Nuggets: Metals

Metals Nuggets 

last updated 23rd Jan 2014

Reactivity Series

Metals have different reactivity and you need to be familiar with the “ranking” of metals according to their ability to react. We call this the reactivity series of metals.
From the most reactive to the least reactive metals, the sequence is 
  1. Potassium, 
  2. Calcium,
  3. Sodium, 
  4. Magnesium, 
  5. Aluminium, 
  6. Carbon, (used to check the reduction of metal oxides)
  7. Zinc, 
  8. Iron, 
  9. Tin, 
  10. Lead, 
  11. Hydrogen, (used to check the displacement of H+ ions in acids)
  12. Copper, 
  13. Mercury
  14. Silver, 
  15. Gold, 
  16. Platinum
Carbon and Hydrogen are not metals but are placed in the series as a form of reference. With Carbon above zinc, it serves as a reminder that carbon can reduce any metal oxides found below Zinc. Carbon cannot reduce any metal oxides above zinc.
Because of their strong reactivity, the reactive metals are commonly found in the form of a compound.  
It is not easy to find a reactive metal existing as an element. Reactive metals tends to lose electrons easily to form ions/compounds. 
Whichever of the 2 metals which becomes atoms, is the least stable and vice versa.
The opposite is true for non-reactive metals. They are common found existing in the form of elements as they are relatively stable and do not tend to lose electrons easily. 
Metal Reactions

You need to be familiar with a few metal reactions.
(a) Metal + Oxygen,
(b) Metals + water,
(c) Metals + Steam,
(d) Metal + Acids,  
(e) Metals + solutions of salts, 
(f) Metal Oxides + Carbon/hydrogen,
Different metals react with oxygen to give different colored flames. Potassium (lilac), sodium (yellow), calcium (red), magnesium (white).
Aluminium does not react with oxygen because it is naturally coated with a layer of oxide. When the layer is removed, a white oxide layer can be seen forming and the metal gets hot.
Zinc and iron fillings glow when heated with air but they do not catch fire.

Metal Extractions
Iron occurs naturally as iron oxide (Fe2O3 hematite) or magnetite (Fe3O4), Zinc occurs naturally as Zinc Sulfide (ZnS)
Metals can be extracted from its compound by various methods (a) electrolysis, (b) reduction, (c) displacement
Electrolysis is used to extract very reaction metals like K, Ca, Na, Mg, Al

Other metals can be extracted by reduction using carbon, hydrogen, carbon monoxide.
Zinc can only be reduced by carbon. It cannot be reduced by hydrogen.
Iron and Blast Furnace
Haematite, coke and limestone added at the top of blast furnace.
Carbon reacts with oxygen to produce carbon dioxide and a lot of heat.
Carbon dioxide further reacts with carbon to produce carbon monoxide.
Carbon monoxide reduces the hematite.
Calcium carbonate undergoes thermal decomposition and forms calcium oxide.
Impurities in the hematite such as sand will react with calcium oxide to form calcium silicate. This is a type of base and acid reaction.
Calcium silicate forms a layer over the molten iron.
Slag is used in road making and as “slag cement” – a final ground slag which can be used in cement, often mixed with Portland cement. 
5 Steps
  1. C + O2 —> CO2 + Heat
  2. CO2 + C —> 2CO
  3. Fe2O3 + 3CO —> 2Fe + 3CO2
  4. CaCO3 —> CaO + CO2
  5. SiO2 + CaO —> CaSiO3 (slag or calcium silicate) 
Steel is an alloy mainly consisting of iron and carbon. The carbon content can vary to give carbon steel with different physical properties.
Stainless steel can be made by adding other metals into the mixture of iron and carbon.
Different Types of Steel
  • cast iron
  • wrought iron
  • mild steel
  • high carbon steel
  • stainless steel
  • titanium steel
  • manganese steel
Carbon Steel % of Carbon Properties Uses
mild steel  less than 0.2% malleable, ductile car bodies, machinery, cables,nails, chains
medium steel between 0.2% – 0.6% tougher and harder than mild steel steel girders, rails
high carbon steel 0.6% – 1.5% tough and hard but brittle cutting tools
Alloy Steel Metals Added Properties Uses
stainless steel chromium and nickel corrosion resistance and acid resistant cutlery,surgicalinstruments
Sacrificial Coating/Protection
A metal is coated over another metal to prevent the underlying metal from corrosion. E.g. metal A coated over metal B and metal B will not corrode. The metal coating must be more reactive than the underlying metal.
e.g. zinc coated over iron : since is zinc is more reactive, zinc corrodes and iron remains unreacted.
Other methods of corrosion prevention
Coat a layer of oil or grease, paint or plastic over iron. This prevents air and water from reacting with the iron.
Coat a layer of less reactive metal over iron. e.g Tin plating
Underground sewage steel pipes are protected from rusting by attaching blocks metal with higher reactivity (e.g. magnesium) to them. The magnesium protects the pipes from corroding as magnesium is more reactive.

Guide: Collection, Drying and Identification of Gases

Guide: Collection, Drying and Identification of Gases 

updated 1 Jan 2013
You will deal with several gases in your chemistry questions. Therefore you need to be familiar with these gases and some of their properties. You need to know how to identify these gases and collect these gases in a dry condition. Most of the time these gases are products of a reaction but they are moist in nature. 
Gas Chemical Formula
Hydrogen H2
Oxygen O2
Nitrogen N2
Ammonia NH3
Carbon Dioxide CO2
Sulphur Dioxide SO2
Chlorine Cl2
Hydrogen Chloride HCl
Gas Colour Odour Tests
H2 colourless odourless extinguishes a lighted splinter with a pop sound
O2 colourless odourless rekindles a glowing splint
CO2 colourless odourless turns limewater chalking upon bubbling (formation of white precipitate)
NH3 colourless pungent smell turns moist red litmus paper blue
Cl2 greenish-Yellow choking smell turns moist blue litmus paper red and then bleaches it
SO2 colourless choking smell turns moist blue litmus paper redturns acidified potassium dichromate solution from orange to green
HCl white misty fumes unpleasant, acrid odour turns moist blue litmus paper red
There are 4 methods of gas collection.
(a) Downward delivery
(b) Upward delivery
(c) Water displacement
(d) Gas Syringe
Limitations and advantages
Each of these methods have their limitations and advantages.
We use downward delivery to collect gases that are heavier than air. (e.g. CO2 , NO2 , SO2 , HCl, Cl2, Br2)
We used upward delivery to collect gases that are lighter than air (e.g. H2, NH3)
We use water displacement when the gases are not easily soluble in water. (H2, N2, O2, CO2)
The gas syringe is useful for us to determine the actual volume collected.
Summary Table for Gas Solubility in Water and Density 
Gas Solubility in Water Density in comparison with air
ammonia soluble less dense
chlorine fairly soluble denser
hydrogen insoluble less dense
nitrogen insoluble less dense
oxygen insoluble denser
carbon dioxide insoluble denser
hydrogen chloride soluble denser
sulfur dioxide soluble denser


The gaseous products may often by moist. There are methods to pass these gases through certain substances (drying agents) to remove the moisture so as to obtain dry gases.
These are some possible drying agents.
1. Concentrated sulfuric acid H2SO4
2. Quick lime (CaO: Calcium oxide)
3. Fused CaCl2(s)
However, the drying agents are not suitable for drying all types of gases. We have to be careful when dealing with gases with acidic nature (HCl) and gases with alkaline nature (NH3)
Summary Table for Drying Agents (Gas)
Gas CaCl2 H2SO4 CaO Reason
CO2 Yes Yes No Acidic gas, reacts with CaO (base)
O2 Yes Yes Yes  
Cl2 Yes Yes Yes  
NH3 No No Yes It reacts with CaCl2, alkaline gas, reacts with H2SO4
H2 Yes Yes Yes  
SO2 Yes Yes No Acidic gas, reacts with CaO (base)


Q&A: Metal Properties

Q&A: Metal Properties 

1. List the general physical properties of metals
2. Define ductile
3. Define malleable
4. Describe the metallic bonding. Illustrate with a diagram.
5.  Explain the high melting and boiling point of metals in terms of its structure.
6. Explain the ductile and malleable properties of metals in terms of their structure. Draw a diagram to illustrate.
7. Explain the good electrical conductivity of metals in terms of their structure.
8. Explain the good thermal conductivity of metals in terms of their structure.
9. Explain the high density of metals in terms of their structure.
1.  These are general properties. It applies for most metals with certain exceptions.
  • high melting and boiling point
  • high density
  • ductile
  • malleable
  • shiny appearance
  • good conductor of heat
  • good conductor of electricity
2. It can be stretched into wires.
3. It can be bent and hammered into sheets.
4. Metallic bonding is the strong forces of attraction between the positive metal ions and the sea of delocalised electrons.These delocalised electrons comes from the valence electrons of the atoms.
5. The metallic bonds in metals are strong hence metals have high melting and boiling points.
6. The layers of atoms are able to slide pass each other easily without breaking the metallic bond.
7. The delocalised electrons are able to move freely to conduct electricity.
8.  The closely packed structure in metals allow heat energy to be easily transferred from one atoms to another through vibration. The sea of delocalised electrons further enhance the heat conduction process through vibration.
9. The atoms in atoms are closely packed resulting in high density.

List: Endothermic/Exothermic Reactions

Examples of Exothermic Reactions: (process that releases energy)

* all reactions that give out heat energy
* reactions that forms bond
* commonly associated with temperature increase(surrounding)

– Combustion (burning)
– Neutralisation
– Haber Process
– rusting of iron
– forming on pairs
– atoms combining to form gaseous molecules
– mixing water with anhydrous salts (anhydrous salts to hydrous salts)
– mixing water with calcium chloride
– condensation
– deposition
– forming of ice (losing thermal energy)
– respiration (burning of food in body)
– double displacement (e.g. precipitation)
– synthesis (2 or more simpler substances forming more complex compounds e.g. Haber Process, except photosynthesis)
– single displacement (displacement of metals or halogens)

Examples of Endothermic Reactions:(process that absorbs energy)

* all reactions that absorbs thermal energy
* all reactions that breaks bonds
* commonly associated with temperature decrease (surrounding)
– Thermal decomposition
– Haber Process (reverse reaction)
– dissolving ammonium and potassium salts (Ammonium Chloride dissolved in water)
– separating ion pairs
– melting solid salts
– melting ice cubes
– evaporation of water
– dissolving hydrated salts in water
– atomisation (breaking of molecules into atoms involves absorption of energy) – Alevel (gaseous atoms)
– photosynthesis

Nuggets: Alkenes

last updated 3rd Oct

  1. Physical Chemical of Alkenes
  2. (i) Insoluble in water
  3. (ii) Low boiling point
  4. General formula for alkanes: CnH2n
  5. Name Formula State
    Ethene C2H4 gas
    Propene C3H6 gas
    Butene C4H8 gas
    Pentene C5H10 liquid
  6. Chemical Properties of Alkenes
  7. The functional group of alkenes is the carbon-carbon double bond. (C = C). There is only one pair of double bonds in alkenes.
  8. (i) Alkenes are more reactive than alkanes because of the presence of the double bond.
  9. (ii) Complete combustion of alkenes produces carbon dioxide and water.
  10. Alkenes produce more soot than alkanes during combustion because of the higher percentage of carbon compared with alkanes.
  11. (iii) Polymerisation. Alkenes can undergo addition polymerization at high pressure (1000 atm) and high temperature (200oC). Thousands of alkene molecules combine to form polymers (e.g. polyethene)
  12. Polyethene is flexible and difficult to break. It is used in production of cling film, plastic bags.
  13. (iv) Addition Reactions. The double bond in alkenes can be broken and other molecules can add on to the carbon atoms to form new products. This is called addition reaction. The product formed is saturated/polyunsaturated (e.g. no more double bonds).
  14. Examples of addition reactions are hydrogenationbrominationhydration.
  15. Hydrogenation is the addition of hydrogen to alkenes. (reaction takes place at 200oC using nickel as catalyst) e.g. C2H4 + H2 –> C2H6 : ethene to ethane.
  16. Hydrogenation is used to convert vegetable oils (unsaturated) into margarine.
  17. Vegetable oil turned into margarine, from a liquid to a solid state, because of the addition of hydrogen which increases the molecular mass. Margarine has larger melting point than vegetable oil.
  18. Unsaturated oil is healthier for consumption. Saturated (trans fat) fat are unhealthy for consumption.
  19. Bromination is the addition of bromine to alkenes. When ethene is passed through liquid bromine, decolorization occurs. (C2H4 + Br2 –> C2H4Br2) The product is 1,2-dibromoethane, a saturated compound.
  20. An organic compound is said to be saturated if it only contains single bonds or when the double bond is broken to form single bond.
  21. Hydration is the addition of steam to alkenes at temperature 300oC, pressure 60atm, with phosphoric acid as catalyst.
  22. Manufacturing of Alkenes by Cracking
  23. Alkenes can be obtained by cracking (break down) of petroleum fractions. 
  24. Cracking is the process of breaking larger hydrocarbons into smaller molecules. 
  25. The products of cracking are smaller alkane and alkene molecules.
  26. Alkenes can also be prepared by thermal cracking of alcohols. Ethane is produced from ethanol: C2H5OH –> C2H4 + H2O
  27. Preparation of Ethene by dehydration of alcohols
  28. Ethanol is heated to 180oC with excess concentrated sulphuric acid. Sulphur dioxide is removed from ethene by passing it through sodium hydroxide solution and ethene is collected over water. 
  29. Manufacturing of Alkenes by Cracking
  30. Cracking is the process in which larger hydrocarbon molecules are broken down into smaller molecules.
  31. When a large alkane molecule is cracked, a mixture of smaller alkanes, alkenes and hydrogen gas could be formed. 
  32. When a large alkane molecule is cracked, sometimes we get a mixture of alkenes and alkane without the hydrogen gas.
  33. E.g. C13H28 –> C2H4 + C3H6 + C8H18 
  34. E.g. C13H28 –> 2C2H4 + C9H20
  35. E.g C13H28 –> 2C2H4 + C3H6 + C6H12 + H
  36. Cracking happens at a temperature of 600oc with catalyst Silicon Dioxide or Aluminium Oxide
  37. The cracking process obtains petrol, alkenes, hydrogen

Physical Changes Vs Chemical Changes

Chemical Reactions Nuggets

Chemical Reactions Nuggets 
last updated 19th Jan 2014

 Physical and Chemical Changes
  1. Physical reactions/changes doesn’t form new substances and such changes are reversible. (e.g. melting ice, boiling water)
  2. In physical reactions, little or no heat/light is taken in or given out.
  3. Chemical reactions involves chemical changes which forms new substances. (e.g. burning of paper, photosynthesis)
  4. In general, these new substances are different in color, texture when compared to the reactants. Their states(solid/liquid/gas) may also be different. **(do note that in some cases, chemical reactions don’t involve any colour change. for e.g. methane burning in oxygen gas to form carbon dioxide and water vapour. In this reaction, all the reactants and products are colourless. CH4(g) + 2O2(g) —> CO2(g) + 2H2O(g))
  5. In chemical changes, the products usually don’t resemble the reactions. (refer to point 4 for exceptional cases)
  6. We use a word equation or a chemical equation to represent chemical reactions. 
  7. Carbon + Oxygen –> Carbon Dioxide (Word Equation)
  8. C + O2 –> CO2 (Chemical equation)
  9. In a chemical reaction, the total mass of reactants before the reaction is equal to the total mass of the products after the reaction. 
  10. In a chemical reaction, no atoms are created or destroyed. They simply re-arrange among themselves to form new molecules. 
  11. Chemical reactions are usually irreversible but some are reversible. (refer to the list below for such examples)
  12. Chemical reactions can be started by heat, mixing, light, electric current.
  13. Some signs (or evidence) of chemical change are:
    • a gas is produced,
    • the temperature changes,
    • a substance disappears,
    • a solid is formed
    • a colour change occurs,
    • a new odour is produced.
Examples of Physical Changes
  • water turning into ice (freezing)
  • salt dissolves in water (dissolution)
  • evaporation (liquid -> gas)
  • Zinc oxide turns from white to yellow upon heating and turns white again upon cooling
  • magnetising and demagnetising a magnet
  • expansion and contraction of materials under the effect of heat
Examples of Chemical Changes that are reversible

( —><— sign refers to reversible reactions)
  • Sulphur dioxide  + Water   —><—  Sulphurous acid 
  • Ammonia + Water —><— Ammonium hydroxide
  • Iron + Steam —><—  Iron oxide + Water 
  • Calcium Carbonate —><—  Calcium Oxide + Carbon dioxide
  • Ammonium Chloride —><—  Ammonia Gas + Hydrogen Chloride Gas (this reaction is commonly mistake for a sublimation process)

Different Types/Categories of Chemical Reactions 

  1. Combination (Synthesis)
  2. Decomposition (thermal, light)
  3. Single Displacement
  4. Double Displacement (metathesis reaction)
  5. Complete Combustion and Incomplete Combustion
  6. Hydrolysis
  7. Oxidation
  8. Reduction
  9. Neutralisation (Acid-Alkali)
  10. Precipitation
  11. Electrolysis
Types of Chemical Reactions by Heat
  • combination
  • combustion
  • thermal decomposition
  • oxidation (may also happen without heat)
Explanation for the different types of Chemical Reactions


This involves 2 or more substances going through a chemical reaction to form a single new product.

  • Fe + S —> FeS (involving elements)
  • C + O2 —> CO(involving elements)
  • Fe + O2 —> Fe2O(involving elements)
  • 2Na + Cl2 —> 2NaCl (involving elements)
  • PCl3 + Cl2 —> PCl(involving compounds and elements)
In this reaction Fe + S —> FeS , we know this is a chemical change because the product looks different form the reactants. Fe is a silvery metal and S is a yellow solid. But the product is a black solid.

Complete Combustion

Combustion is also known as burning. Combustion is also known as a type combination reaction. In a combustion reaction, oxygen is combined with other reactants in the presence of heat. The end product of combustion reactions are CO2 and H2O.

  • combustion of methane
  • CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
  • burning of naphthalene
    C10H8 + 12O2 → 10CO2 + 4H2O

  • combustion of ethane
    2C2H6 + 7O2 → 4CO2 + 6H2O
Incomplete Combustion
Combustion reactions can also sometimes be incomplete. In such cases, carbon dioxide (CO2) will not be formed. Instead, carbon monoxide (CO) is formed. CO is a poisonous gas. It is also odourless which makes it hard to be detected. 
Single Displacement

Single displacement reactions has this general pattern. A + BC —> AB + C 
In single displacement reactions, one of the elements in the compound is replaced by another element.

Li(s) + NaCl(aq) > LiCl(aq) + Na(s)

F2(aq) + 2KI(aq) —> I2(aq) + 2KF(aq)

Double Displacement

The general pattern of a double displacement reaction is:

AB + CD —> AD + CB

  • 2KI(aq) + Pb(NO3)2(aq) —> PbI2(s) + 2KNO3(aq)
  • MgCl2(aq) + Ca(OH)2(aq) > Mg(OH)2(s) + CaCl2(aq)

Nuggets: Alkanes

last update 3rd Oct
  1. Organic chemistry is the study of the properties of compounds containing carbon. (carbon compounds)
  2. hydrocarbon is a compound containing carbon and hydrogen only.
  3. An organic compound contains the element carbon. (except CO, CO2, and carbonates)
  4. Organic compounds may also contain other elements like oxygen and nitrogen.
  5. saturated compound contains only single covalent bonds
  6. An unsaturated compound contains double covalent bonds.
  7. homologous series is a group of compounds with a general chemical formula, which follows a regular structural pattern in which each successive member differs from the other by a -CH2- group.
  8. Most chemical properties of organic compound (except for alkanes) are due to the presence of the function group. (e.g. alkenes have the C = C functional group)
  9. General formula for alkanes: CnH2n+2
  10. Name Formula State
    Methane CH4 gas
    Ethane C2H6 gas
    Propane C3H8 gas
    Butane C4H10 gas
    Pentane C5H12 liquid
    Hexane C6H14 liquid
    Heptane C7H16 liquid
    Octane C8H18 liquid
    Nonane C9H20 liquid
    Decane C10H22 liquid
  11. C1-4 (gas), C5-17(liquid), C18 onwards (solid: e.g bitumen)
  12. Characteristics of an alkane homologous series:
  13. (a) all members share a general chemical formula,
  14. (b) each successive member differs in its molecular formula by the addition of a -CH2- gropup,
  15. (c) its relative molecular mass differs from each group by 14,
  16. (d) all members have similar chemical properties ( (i) undergoes combustion to form CO2 and H20, (ii) undergoes substitution reaction),
  17. (e) physical properties of each member changes gradually as the number of carbon atoms increases. (as the carbon atom increases, the boiling point also increases, the compound becomes less flammable and become less volatile, the compound gets more viscous as the longer molecules gets entangled more easily)
  18. Physical Properties of Alkanes
  19. (i) insoluble in water
  20. (ii) does not conduct electricity
  21. (iii) low melting and boiling point for C1-4.
  22. Chemical Properties of Alkanes
  23. (i) All alkanes are generally unreactive because the C-C and C-H covalent bonds are strong.
  24. However they can undergo
  25. (ii) combustion: easily flammable when in gaseous and liquid states. less flammable in solid state. When solid alkane(alkanes with larger molecule size, e.g. candle wax) is burnt, it produces a smoky flame due to incomplete combustion of carbon atoms in the molecules. The smoky flame is caused by soot (carbon). This combustion is an exothermic reaction (energy released). Incomplete combustion also produces CO instead of CO2. In an incomplete combustion, the products formed are CO and H2O. In some other cases, incomplete combustion can also lead to the formation of carbon. (e.g. CH4 + O2 -> C + 2H20)
  26. (iii) (Halogenation) Substitution reactions with halogens (chlorine/bromine)
  27. Alkanes react with halogens in the presence of ultraviolet light(UV)
  28. The UV light acts as a catalyst. The UV light breaks up the Cl2 bonds.
  29. CH4 + Cl2 –> CH3Cl + HCl (hydrogen atom is replaced by chlorine atom)
  30. Substitution reaction can be repeated to form CH2Cl2, CHCl3, CCl4
  31. Isomerism (only for Pure Chem)
  32. Isomerism is where two or more organic molecules have the same molecular formula but different structural formulae. These different molecules are called isomers.
  33. Isomers have different melting and boiling points.
  34. The branched-chain or side-chain of an alkane is called alkyl groups. The formula of the alkyl groups is CnH2n+1