Why latex does not coagulates when it is inside the rubber tree?

Latex is a colloidal solution. It consists of rubber particles dispersed in water. Each rubber particles is made up of many long-chain rubber molecules enclosed by a protein-like membrane which is negatively charged. The negative charges around the rubber particles cause repulsion between these particles to occur when they are near each other. Hence the repulsion between the negatively-charges particles prevents the rubber particles from coming close to each other. Hence latex could not coagulate.  

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SPM Form 5: Chemicals for Consumers (Checklist)

• Saponification: Alkaline hydrolysis of an ester where the ester is boiled with sodium hydroxide, NaOH or potassium hydroxide, KOH solution to produce an alcohol and a sodium or potassium salt of a carboxylic acid.
• Soap is not an effective cleansing agent in hard water but is only effective in soft water. Detergent is an effective cleansing agent in both hard water and soft water.  
• Soft water: Water that does not contain mineral salts such as magnesium salt or calcium salts.
• Hard water: Water that contain mineral salts.
• Food additives are substances added to food to preserve flavor or enhance its taste and appearance. Types of additives are preservatives, antioxidants, flavourings, stabilizers, thickening agents and food dyes (colourings).
• Medicines: a compound or preparation used for the treatment or prevention of disease, especially a drug or drugs taken by mouth. Types of modern medicines are analgesics, antibiotics, stimulants, antidepressant and antipsychotic. 

SPM Form 5: Carbon compounds (Checklist)

• Organic compounds: Carbon-containing compounds that can be obtained from living things.
• Inorganic compounds: Non-carbon-containing compounds that can be obtained from non-living things.
• Hydrocarbons: Compounds that contain only carbon and hydrogen. In a saturated hydrocarbon, all the bonds are single bonds. Alkane is another name for a saturated hydrocarbon. Unsaturated hydrocarbons are a hydrocarbon that contains one or more double or triple bonds are an unsaturated hydrocarbon. There are three types of unsaturated hydrocarbons alkenes, alkynes, and aromatic hydrocarbons.
• Alkanes: Hydrocarbons having the general formula C_nH_2n+2, where n = 1, 2
• Alkenes: Hydrocarbons that contain one or more carbon-carbon double bonds. They have the general formula C_nH_2n.      
• Alcohol: An organic compound containing the hydroxyl group -OH.
• Aldehydes: Compounds with a carbonyl functional group and the general formula RCHO, where R is an H atom, an alkyl, or an aromatic group.       
• Carboxylic acids: Acids that contain the carboxyl group -COOH.
• Esters: Compounds that have the general formula R’COOR, where R’ can be H or an alkyl group or an aromatic group and R is an alkyl group or an aromatic group.
• Homologous series: A series of compounds in which each member differs from the next by a specific number and kind of atoms.
• Esterification: A reaction between an alcohol and a carboxylic acid in the presence of concentrated sulphuric acid as a catalyst through boiling to produce an ester and water. e.g. reaction of ethanol, C₂H₅OH with ethanoic acid, CH₃COOH in the presence of concentrated sulphuric acid, H₂SO₄ produces ethyl ethanoate, CH₃COOC₂H₅ which is an ester with a pleasant fragrant smell.
• Fats: Solid triester of glycerol and mostly saturated fatty acids.
• Vulcanization of rubber:  A process whereby rubber is reacted with sulphur to enable the formation of cross-linkages by sulphur atoms between the rubber molecules through covalent bonds. Rubber can be vulcanized by dipping natural rubber sheets into disulphur dichloride solution in methylbenzene. Vulcanized rubber is hard whereas unvulcanized rubber is soft. Vulcanized rubber is more elastic than unvulcanized rubber.
• Coagulation of latex: A process of converting liquid latex to solid natural rubber by adding an acid. Coagulation of latex occurs rapidly when an acid is added to the latex. Latex does not coagulate when an alkali is added to the latex. Alkaline solutions contain hydroxide ions which neutralize the acid produced by bacteria. Hence, it prevents the latex from coagulating. 

SPM Form 5: Rate of reaction (Checklist)

• Rate of reaction: The decrease in the amount of reactant used or the increase in the amount of product obtained in a given time. The rate of reaction measures how much product is formed in a certain time. Some reactions are slow, such as rusting, and some are fast, like burning.
• Average of reaction: The average value of the rate of reaction within a specified period of time.
• Rate of reaction at a given time: The actual rate of reaction at a specific time.
• Collision theory: A model that explains reaction rate as the result of particles colliding with a certain minimum energy.
• Effective collision: A collision in which the particles meet with sufficient energy and an orientation that allows them to react.
• The rate of reaction decreases with the increase of time.
• The rate of reaction increases as the total surface area of a solid reactant increases.
• The rate of reaction increases as the concentration of reactant increases.
• The rate of reaction increases as the temperature of a reactant increases.
• The rate of reaction increases when a positive catalyst is used and the rate of reaction increases when the amount of a positive amount used increases.
• Catalyst: A substance which alters the rate of a chemical reaction while it remains chemically unchanged at the end of the reaction. 

SPM Chemistry Form 4 and Form 5: Checklist (Definitions)

Form 4:

Chapter 1: Introduction to Chemistry

Chapter 2: The Structure of the Atom

Chapter 3: Chemical Formulae and Equations

Chapter 4: Periodic Table of Elements

Chapter 5: Chemical Bonds

Chapter 6: Electrochemistry

Chapter 7: Acids and Bases

Chapter 8: Salts

Chapter 9: Manufactured Substances in Industry

Form 5:

Chapter 1: Rate of Reaction

Chapter 2: Carbon compounds

Chapter 3: Oxidation and reduction

Chapter 4: Thermochemistry

Chapter 5: Chemical for Consumers

SPM Form 4: Salts (Checklist)

• Salts: An ionic compound formed when the hydrogen ion in an acid is replaced by a metal ion or an ammonium ion. Salts are formed as the product of an acid reaction with an alkali. Soluble salts dissolve in water. Insoluble salts do not dissolve in water.
• Recrystallization: A technique used to purify crystals of a soluble salt by carrying out the crystallization process again on these crystals.
• Precipitation reaction: A reaction which involves the reaction between two aqueous reactants to form an insoluble substance.
• Continuous variation method: An experiment which involves the reaction of a solution at fixed volume and another solution with volumes that varies uniformly.
• All carbonate salts are decomposed on heating to liberate carbon dioxide gas, CO₂ except sodium carbonate, Na₂CO₃ and potassium carbonate, K₂CO₃.
• All nitrate salts are decomposed on heating to liberate nitrogen dioxide gas, NO₂ and oxygen gas, O₂ except sodium nitrate, NaNO₃ and potassium nitrate, KNO₃ which liberate oxygen gas, O₂ only.
• The presence of carbonate ions can be confirmed by adding dilute acids.
• The presence of sulphate ions can be confirmed by adding acidified barium chloride, BaCl₂ solution.
• The presence of chloride ions can be confirmed by adding acidified silver nitrate, AgNO₃ solution.
• The presence of nitrate ions can be confirmed by adding dilute sulphuric acid, H₂SO₄ followed by iron (II) sulphate, FeSO₄ solution and a little concentrated sulphuric acid, H₂SO₄.
• The identify of cation can be determined by using sodium hydroxide, NaOH solution and ammonia, NH2 solution except aluminium ion, Al³⁺ and lead (II) ion, Pb²⁺.
• Ammonium ion can be tested by heating with a strong alkali or adding Nessler’s reagent.
• Iron (II) ion and iron (III) ion can be tested by using potassium hexacyanoferrate (II), K₄Fe(CN)₆ solution, potassium hexacyanoferrate (III), K₃Fe(CN)₆ or potassium thiocyanate, KSCN solution.
• Lead (II) ions can be tested by using potassium iodide, KI solution.

SPM Form 4: Manufactured Substances in Industry (Checklist)

• Sulphuric acid (H₂SO₄) is a highly corrosive acid made from sulfur dioxide; widely used in the chemical industry.
• Contact process: the industrial process for making sulfuric acid.
• Ammonia (NH₃) is colorless, pungent gas composed of nitrogen and hydrogen.

• Haber process: the process for making ammonia from nitrogen and hydrogen, in industry.
• Ammonium fertilizer: A salt that is prepared from the reaction between ammonia and an acid. Ammonium sulphate, (NH₄)₂ SO₄ which is an ammonium fertilizer can be prepared from the reaction between ammonia, NH₃ solution and sulphuric acid, H₂SO₄.
• Alloys: A mixture of two or more elements with certain fixed composition in which the major component is a metal. Alloy is harder than pure metal.
• Metal corrosion: The gradual destruction of a metal by reaction with its environmental. Iron rusts faster than steel. Stainless steel does not rust.
• Polymer: a compound containing very large molecules, formed by polymerization.
• Synthetic polymers are derived from petroleum oil, and made by scientists and engineers. Examples of synthetic polymers include nylon, polyethylene, polyester, Teflon, and epoxy.
• Natural polymers occur in nature and can be extracted and often water-based. Examples of naturally occurring polymers are silk, wool, DNA, cellulose and proteins.
• Glass: A homogeneous material with a random, liquid-like molecular structure.
• Ceramic is a hard, unreactive material that can withstand high temperatures, made by baking clay in a kiln; ceramics are non-conductors.
• A "composite" material is when two or more different materials are combined together to create a superior and unique material.

SPM Form 4: Periodic Table of elements (Checklist)

• Periodic Table: the table showing the elements in order of increasing proton number; similar elements are arranged in columns called groups.
• Group: A vertical column of elements in the Periodic table.
• Period: A horizontal row of the Periodic Table; its number tells you how many electron shells there are.
• Alkali metals: the Group I elements of the Periodic Table, which include lithium, sodium, potassium, rubidium, caesium and francium.
• Alkaline earth metals: the Group II elements of the Periodic Table, which include beryllium, magnesium, calcium, strontium, barium, and radium.
• Halogens: the Group VII elements of the Periodic Table, which include fluorine, chlorine, bromine, iodine and astatine.
• Noble gases: the Group 18 elements of the Periodic Table; they are called ‘noble’ or inert gases because they are so unreactive, which include helium, neon, argon, krypton, xenon and radon.
• Transition elements: the elements in the wide middle block of the Periodic Table (elements in group 3 to group 12).
• Metal: an element that shows metallic properties (for example conducts electricity, and forms positive ions)
• Non-metal - an element that does not show metallic properties: the non-metals lie to the right of the zig-zag line in the Periodic Table.
• Amphoteric oxide: An oxide that exhibits both acidic and basic properties.
• The atomic radius is a term used to describe the size of the atom.
• The ionization energy is the energy required to completely remove an electron from a gaseous atom or ion.
• Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom.
• Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons.

SPM Form 4: Chemical Formulae and Equations (Checklist)

• Relative atomic mass, A_r, is defined as the mass of one atom of an element relative to 1/12 of the mass of an atom of carbon-12, which has a mass of 12.00 atomic mass units.
• Relative molecular mass, M_r, is defined as the mass of one molecule of an element or compound relative to 1/12 of the mass of an atom of carbon-12, which has a mass of 12.00 atomic mass units.
• Empirical formula is a formula showing the simplest ratio of atoms present.
• Molecular formula is a formula showing the actual number of atoms of each element present in one molecule.
• Chemical equation is a representation of a chemical reaction in words or using chemical formulae.
• The molar mass is the mass of one mole of a sample.                
• Molar volume is the volume of one mole of a substance. The molar volume of an ideal gas at STP is 22.4 L/mol.
• Avogadro's constant is the number of particles found in one mole of a substance. Avogadro’s constant=6.0221 x 10²³ particles per mole.
• Avogadro’s Law: equal volume of all gases measured under the same conditions of temperature and pressure contain equal numbers of molecules.

SPM Form 4: Acids and Bases (Checklist)

• Acid: A chemical substance which ionizes in water to produce hydrogen ions. An acid will only show its acidic properties when it is dissolved in water. E.g. ethanoic acid CH₃COOH only exhibit acidic properties when water is present.
• Alkali: A chemical substance which ionizes in water to produce hydroxide ions. An alkali will only show its alkaline properties when it is dissolved in water. E.g. barium hydroxide, Ba(OH)₃ only exhibit alkaline properties when water is present.
• A strong acid is an acid that is completely dissociated in an aqueous solution. E.g. HCl (hydrochloric acid), H₂SO₄ (sulfuric acid), HNO₃ (nitric acid).
• A weak acid is an acid that is partially dissociated in an aqueous solution. E.g. CH₃COOH
• A strong base is a base that is completely dissociated in an aqueous solution. E.g. KOH, NaOH
• A weak base is a base that is partially dissociated in an aqueous solution. E.g. NH₄OH
• Molarity: The number of moles of solute in 1 dm³ solution. When the molarity of an acid increases, its pH value decreases. When the molarity of an alkali increases, its pH value increases.
• pH is a measure of hydrogen ion concentration; a measure of the acidity or alkalinity of a solution. Aqueous solutions at 25°C with a pH less than seven are acidic, while those with a pH greater than seven are basic or alkaline. A pH level of is 7.0 at 25°C is defined as 'neutral' because the concentration of H₃O⁺ equals the concentration of OH⁻ in pure water.
• The two factors which determine the pH value of an acid or an alkali are degree of ionization and concentration of the acid or alkali.
• Neutralization: A reaction between an acid and an alkali or a base to produce salt and water.
• Acid-base titration: A quantitative analysis to determine the volume of an acid required to exactly neutralize a fixed volume of an alkali with the help of a suitable indicator.
• End-point: A stage achieved in titration where the volume of acid added exactly neutralizes a fixed volume of an alkali.

SPM Form 4: Introduction to Chemistry (Checklist)

Scientific method:

Making observation → Making inference → Identifying problem → Making hypothesis → Identifying variables → Controlling variables → Planning an experiment → Collecting data → Interpreting data → Making conclusion → Writing report

• Fixed variables: the factor that is kept constant throughout the experiment
• Manipulated variables: the factor that is purposely changed in an experiment
• Responding variables: the factor that changes with the manipulated variables
• Hypothesis: a general statement about the relationship between a manipulated  variable and a responding variable in order to explain an event or phenomenon. 

SPM Form 5: Thermochemistry (Checklist)

• Thermochemistry is the scientific study of heat that is released or absorbed during chemical changes.
• Exothermic reaction: A chemical reaction that involves liberation of heat to the surroundings. The increase in the temperature of the reacting mixture indicates an exothermic reaction.
• Endothermic reaction:  A chemical reaction that involves absorption of heat from the. The decrease in the temperature of the reacting mixture indicates an endothermic reaction.
• Heat of precipitation: The change in heat energy when one mole of a precipitate is formed from their ions in aqueous solutions.
• Heat of displacement: The change in heat energy when one mole of a metal is displaced from its salt solution by a more electropositive metal.
• The higher the position of the more electropositive metal is in the electrochemical series, the bigger the value of heat displacement of copper by the metal.
• Heat of neutralization: The change in heat energy when one mole of water is formed from the neutralization of an acid and an alkali.
• Heat of neutralization for reaction between any strong acid and any strong alkali is a constant.
• The values of heat neutralization for reactions between strong acids and strong alkalis are higher than that of reactions between strong acids and weak alkalis, weak acids and strong alkalis or weak acids and weak alkalis. 
• Heat of combustion: The change in heat energy when one mole of substance is completely combusted in excess oxygen under standard conditions.
• The greater the number of carbon atoms per molecule of an alcohol, the greater the value of heat of combustion of the alcohol.

SPM Form 5: Oxidation and Reduction (Checklist)

• Oxidation: A reaction that occurs when a substance gains oxygen.
• Oxidizing agent: The substance that oxidizes other reacting substances in a reaction while it is being reduced.
• Reduction: The reaction that occurs when a substances loses oxygen.
• Reducing agent: The substance that reduces other reacting substances in a reaction in a reaction while it is being oxidized.
• Redox reaction: A chemical reaction in which oxidation and reduction occur simultaneously.
• Metal displacement reaction: A reaction where a more electropositive metal replaces a less electropositive metal from its salt solution. In a metal displacement reaction, a more electropositive metal acts as the reducing agent whereas the ion of the less electropositive metal in the salt solution acts as the oxidizing agent.
• Halogen displacement reaction: A reaction where a more reaction halogen replaces a less reactive halogen from its halide solution. In a halogen displacement reaction, a more reactive halogen acts as the oxidizing agent whereas the halide ion of the less reactive halogen acts as the reducing agent.
• Rusting of iron = the corrosion of ion.  Rusting can be prevented when iron is in contact with a more electropositive metal. Rusting occurs faster when iron is in contact with a less electropositive metal.
• Reactivity series of metals: A list of metal that are arranged according to their reactivity with oxygen. The descending order of reactivity of metal with oxygen is magnesium, zinc, iron, lead and copper.
• Carbon is placed between aluminium and zinc in the reactivity series. Magnesium is placed higher than carbon in the reactivity series. Carbon acted as a reducing agent because the metal oxide was reduced to metal in the presence of carbon.
• Hydrogen is placed between zinc and iron in the reactivity series. Hydrogen gas acted as a reducing agent because the metal oxide was reduced to metal in the presence of hydrogen gas.
• In an electrolytic cell, oxidation takes place at the anode whereas reduction takes place at the cathode.
• In a chemical cell, oxidation takes place at the negative terminal (anode) whereas reduction takes place at the positive terminal (cathode). 

SPM Form 4: Electrochemistry (Checklist)

• Electrode: A conductor in the form of a wire, rod or plate which carry electric current in and out of an electrolyte during electrolysis.
• Electrolyte: A substances that can conduct electricity in molten state or aqueous solution and is decomposed by electric current.
• Non-electrolyte: A substances that cannot conduct electricity in molten state or aqueous solution.
• Anion: A negatively-charged ion.
• Cation: A positively-charged ion.
• Anode: An electrode which is connected to the positive terminal of the source of electricity during electrolysis. (donation of electrons)
• Cathode: An electrode which is connected to the negative terminal of the source of electricity during electrolysis. (acceptance of electrons)
• Electrolysis: The process whereby a compound is broken down into its constituent elements when electricity is passed through an electrolyte.
• Aqueous solution: A solution produced when a compound is dissolved in water.
• Electrochemical series: An arrangement of metals based on the tendency of each metal to donate electrons.
• The lower the position of an ion in the electrochemical series, the higher is the tendency of the ions to be discharged.
• Purification of metals: The process of obtaining a pure metal from an impure metal through electrolysis. E.g. Impure copper can be purified through electrolysis when the impure copper is used as the anode and a pure copper is used as the cathode.
• Electroplating of metals: The process of coating a layer of metal onto another metal using electrolysis.
• Simple voltaic acid: A cell that converts chemical energy to electrical energy. The chemical reactions in a simple voltaic cell produce electricity. No current flow will flow if both electrodes are made of the same metal.
• Electropositivity: A measurement of the ability of an atom to donate electrons to form a positive ion.
• Displacement reaction: A reaction where a more electropositive metal displace another metal from its salt solution. A metal which has a higher position in the electrochemical series is able to displace the metal below it in the series from their salt solutions.

SPM Form 4: Chemical Bonds (Checklist)

• An ionic bond is a chemical link between two atoms caused by the electrostatic force between oppositely-charged ions in an ionic compound. Examples: an ionic bond between the sodium and chloride ions in table salt, NaCl.

• A covalent bond is a chemical link between two atoms in which electrons are shared between them. Examples: a covalent bond between the oxygen and hydrogen in a water molecule (H₂O).

• Ionic compounds: a compound that consists of positive ions and negative ions arranged in the crystal lattice and are attracted to each other by strong electrostatic forces. Ionic compounds such as magnesium oxide, MgO, sodium chloride, NaCl and iron(III) chloride, FeCl₃ can be respectively prepared through a reaction between a metal and a non-metal.
• Covalent compounds are those in which the elements share electrons via covalent bonds.
• Ionic compounds have high melting and boiling point whereas covalent compounds have low melting and boiling points.
• Ionic compounds are mostly soluble in water but insoluble in organic solvents. Covalent compounds are mostly insoluble in water but soluble in organic solvents.
• Ionic compounds can conduct electricity in molten state but cannot conduct electricity in solid state. Covalent compounds cannot conduct electricity in solid and molten states.
• Melting points: the temperature at which a solid turns into a liquid at a fixed pressure.
• Boiling point: the temperature at which the vapor pressure of a liquid equals the external pressure surrounding the liquid.
• Solubility: the maximum quantity of a substance that may be dissolved in another or the maximum amount of solute that may be dissolved in a solvent.

SPM Chemistry Syllabus Form 5

Chapter 1: Rate of reaction

1.1 Rate of reaction

1.2 Factors affecting the rate of reaction

1.3 Collision theory

Chapter 2: Carbon Compounds

2.1 Carbon compounds

2.2 Alkanes

2.3 Alkenes

2.4 Isomerism

2.5 Alcohols

2.6 Carboxylic Acids

2.7 Esters

2.8 Fats

2.9 Natural rubber

Chapter 3: Oxidation and Reduction

3.1 Redox reactions

3.2 Rusting as a redox reaction

3.3 The reactivity series of metals and its application

3.4 Redox reactions in electrolytic and chemical cells

Chapter 4: Thermochemistry

4.1 Energy changes in chemical reactions

4.2 Heat of precipitation

4.3 Heat of displacement

4.4 Heat of neutralization

4.5 Heat of combustion

4.6 The existence of various energy sources

Chapter 5: Chemical for Consumers

5.1 Soap and detergent

5.2 Food additives

5.3 Medicines

5.4 The existence of chemicals

SPM Chemistry Syllabus Form 4

Chapter 1: Introduction to Chemistry

1.1 Chemistry and its importance

1:2 Scientific methods

Chapter 2: The Structure of the Atom

2.1 Matter

2.2 The atomic structure

2.3 Isotopes

2.4 The electronic structure of an atom

Chapter 3: Chemical Formulae and Equations

3.1 Relative atomic mass and relative molecular mass

3.2 The mole and the number of particles

3.3 The mole and the mass of substances

3.4 The mole and the volume of gas

3.5 Chemical Formulae

3.6 Chemical Equations

Chapter 4: Periodic Table of Elements

4.1 The Periodic Table of Elements

4.2 Group 18 elements

4.3 Group 1 elements

4.4 Group 17 elements

4.5 Elements in a period

4.6 Transition elements

Chapter 5: Chemical Bonds

5.1 Formation of compounds

5.2 Ionic bonds

5.3 Covalent bonds

5.4 The properties of ionic and covalent compounds

Chapter 6: Electrochemistry

6.1 Electrolytes and non-electrolytes

6.2 Electrolysis of molten compounds

6.3 Electrolysis of aqueous solutions

6.4 Electrolysis in industries

6.5 Voltaic cells

6.6 The electrochemical series

Chapter 7: Acids and bases

7.1 Acids and bases

7.2 The strength of acids and alkalis

7.3 Concentrations of acids and alkalis

7.4 Neutralization

Chapter 8: Salts

8.1 Salts

8.2 Qualitative analysis of salts

Chapter 9: Manufactured Substances in Industry

9.1 Sulphuric acid

9.2 Ammonia and its salts

9.3 Alloys

9.4 Synthetic polymers

9.5 Glass and polymers

9.6 Composite materials