3

U ses @ Predict chemical reactions that can occur@ Helps in extraction of metals and prevention of rusting

Displacement ReactionsA metal can displace another metal from a compound Displacement of Metals from Solutions Displacement of Metals from Metal Oxides A more reactive metal will displace the ions of a less reactive metal in the reactivity series from solution

@ 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+ + 2e-

Cu 2+ (aq) + 2e- Cu (s)

A more reactive metal will take the oxygen from the oxide of a less reactive metal

@ 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) + 2e-

Cu2+ + 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 Reaction of Metal Oxides with Hydrogen 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)

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) PotassiumSodiumCalciumMagnesium

Oxides that can be reduced by carbon (below carbon) Zinc IronLeadCopperSilver

@ This method is often used to extract metals from their ores

Oxides that cannot be reduced by hydrogen (can only be reduced by electrolysis) PotassiumSodiumCalciumMagnesium Zinc

Oxides that can be reduced by hydrogen IronLeadCopperSilver

@ 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.

E lectrolysis @ 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) + SO4

2- (aq) Mg2+ (aq) + SO42- (aq) + Cu (s)

@ Take out the spectators: Mg (s) + Cu 2+ (aq) Mg 2+ (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 Observation Reason PotassiumSodium

Unaffected by heat Does not decompose

Too reactiveStable (eg Na2O3)

CalciumMagnesiumAluminiumZincIron (II)LeadCopper (II)

Decomposes into metal oxide and carbon dioxide CuCO3 CuO + CO2

Silver 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 ApplicationMethod Application Remarks Painting Large objects e.g. vehicles, bridges Metal will rust if the paint is scratchedOiling/ Greasing Machinery and tools Also helps in lubrication Plastic Layer Small iron/ steel objects e.g. wires Metal will rust if plastic layer is scratched Zinc-plating Buckets, dustbins, kitchen sinks Metal doesn’t rust even if zinc layer is damaged Tin-plating Food cans Metal rusts if coating is scratched Chrome-plating Taps, kettles, bicycles Also makes it shinyMore reactive metal block Underground pipes, ships More reactive metal will corrode instead of ironStainless Steel Cutlery, surgical instruments Form chromium oxide layer on the metal

RedoxA reaction in which one substance is oxidized while another is reduced at the same time is called a Redox reaction.

Oxidation ReductionOXYGEN

Gain of Oxygen by a substancee.g. 2Ca (s) + O2 (g) 2CaO (s) Calcium gained oxygen, thus, it has been oxidised

Loss of Oxygen by a substance e.g. Zn (s) + CuO (s) ZnO (s) + Cu (s) Copper loses oxygen, thus it has been reduced

HYDROGENLoss of hydrogen by a substance e.g. 2NH3 (g) + 3CuO (s) N2 (g) + 3Cu (s) + 3H2O (l) Ammonia has lost hydrogen, thus it has been oxidised

Gain of hydrogen by a substance e.g H2S (g) + Cl2 (g) 2HCl (g) + S (s) Chlorine gains hydrogen, thus it is reduced

ELECTRONSLoss of electrons from a substance Put in the charges (even for compounds) e.g. 2 Na (s) + Cl2 (g) 2Na+ Cl- (s) Loss of electrons = positive charge Sodium loses electrons, thus it has been oxidised

Gain of electrons by a substance e.g. 2 Na (s) + Cl2 (g) 2Na+ Cl- (s)Gain of electrons = negative charge Chlorine gains electrons, thus it is reduced

*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. Write the before and after of the substance 2. Where should electrons be added such that the reaction makes sense? 3. Must include STATE SYMBOLS (molten/melted = liquid (l), solution = aqueous (aq)

Na (s) Na+ (s) + _______To equate the reaction, the ___ must be a e-

Sodium 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 Causes another substance to get oxidized Gets reducedRA Causes another substance to get 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 GreenK2Cr2O7 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 ColourlessKMNO4 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. Calculate the oxidation states of each element before and after the reaction2. Identify which element undergoes oxidation (reducing agent) and which undergoes reduction (oxidising agent) 3. Write the half equations for oxidation and reduction 4. 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 H2O to make the other side have x oxygen atoms too)

@ Then, add H+ to balance out the extra hydrogen 5. 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

6. Equate the number of electrons for both half reactions 7. Equate the 2 half-reactions 8. Cancel all substances that appear on both sides

Disproportionation Reaction

2H2O2 2H2O + O2

O.S. of O in H2O2 = -1O.S. of O in H2O = -2O.S. of O in O2 = 0

O.S. of O increases from -1 in H2O2 to 0 in O2 so O is OXIDISEDO.S. of O decreases from -1 in H2O2 to -2 in H2O so O is REDUCEDO has been both oxidized and reduced, so it underwent a disproportionation reaction.