+kimia+

Other processes and reactions
Galvanising	This means to coat iron or steel with a layer of zinc to stop it rusting (more details on Metal Reactivity page)
Haber Process	The synthesis of ammonia by combining nitrogen and hydrogen using high temperature, pressure and an iron catalyst. (all the details)
Contact Process	Part of the manufacture of sulphuric acid from: (all the details) sulphur ==> sulphur dioxide ==V catalyst==> sulphur trioxide* ==> sulphuric acid. S(s) + O2 (g) ==> SO2 (g), * 2SO2 (g) + O2 (g) ==> 2SO3 (g), followed by, indirectly, SO3 + H2O ==> H2SO4
Double decomposition	Double decomposition is chemical reaction that takes place between two compounds, in which the first part of one compound combines with the second part of another compound. The bits left over combine to form the second compound. One of the compounds is usually insoluble. e.g. if you mix solutions of sodium chromate with lead nitrate you get a yellow precipitate of lead chromate and sodium nitrate is left in solution. sodium chromate + lead nitrate ==> lead chromate + sodium nitrate Na2CrO4(aq) + Pb(NO3)2(aq) ==> PbCrO4(s) + 2NaNO3(aq) This is the yellow pigment Van Gogh used in his paintings and you see it as the road markings you don't park on!
Catalytic Conversion (car exhaust)	One way of reducing pollutants from a car exhaust is to use transition metal catalysts (Pt+Rh set on a heat resistant base = the catalytic converter). A platinum/rhodium catalyst converts toxic carbon monoxide and nitrogen monoxide gases into harmless carbon dioxide and nitrogen gases. 2CO (g) + 2NO (g) == Pt/Rh ==> 2CO2 (g) + N2 (g)
Esterification	Combining an organic carboxylic acid with an alcohol, produces an pleasant smelling ester (which are used in perfumes and flavourings) and water. e.g. ethanoic acid + ethanol ethyl ethanoate + water + + H2O Its an equilibrium, 2/3 rds conversion, and the reaction is catalysed by a few drops of concentrated sulphuric acid.
Rusting	Iron (or steel) corrodes more quickly than most other transition metals and readily does so in the presence of both oxygen (in air) and water to form an iron oxide. This means rusting is an oxidation reaction. iron + oxygen + water ==> hydrated iron(III) oxide 4Fe(s) + 3O2(g) + xH2O(l) ==> 2Fe2O3.xH2O(s) i.e. rust is an orange-brown solid of hydrated iron(III) oxide formed from the reaction with oxygen and water (the equation is not meant to be balanced and the amount of water x is variable, from dry to soggy!). For more details of the chemistry of rusting and its prevention go to the corrosion section on the Metal Reactivity Series page.
Substitution	A substitution reaction is where one part of a molecule is replaced by something else. e.g. when chlorine reacts with methane, a hydrogen atom in methane is replaced by a chlorine atom from the chlorine molecule. methane + chlorine ==> chloromethane + hydrogen chloride CH4 + Cl2 ==> CH3Cl + HCl
Addition	An addition reaction is when one molecule adds to another molecule resulting in a single product e.g. ethene + bromine ==> 1,2-dibromoethane C2H4 + Br2 ==> C2H4Br2

Oxidation and Reduction +Chemistry+

For a summary of the metals chemical reactions with air/oxygen, acids and oxides/salts (displacement), including word equations and balanced symbol equations, all in the context of the reactivity series just click on its name from this alphabetical order list ... aluminium .. caesium .. calcium .. copper .. francium .. gold .. iron .. lead .. lithium .. magnesium .. platinum .. potassium .. rubidium .. silver .. sodium .. tin .. zinc (but the notes are in reactivity order) and other sub-sections on this page: METAL CORROSION- RUSTING * DISPLACEMENT REACTIONS and OXIDATION - REDUCTION explained

OXIDATION and REDUCTION - REDOX REACTIONS
OXIDATION - definition and examples	REDUCTION - definition and examples
(a) The gain or addition of oxygen by an atom, molecule or ion e.g. ... (1) S ==> SO2 [burning sulphur - oxidised] (2) CH4 ==> CO2 + H2O [burning methane to water and carbon dioxide, C and H gain O] (3) NO ==> NO2 [nitrogen monoxide oxidised to nitrogen dioxide] (4) SO2 ==> SO3 [oxidising the sulphur dioxide to sulphur trioxide in the Contact Process for making sulphuric acid]	(b) The loss or removal of oxygen from a compound etc. e.g. ... (1) CuO ==> Cu [loss of oxygen from copper(II) oxide to form copper atoms] (2) Fe2O3 ==> Fe [iron(III) oxide reduced to iron in blast furnace] (3) NO ==> N2 [nitrogen monoxide reduced to nitrogen, catalytic converter in car exhaust] (4) SO3 ==> SO2 [sulphur trioxide reduced to sulphur dioxide]
(c) The loss or removal of electrons from an atom, ion or molecule e.g. (1) Fe ==> Fe2+ + 2e- [iron atom loses 2 electrons to form the iron(II) ion, start of rusting chemistry] (2) Fe2+ ==> Fe3+ + e- [the iron(II) ion loses 1 electron to form the iron(III) ion] (3) 2Cl- ==> Cl2 + 2e- [the loss of electrons by chloride ions to form chlorine molecules]	(d) The gain or addition of electrons by an atom, ion or molecule e.g. ... (1) Cu2+ + 2e- ==> Cu [the copper(II) ion gains 2 electrons to form neutral copper atoms, electroplating or displacement reaction) (2) Fe3+ + e- ==> Fe2+ [the iron(III) ion gains an electron and is reduced to the iron(II) ion] (3) 2H+ + 2e- ==> H2 [hydrogen ions gain electrons to form neutral hydrogen molecules, electrolysis of acids or metal-acid reaction]
(e) An oxidising agent is the species that gives the oxygen or removes the electrons	(f) A reducing agent is the species that removes the oxygen or acts as the electron donor
REDOX REACTIONS - in a reaction overall, oxidation and reduction must go together
(g) Redox reaction analysis based on the oxygen definitions
(1) copper(II) oxide + hydrogen ==> copper + water CuO(s) + H2(g) ==> Cu(s) + H2O(g) copper oxide reduced to copper, hydrogen is oxidised to water hydrogen is the reducing agent (removes O from CuO) copper oxide is the oxidising agent (donates O to hydrogen) (2) iron(III) oxide + carbon monoxide ==> iron + carbon dioxide Fe2O3(s) + 3CO(g) ==> 2Fe(l) + 3CO2(g) the iron(III) oxide is reduced to iron, the carbon monoxide is oxidised to carbon dioxide CO is the reducing agent (O remover from Fe2O3) the Fe2O3 is the oxidising agent (O donator to CO)] (3) nitrogen monoxide + carbon monoxide ==> nitrogen + carbon dioxide 2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g) nitrogen monoxide is reduced to nitrogen carbon monoxide is oxidised to carbon dioxide CO is the reducing agent and NO is the oxidising agent (4) iron(III) oxide + aluminium ==> aluminium oxide + iron (the Thermit reaction) Fe2O3(s) + 2Al(s) ==> Al2O3(s) + 2Fe(s) iron(III) oxide is reduced and is the oxidising agent aluminium is oxidised and is the reducing agent
(h) Redox reaction analysis based on the electron definitions
(1) magnesium + iron(II) sulphate ==> magnesium sulphate + iron Mg(s) + FeSO4(aq) ==> MgSO4(aq) + Fe(s) this is the 'ordinary molecular' equation for a typical metal displacement reaction, but this does not really show what happens in terms of atoms, ions and electrons, so we use ionic equations like the one shown below. The sulphate ion SO42-(aq) is called a spectator ion, because it doesn't change in the reaction and can be omitted from the ionic equation. No electrons show up in the full equations because electrons lost by x = electrons gained by y!! magnesium + iron(II) ion ==> magnesium ion + iron Mg(s) + Fe2+(aq) ==> Mg2+(aq) + Fe(s) the magnesium atom loses 2 electrons (oxidation) to form the magnesium ion, the iron(II) ion gains 2 electrons (reduced) to form iron atoms. Mg is the reducing agent (electron donor) and the Fe2+ is the oxidising agent (electron remover or acceptor) Displacement reactions involving metals and metal ions are electron transfer reactions. (2) zinc + hydrochloric acid ==> zinc chloride + hydrogen Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g) the chloride ion Cl- is the spectator ion zinc + hydrogen ion ==> zinc ion + hydrogen Zn(s) + 2H+(aq) ==> Zn2+(aq) + H2(g) Zinc atoms are oxidised to zinc ions by electron loss, so zinc is the reducing agent (electron donor) hydrogen ions are the oxidising agent (gaining the electrons) and are reduced to form hydrogen molecules (3) copper + silver nitrate ==> silver + copper(II) nitrate Cu(s) + 2AgNO3(aq) ==> 2Ag + Cu(NO3)2(aq) the nitrate ion NO3- is the spectator ion copper + silver ion ==> silver + copper(II) ion Cu(s) + 2Ag+(aq) ==> 2Ag(s) + Cu2+(aq) copper atoms are oxidised by the silver ion by electron loss electrons are transferred from the copper atoms to the silver ions, which are reduced the silver ions are the oxidising agent and the copper atoms are the reducing agent (4) iron(II) chloride + chlorine ==> iron(III) chloride (5) halogen (more reactive) + halide salt (of less reactive halogen) ==> halide salt (of more reactive halogen) + halogen (less reactive) X2(aq) + 2KY(aq) ==> 2KX(aq) + Y2(aq) X2(aq) + 2Y-(aq) ==> 2X-(aq) + Y2(aq) where halogen X is more reactive than halogen Y, F > Cl > Br > I X is the oxidising agent (electron acceptor) KY is the reducing agent (electron donor) See GCSE Group 7 The Halogens - displacement reaction notes (6) Electrode reactions in electrolysis are electron transfer redox changes at the negative cathode positive ions are attracted: metal ions are reduced to the metal by electron gain: Mn+ + ne- ==> M n = the numerical charge of the ion and the number of electrons transferred or 2H+(aq) + 2e- ==> H2(g) (for the discharge of hydrogen) at the positive anode negative ions are attracted: negative non-metal ions are oxidised by electron loss e.g. for oxide ions: 2O2- - 4e- ==> O2 or 2O2- ==> O2 + 4e- for hydroxide ion: 4OH- - 4e- ==> O2 + 2H2O or 4OH- ==> O2 + 2H2O + 4e- for halide ions (X = F, Cl, Br, I): 2X- - 2e- ==> X2 or 2X- ==> X2 + 2e-

2009 SPM EXAM TIPS – CHEMISTRY +Soalan spOt+

2009 SPM EXAM TIPS – CHEMISTRY

PAPER 2

SECTION A

1. Empirical formula. (precaution, draw apparatus, observation, calculation).
2. Periodic table.
3. Carbon compound. (alkane, alkane, alcohol.)
4. Salt
5. Redox and electrolysis. (define term)
6. Soap and detergent (how to make soap, adv, disadvantage.)
7. Contact & Haber process.
8. Thermochemistry

SECTION B

1. How to make soluble / insoluble salt.




(Identity anion / cation.)

1. Rate of reaction. (catalyst, surface area, concentration.)
2. Bond (ionic , covalent)
3. Combination of (alloy ,glass, polymer) or




Rubber (polymer & vulcanisation)

Chemistry Study Guides

Read for FREE Online

Fundamentals of Chemistry

Units, Scientific Notation, and Significant Figures The Periodic Table Atomic Structure

Bonding

Introduction to Chemical Bonding Properties of Chemical Bonds Ionic Bonds Covalent Bonds Molecular Orbitals Review of Chemical Bonding

Stoichiometry

Introduction to Stoichiometry Stoichiometric Calculations Real World Stoichiometric Reactions Review of Stoichiometry

Gases

Gas Pressure Ideal Gases Kinetic Molecular Theory Real Gases Review of Gases

Solutions

Introduction to Chemical Solutions Composition of Solutions Colligative Properties of Solutions Solubility Review of Chemical Solutions

Acids & Bases

Introduction to Acids and Bases Fundamentals of Acids and Bases pH Calculations Buffers Titrations Review of Acids and Bases

Electrochemistry

Introduction to Electrochemistry Galvanic Cells Electrolytic Cells Thermodynamics Review of Electrochemistry

Reaction Kinetics

Introduction to Reaction Kinetics Reaction Rate Laws Reaction Mechanisms Review of Reaction Kinetics

Organic I

Organic Atomic Structure Organic Molecular Orbitals Organic Covalent Bonding

Organic II

Structure of Alkanes Conformations of Alkanes Carbocycles

Organic III

Enantiomers and Diastereomers Stereoisomers Stereochemistry

Organic IV

Introduction to Organic Chemistry IV Sn1E1 Reactions Sn2E2 Reactions Review of Organic Chemistry IV

Summary of the structure of the Periodic Table

1. Summary of the structure of the Periodic Table

1a. The basic structure of the Periodic Table

See the notes 1. to 4. in the full Periodic Table at the end of this page.

• The idea of the Periodic Table is to arrange the elements in a way that enables chemists to understand patterns in the properties of elements, but some reminders first.
• An ATOM is the smallest particle of a substance which can have its own characteristic properties, BUT atoms are built up of even more fundamental sub-atomic particles - the electron, proton and neutron and the structure of an atom ultimately determines its properties.
• An ELEMENT is a pure substance made up of only one type of atom, 92 of the elements in the Periodic Table (part of which is shown above) naturally occur, from hydrogen H (element 1) to uranium U (element 92).
• Note that each element has symbol which is a single capital letter like H or U or a capital letter + small letter e.g. cobalt Co, chlorine Cl or sodium Na.
• The majority of elements are readily divided into two types with common characteristic physical and chemical properties.
  • Most elements on the left are METALS and their typical properties are described in section 2a. e.g. elements 3 to 4 (lithium to beryllium), elements 11 to 13 (sodium to aluminium), elements 19 to 31 (potassium to gallium), elements 37 to 50 (rubidium to tin).
  • The elements on the right are NON-METALS and their typical properties are described in section 2b. e.g. elements 1 to 2 (hydrogen to helium), elements 5 to 10 (boron to neon), elements 15 to 18 (phosphorus to argon), elements 35 to 36 (bromine to krypton).
  • BUT a few elements are referred to as SEMI-METALS which have mixed metal/non-metal character and not so easy to classify, see section 2c. They occur in a diagonal band (down and L to R) e.g. silicon (₁₄Si), germanium (₃₂Ge), arsenic (₃₃As) and tellurium (₅₂Te).
• The elements are laid out in order of Atomic (proton) Number* (*see atomic structure page).
  • Originally they were laid out in order of 'relative atomic mass' (the old term was 'atomic weight').
  • This is not correct for some elements now that we know their detailed atomic structure (detailed GCSE notes) in terms of protons, neutrons and electrons, and of course, their chemical and physical properties.
  • For example: Argon (at. no. 18, electrons 2,8,8) has a relative atomic mass of 40. Potassium (at. no. 19, electrons 2,8,8,1) has a relative atomic mass of 39. Argon, in terms of its physical, chemical and electronic properties is clearly a Noble Gas in Group 8 (0). Likewise, potassium is clearly an Alkali Metal in Group 1.
• Many of the similarities and differences in the properties of elements can be explained by the electronic structure of the atoms (electron configuration = electron arrangement in shells or energy levels, so watch out the varying phrases used!).
• The idea of the Periodic Table is to arrange the elements in a way that enables chemists to understand patterns in the properties of elements.
• The main structural features of the periodic table are ...
  • to produce columns of similar elements called Groups.
    • They are usually similar chemically and physically BUT there are often important trends in physical properties and chemical reactivity up/down a group.
  • The resulting complete horizontal rows are called Periods and usually consist of a range of elements of different character.
    • There are important trends from left to right across a period e.g. the most important overall change is from metallic ==> non-metallic element character.
    • Certain 'horizontal blocks' of elements within a period, which have specific chemical features in common, may be known as a particular block or series e.g. from ₂₁Sc to ₃₀Zn are called the 1st Transition Metal Series within period 4.

Notes! Notes! Notes!++Chemistry++

SPM Physics Form 4 02 Structure of Atom I | Answer 02 Structure of Atom II | Answer 02 Structure of Atom III 02 Structure of Atom IV 05 Chemical Bond I 05 Chemical Bond II 05 Chemical Bond III 05 Chemical Bond IV 06 Electrolyte (Card) 07 Colour of Ion - Copper (Card) 07 Colour of Ion - Iron (Card) (iPaper) (iPaper) (iPaper) (iPaper) SPM Physics Form 5 Rate of Reaction I Rate of Reaction II Carbon Compound I Carbon Compound II Complete Formulae List (Highly recomended) Chemical Reaction List - Form 4 Chemical Reaction List - Form 5

( iPaper is a document format built for the Web. Virtually any document format is convertable to iPaper, which is embeddable into any website or blog. You need Flash Player to open an iPaper file. The file size may be a little too big for dial up internet connection.)

Collision theory ( chemistry)

The collision theory states that for a chemical reaction to occur, the reacting particles must:

• collide with each other
  
• possess activation energy
  
  

Activation energy is the minimum energy required by the reacting particles in order for reaction occur. The figure below shows the energy profile diagram for exothermic and endothermic reaction.

Exothermic Reaction:

Endothermic reaction:

The collision that are succesful in producing a chemical reaction are called effective collisions.

Collision theory is used to explain the factors that affect the rates of reactions. Any factor that increases the frequency of effective collision will increase the rate of reaction.

Factor affecting the rate of reaction:

• Size of particle: dcrease in particle size will increase the total surface area exposed to collision.
• Concentration: increase in concentration increases the number of perticles per unit volume. So will give more chance to effective collision occur.
• Temperature: increase in temperature increase the number of paticles possesing activation energy.
• Catalyst: will lowers the activation energy.
  

[Chemistry Form 5] Collision Theory

• The collision theory is based on the kinetic theory, which states that molecules are in constant random motion.
  
  
• In order for a reaction to occur, the reactant molecules must collide in the correct orientation, and with enough energy to form products.
  
  
• Collision between particles that produce a chemical change are called "effective collisions". This is when the reactant molecules collide with enough kinetic energy to break their original bonds and then form new bonds in the product molecules.
  
  
• Effective collision frequency is the number of effective collisions per unit time.
  
  
• The activation energy E_a, is the minimum energy the colliding reactant particles must have before any collision between them can result in a chemical reaction.

The Collision theory, proposed by Max Trautz and William Lewis in 1916 and 1918, qualitatively explains how chemical reactions occur and why reaction rates differ for different reactions.

This theory is based on the idea that reactant particles must collide for a reaction to occur, but only a certain fraction of the total collisions have the energy to connect effectively and cause the reactants to transform into products.

This is because only a portion of the molecules have enough energy and the right orientation (or "angle") at the moment of impact to break any existing bonds and form new ones. The minimal amount of energy needed for this to occur is known as activation energy.

Particles from different elements react with each other by releasing activation energy as they hit each other.

If the elements react with each other, the collision is called successful, but if the concentration of at least one of the elements is too low, there will be fewer particles for the other elements to react with and the reaction will happen much more slowly.

As temperature increases, the average kinetic energy and speed of the molecules increases but this only slightly increases the number of collisions.

The rate of the reaction increases with temperature increase because a higher fraction of the collisions overcome the activation energy.

Collision theory is closely related to chemical kinetics.

Reaction rate tends to increase with concentration -
a phenomenon explained by collision theory

Oxidation and Reduction

Oxidation and Reduction

Form 4 Chapter 4 - Periodic Table of Elements 3

Final part for this chapter:

More about Periods in the Periodic Table:
Reactions:
1) metal oxide + water --> basic solution
2) non-metal oxide + water --> acidic solution
3) oxide + acidic/basic solution will produce different results, as seen below.
- metal oxide + acid --> salt + water
- metal oxide + base --> no reaction
- non-metal oxide + acid --> no reaction
- non-metal oxide + base --> salt + water
- amphoteric + acid --> salt + water
- amphoteric + basic --> salt + water
(An example of an amphoteric substance is aluminium oxide, Al(OH)3.)

Transition metals

Characteristics:
- high b.p & m.p
- high density
- conductors
- shiny
- ductile
- malleable
- can form a coloured solution in a compound
Examples:
Fe 2+ ions - Green
Fe 3+ ions - Brown
Cu 2+ (Copper) ions - Blue
Co 2+ (Cobalt) ions - Pink
CrO4 2+ (Chromate) ions - Yellow

- catalyst in a reaction
Examples:
Note that <--X--> indicates a reversible reaction with transition metal X.
a) N2 + 3H2 <--Fe--> 2NH3 (Found in the Haber Process)
b) Zn + H2SO4 --CuSO4--> ZnSO4 + H2
c) Contact Process
- S + O2 --> SO2
- 2SO2 + O2 <--V2O5--> 2SO3 (Vanadium Pentoxide / Vanadium (V) oxide)
- SO3 + H2SO4 --> H2S2O7
- H2S2O7 + H2O --> 2H2SO4
Note that the Contact Process involves the production of sulphuric acid. (Chapter 9)

- forms complex ions/compounds (KMnO4 - potassium permanganate, K2Cr2O7 - potassium dichromate, K4Fe(CN)6 - potassium hexacyanoferrate(II).)
- compounds of transition metals contains more than one oxidation number
Example: Iron(II) ions, Iron(III) ions, Copper(I) ions, Copper(II) ions.

And that concludes chapter 4.

Form 4 Chapter 2 - The Structure of the Atom

Chemistry SPM notes for Chapter 2.

Chapter 1 should be simple enough, usually 0-2 questions will come out in Paper 1 only.

Note that you should read through some references as the following notes will not be detailed.

I will just focus on some of the parts. (Disclaimer: I am not predicting which questions will come out in SPM. You can copy & print out the notes below for easier reading/preparation for SPM.)

Experiment to investigate the kinetic theory of matter (Paper 1, Paper 3):

1) Diffusion in a gas - Liquid bromine, gas jar, gas jar cover

2) Diffusion in a liquid - Water, potassium permanganate crystals, petri dish

3) Diffusion in a solid - Jelly, boiling tube, potassium permanganate crystal, stopper

Calculation of a size of oil molecule (Removed from syllabus, just read if you want to. Thanks for pointing it out!):

Volume of 50 drops = V

Volume of 1 drop = v/50

Tiny droplets/Small particles = (V/50)/n

Diameter of small particles = d = 2r

Volume of oil particle = πr^2h = (V/50)/n

Since h is the size of particle,

Note that burette reading should always be in 2 d.p(decimal places)

Changes in the states of matter:

(Quite straightforward here...)

Sublimation - iodine, ammonium chloride, solid carbon dioxide (dry ice).

Experiment: To determine the melting/cooling point of naphthalene (Paper 2, Paper 3)

Naphthalene can be replaced with acetamide.

*When plotting a graph, ensure that:

a. axes are labelled with units

b. points are transferred correctly

c. curve is smooth

To explain the graph (for naphthalene):

During heating:

- the kinetic energy of molecules increases

- the distance between the molecules increases

- the forces of attraction between the molecules become weaker

During cooling:

- the kinetic energy of molecules decreases

- the distance between the molecules decreases

- the forces of attraction between the molecules become stronger

Atomic Structure (Paper 1, Less likely in Paper 2)

John Dalton

- all elements are made up of small indivisible particles called atoms

J.J.Thompson

- discovered negatively charged particles called electrons

Ernest Rutherford

- discovered positively charged particles called protons

Neils Bohr

- atoms have electrons arranged in orbits called electron shells surrounding the nucleus

James Chadwick

- discovered electrically neutral particles called neutrons

Nucleon number = Number of protons + Number of neutrons

(A - nucleon number, Z - proton number, X - element)

You should know most of the elements by now. (Eg. Tin, Silver, Gold, Mercury)

.

Isotopes (Paper 2)

Isotopes are atoms of the same element with the same proton number but different nucleon numbers. (same proton number but different neutron number also accepted.)

.

Uses of Isotopes (Paper 1, Less likely in Paper 2)

Cobalt-60

- kill cancer cells

Carbon-14

- determine absorption rate of phosphorus in plants

- carbon dating

Uranium-235

- nuclear energy

.

End of Chapter 2.
Source :http://anendlesscalm.blogspot.com/search/label/SPM%20Chemistry

Reaction & changes

Physical changes

When a substance undergoes physical changes, its physical composition changes but its chemical composition remain the same.

Eg: Ice, water and water vapour may exist in different physical states - solid, liquid and gas respectively - but they are still made up of the same chemical composition, H

₂O.

Also, physical changes are reversible.

Eg: Water turns into ice when cooled. To get back the water, heat the ice.

Forces like temperature and pressure can effect physical changes.


Chemical changes

The chemical composition of a substance is changed when it undergoes chemical changes. As a result, a new substance is formed. The process is irreversible.

So, how can a chemical change in a substance be identified?

• Change in colour.
• Release or absorption of energy or heat.
• Release of odour.
• Production of gases or solids.

Note: In a chemical reaction, matter is neither destroyed nor created. The particles of an original substance are rearranged to form a new substance. The number of particles that exists before and after the reaction is the same.

Exothermic vs Endothemic

In chemical reactions, heat is either released (exothermic reaction) or absorbed (endothermic reaction). Any change of physical state from solid to liquid is an endothermic reaction as heat is absorbed to push the particles further apart to change its physical state.

However, any change of physical state from liquid to solid is an exothermic reaction as heat is released to push the particles closer together.

To determine whether a reaction is endothermic or exothermic, use a thermometer.

Analysing electrolysis of molten compounds

Malaysia SPM Form 4 Chemistry, Chapter 6: Electrochemistry.

Electrolysis : Decomposing Using the Passage of Electricity

• There are 2 types of electrochemical cell, namely

1. Electrolytic Cell
2. Voltaic Cell (or sometimes is called Galvanic Cell)

• In electrolytic cell, electric current is flowed through an electrolyte to produce a chemical reaction.
• In electrolytic cell, electrical energy is converted into chemical energy, and the process is called electrolysis.

• In voltaic cell, chemical is used to produce electricity.
• In this cell, chemical energy is converted into electrical energy.

The Electrolytic Cell

• A suitable apparatus for electrolysis is shown in Figure above.
• As we can see, the electrode connected to the positive terminal of the cell is positive electrode and is given a name, anode.
• The electrode connected to the negative terminal of the cell is negative electrode and is called the cathode.

• When electricity is passed through an electrolyte, chemical reaction happens.
• In this reaction, chemical is splitting up into 2 new substances.
• All electrolytes are ionic, which means they are composed of positively and negatively charged ions.
• On passing an electric current through the electrolyte, these ions move towards the oppositely charged electrode.

• Most negatively charged ions are non-metal ions, such as oxide (O^2-, chloride (Cl^-), Iodide (I^-), etc.
• During electrolysis, negatively charged ions move towards the positive electrode(anode). The negative ions lose their electron(s) to the anode, which is positively charged.
• The electron(s) is then move to the cathode through the external circuit (the wire).
• The positively charged ions move towards the negative electrode(cathode').
• These positive ions are metal ions, such as copper (Cu²⁺), silver (Ag⁺), lead (Pb²⁺), etc, or hydrogen (H⁺).
• At cathode, positive ions gain electron(s) from the cathode, which has an excess of electrons and therefore an overall negative charge.

• This process results in the chemical decomposition of the electrolyte. It also allows electrons to travel from the cathode to the anode and hence allows conduction of electricity.
• During the electrolysis, electrical energy is supplied to the system to produce a chemical reaction.
• Therefore, during electrolysis, electrical energy convert into chemical energy.

Example 1:Electrolysis of MOLTEN Lead (II) Bromide

• This is composed of lead(II) ions, Pb^{2 +} , and bromide ions, Br^-. Its chemical formula is therefore PbBr₂.
• A suitable apparatus which could be used to carry out this electrolysis is shown in Figure above.
• The bulb helps to show when electricity is flowing in the circuit, and until the lead(II) bromide is completely molten, the bulb does not light up . This confirms that electrolytes have to be molten for the ions to start to move to the electrodes and thereby conduct electricity.

At the Cathode	At the Anode
Observation When electricity is flowing, a silvery deposit of lead metal forms on the cathode. In fact, as it is molten, it is more likely to drip off in a molten blob.	Observation When electricity is flowing, brown fumes of bromine gas are seen at the anode.
Half equation Pb2+ + 2e ---> Pb	Half equation 2Br- ---> Br2 + e
Explanation The lead(II) ions, as they are positive, move to the negative cathode, where each ion gains two electrons to form a lead atom. Any reaction at a cathode involved is again in electrons. This is called reduction or more exactly, cathodic reduction .	Explanation The bromide ions, as they are negative, move to the positive anode, where each loses an electron to form a bromine atom. Then two of these newly formed atoms combine to form bromine gas. Any reaction at an anode involves a loss of electrons.

• In summary, the lead(II) bromide is split into its component elements :
  

PbBr₂ ---> Pb + Br₂

Example 2: Electrolysis Of Molten Lead(II) Oxide

Understanding properties of electrolytes and non-electrolytes

Malaysia SPM Form 4 Chemistry, Chapter 6: Electrochemistry.

Electrolytes And Non-electrolyte

Conductor

A conductor is a substance which conducts electricity but is not chemically changed during the conduction .

Insulators

A non-conductor is a substance which does not allow the passage of electricity. Sometimes these non-conductors are used to protect something from electricity. They are then called insulators .

Electrolytes

Electrolytes are compounds which when molten or dissolved in water conduct electric current and are decomposed in the process .

Non-electrolyte

A non-electrolyte is a liquid which does not allow the passage of electricity.

Examples of non-electrolytes, weak electrolytes and strong electrolytes

• Examples of electrolytes are acids, alkalis and salts dissolved in water or molten salts .
• All these are ionic substances.
• Solid ionic substances do not conduct electricity, as their ions are held together in fixed positions by strong forces .
• In order for the ions to move, the solid must be dissolved in water or made molten, thereby destroying the ionic lattice .

Understanding formation of compounds ( chemistry f.4)

(9Stability of Noble Gases

What is Chemical Compound

• A chemical compound is a substance that is formed by more than one elements that bond together chemically in a fixed proportions.
• In periodic table, there are only 118 elements, and about 1/3 of them are synthetic elements.
• Only a few substances exist as element (Not Compound) in nature.
• The table below shows some examples of substance exist as element in nature.

Element exist as monoatomic gas	Element exist as diatomic Molecule(gas)	Element exist as solid
Helium (He) Neon (Ne) Argon (Ar) Krypton (Kr) Xenon (Xe) Radon (Rn)	Oxygen (O2) Nitrogen (N2)	Carbon (Graphite & Diamond) Gold Silver Platinum

• In nature, we can find millions of substances, which means most of the chemical substances exist as compound in nature.
• In short, elements tend to form compound in nature.

Why Elements Tend to Form Compound?

• A compound is formed by 2 or more elements hold together by a force called chemical bond.
• Before studying why elements like to bond together, we need to know why certain elements such as Helium and Neon do not form any bonds with other elements.

Why Noble Gases Don't Form Compound

• In previous chapter, we have discussed that Group 18 elements (Noble Gases) exist as monoatom in nature.
• They are inert in nature and do not react with any other elements (or themselves) to form any chemical compounds.
• In other words, they are chemically very very stable (or chemically very very non-reactive).

Duplet and Octet Electron Arrangement

• The charge on the nucleus and the number of electrons in the valence shell determine the chemical properties of an atom.
• The stability of noble gas is due to their electrons arrangement.
• The diagram above shows the first four elements of Noble Gas.
• We can see that the outer most shell (valence shell) of Helium has 2 electrons. We call this duplet electron arrangement. We should take notes that the maximum number of electrons can be filled in the first shell is 2 electrons, which means 2 electrons in the first shell is considered FULL.
• The valence shell all other Group 18 elements (including Xenon and Radon which is not shown in the diagram) has 8 electrons, and we call this octet electron arrangement.
• When the electron arrangement of an atom is duplet or octet, the energy of the electrons is very low, and it is very difficult (even though it is not impossible) to add or remove electrons from the atom.
• This explain why noble gases are reluctant to react with all other elements.

The Octet Rule

• So far we have learnt that the electon arrangement of noble gases are octet duplet, and this is the most stable electron arrangement of an atom.
• Atoms of other main group elements which is not octet tend to react with other atoms in various ways to achieve the octet.
• The tendency of an atom to achieve an octet arrangement of electrons in the outermost shell is called the octet rule.
• If the outermost shell is the first shell, then the maximum number of electrons is two, and the most stable electron arrangement will be duplet.
• A configuration of two electrons in the first shell, with no other shells occupied by electrons, is as stable as the octet electron arrangement and therefore is also said to obey the octet rule.

Important Notes Most of the elements (except noble gases) are chemically not stable. It is the aim of every atom to achieve the duplet or octet electron arrangement. This makes them very stable. It is only the valence electrons in the outermost shell involved in bonding. The electrons in the inner shells are not involved. The maximum number of electrons in the first shell is two. This is called a duplet. The maximum number in the second shell is eight. This is called an octet.

How Atoms Achieve Duplet or Octet Electron Arrangement?

• Atoms can achieve duplet or octet electron arrangement in 3 ways:

1. throw away the excess electron(s)
2. receiving electron(s) form other atom if they are lack of electron(s)
3. sharing electron

• 2 types of chemical bonds are commonly formed between atoms, namely

1. Ionic Bond
2. Covalent Bond

The Ionic Bond

• By releasing or receiving electron(s), the atoms will become ions and consequently form ionic bond between the ions.
• Ionic bonds are always form between metal and non-metal. For example, sodium (metal) react with chlorine (non-metal) will form an ionic bond between sodium ion and chloride ion.
• The compounds formed is called the ionic compound.
• Some time, an ionic bond is also called electrovalent bond.

[edit] The Covalent Bond

• By sharing electron(s), the atoms will form covalent bond between the atom and the molecule formed is call the covalent molecule.
• Covalent bond is always formed between non-metal with another non-metal.