1. CLASSES OF CHEMICAL REACTION
SMK NEGERI 13 BANDUNG

2. GENERAL CLASSIFICATION
Precipitation
Hydrolysis
Decomposition
Ionisation
Dissociation
Acid-Base
Oxidation-Reduction
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3. AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Precipitation
Any solution in which an insoluble product form; e.g., silver chloride precipitates when silver nitrate solution is added to sodium chloride solution
AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
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4. Al3+(aq) + H2O(l) → AlOH2+(aq) + H+(aq)
Hydrolysis
Any reaction in which water reacts with some other substance; e.g.,
Al3+(aq) + H2O(l) → AlOH2+(aq) + H+(aq)
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5. CuSO4(s) → CuO(s) + SO3(g)
Decomposition
A reaction in which a substance breaks down or ‘decomposes’ into two or more simple substance; e.g., copper (II) sulfate decomposes to copper (II) oxide and sulfur trioxide
CuSO4(s) → CuO(s) + SO3(g)
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6. HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)
Ionisation
A reaction in which ions are formed; e.g., the reaction of gaseous hydrogen chloride with water
HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)
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7. Dissociation
A reaction, usually in the gaseous state or in solution, in which one molecule ‘dissociate’ into two fragments; e.g., in the ‘dissociation’ of N2O4 into NO2
N2O4(g) → 2NO2(g) at 100 oC
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8. ACID-BASE REACTION Historical Background
Bronsted Theory of Acid and Bases
Acid-Base Indicators
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9. Historical Background
In 1675, Robert Boyle geve some properties of acids and alkalis.
The properties of acids included their sour taste, and their action on vegetable colours such as juice of violets, extracts of cochineal and Brazil wood.
Alkalis (bases) restore the original colours that had been changed by acids.
Alkalis also reacted with acids so that the caracteristic properties of each dissapeared and a salt was formed.
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10. Historical Background
The work of Graham and Liebig in the 19th century led to the idea that acids were compounds of hydrogen, and that the hydrogen in an acid could be ‘displaced’ by metal.
Acid were substances that gave up hydrogen ions, H+, in aqueous solution.
Base were substances that gave up hydroxide ions, OH-, in aqueous solution.
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11. Bronsted Theory of Acids and Bases
A definition of terms acid and base, applicable to other solvents as well as water and from which a clear cut scheme for comparing acid and base strengths can be made, was devised by Bronsted in 1923.
Acids are proton donors
Base are proton acceptors
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12. Acid-base Indicators
Acid-base indicators are substances use to determine whether a particular solution in acidic or alkaline.
The first indicators were vegetable extracts of various sorts, the best known is litmus that extracted from a lichen. Litmus takes on a red colour in acidic solution and a blue colour in alkaline solution.
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13. Acid-base Indicators
Two of commonest indicators in use today are methyl orange and phenolphtalein.
Mehtyl orange appears red in acidic solutions and yellow in alkaline solution.
Phenolphtalein is colourless in acid and bright pink in alkali.
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14. OXIDATION-REDUCTION REACTIONS
Historical Background
Oxidation Number
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15. Historical Background
In the 17th century, scientist became interested in the power of steam and this further enlieved interest in the subject of combustion.
Chemists now began to classify reactions in which any element gained oxygen as oxidation, and reactions in which a compound lost oxygen as reduction.
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16. Historical Background
Today definition of oxidation and reduction have been broadened, and now the most general definitions are:
oxidation involves loss of electrons
reduction involves gain of electrons
A substance that will oxidise another is called an oxidisier or oxidant, and the oxidant must gain electrons when it react.
A substance that will reduce another is called an reducer or reductant, and the reductant must donates electrons when it react.
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17. Oxidation Numbers
The oxidation numbers of the elements in compounds that are known to be made up of ions, or that dissociate to form ions in solution, are simply the charges on the ions. Thus the oxidation numbers in the following compounds are shown in colour;
Na Cl Fe Cl3 Mg O
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18. Oxidation Numbers
The oxidation numbers of the element in compounds or entities that are normally regarded as being composed of ions – e.g., C and O in CO2 are obtained by assuming that the atoms have positive and negative charges, which can calculated according the five rules.
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19. Oxidation Numbers The five rules are:
The more electronegative atom will have the negative oxidation number.
In a neutral molecule, the algebraic sum of the individual oxidation numbers is zero; in an ion, the algebraic sum of the oxidation numbers is the charge of the ion.
Atoms of elementary substance have an oxidation number of zero; e.g., the oxidation number in Cl2 is zero.
Hydrogen is generally given an oxidation number of +1 in its compounds, except when its combine with a metal.
Electronegative atoms have oxidation numbers equal to the charges present on normal negative ions of the element, provide the atom in question is the most electronegative atom in the compound.
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