Chemistry [Ch-1] - Chemical Reactions and Equations

properties of chemical reaction

properties of chemical reaction

• it is a process in which new substances with new properties are formed

• a rearrangement of atoms takes place between the reactants

• it involves breaking of old chemical bonds between the rearranged atoms

magnesium ribbon, after being cleaned with sand paper, burns in air with a white flame to form magnesium oxide

Mg (ribbon) + O2 →[heat] MgO (white powder)

reason why the magnesium ribbon is cleaned by rubbing with a sand paper before burning in air

this is done to remove the protective layer of magnesium oxide from the surface of magnesium ribbon so that it may easily combine with oxygen in air

examples of chemical reactions that occur in daily life

souring of milk, formation of curd, cooking of food, digestion of food, respiration, fermentation of grapes, rusting of iron, burning of fuels, burning of candle wax and ripening of fruits

list of characteristics of chemical reactions

these characteristics help us tell whether a chemical reaction has taken place-

• evolution of a gas

• formation of a precipitate

• change in color

• change in temperature

• change in state

explain the evolution of a gas

• it is the process of gas being separated from its compound and rising up in a chemical reaction

• we can also observe a rise in temperature in this phenomenon

examples of evolution of a gas

• the chemical reaction between zinc (Zn) and dilute sulphuric acid (H2SO4) shows evolution of H2

• the chemical reaction between sodium carbonate (Na2CO3) and dilute hydrochloric acid (HCl) shows evolution of CO2

explain the formation of a precipitate

• a precipitate is a solid product which separates out from the solution during a chemical reaction

• it can be formed by-

  • mixing aqueous solutions when one of the products is insoluble in water

  • by passing a gas into an aqueous solution of a substance

examples of formation of precipitate

• lead nitrate solution (Pb(NO3)2) + potassium iodide solution (KI) gives a yellow precipitate of lead iodide (PbI2)

• sulphuric acid (H2SO4) + barium chloride (BrCl2) gives white precipitate of barium sulphate (BrSO4)

• sodium hydroxide solution (NaOH) + copper sulphate solution (CuSO4) gives blue precipitate of copper hydroxide (Cu(OH)2)

explain and give examples of change of color in chemical reactions

• some reactions are characterised by a change in color of the reactants

• for example;

  • potassium permanganate solution [KMnO4] (purple) + citric acid [C6H8O7] (lemon juice etc) will make it colorless

  • sulphur dioxide gas [SO2] + acidified potassium dichromate (orange) [K2Cr2O7] solution will make it green

  • potassium iodide solution + lead nitrate solution will change from colourless to yellow

explain the change in temperature in chemical reactions

• some reactions are characterised by a change in temperature

• this is can happen in two scenarios

  • chemical reactions produce heat energy, which causes the temperature of the reaction to increase. this is called an exothermic reaction

  • chemical reactions absorb heat energy, which causes the temperature of the reaction to decrease. this is called an endothermic reaction

examples of exothermic reactions

• quick lime (CaO) + water (H2O) gives slaked lime (Ca(OH)2) and heat

• burning of magnesium wire in air to form magnesium oxide

• decomposition of vegetable matter into compost

• respiration

• burning of natural gas

• all combustion reactions

examples of endothermic reactions

• barium hydroxide (Ba(OH)2). + ammonium chloride (NH4Cl) takes in heat to give barium chloride (BaCl2) + water (H2O) + ammonia (NH3)

• decomposition of calcium carbonate

• photosynthesis (sunlight energy is absorbed)

• electrolysis of water to form hydrogen and oxygen (electric energy is absorbed)

explain and give examples of the change in state in chemical reactions

• some reactions are characterised by a change in state between solid, liquid and gas

• for example;

  • when wax is burned, CO2 and H2O is formed. thus, the state changes from solid to gas and liquid

properties of chemical equation

• the method of representing a chemical reaction with the help of symbols and formulae of the substance involved in it

• the substances which combine are known as reactants

• the new substances produced are known as products

definition of balanced equation

balanced equation has an equal number of atoms of each element in LHS & RHS. this is because of the law of conservation of mass, where atoms can neither be created nor destroyed in a reaction

definition of unbalanced equation

unbalanced equations have an unequal number of atoms of each element in LHS & RHS. in other words, they have unequal masses of various elements in reactants and products

list of elements that exist in diatomic form

oxygen (O2), hydrogen (H2), nitrogen (N2), fluorine (F2), chlorine (Cl2), bromine (Br2), iodine (I2)

NOTE: all the other elements exist in monoatomic form. take note that carbon (C) is also monoatomic, despite seeming like its not

chemical reaction involved in launching of space shuttles

liquid hydrogen burning in liquid oxygen to form water and a huge amount of energy to lift the shuttle

three ways to make chemical equations more informative

• indicating physical states

• indicating heat changes

• indicating conditions under which reaction takes place

method of indicating physical states of reactants and products in an equation

there are 4 physical states that can be indicated in a reaction by putting the symbol next to the respective compounds/elements

• solid state (s)

• liquid state (l)

• aqueous solution [solution made in water] (aq)

• gaseous state (g)

method of indicating precipitate in a chemical reaction

• since the precipitate is a solid substance, it is indicated by (s)

• (optional) mention [(insert color) ppt.] below the respective compound

method of indicating heat changes in an equation

• for exothermic reactions: ‘+ heat’ or ‘+ heat energy’ is to be written on the RHS

• for endothermic reactions: a delta symbol (triangle) is to be drawn over the arrow

method of indicating conditions under which the reaction takes place

• when ‘heated in the presence of X’ appears in the word equation, then X is called the catalyst of the reaction, and is written below the arrow mark

• other conditions such as atmospheric pressure and temperatures can also be written in the same way

chemical formula of flame in gas stove

CH4 (methane) (g) + O2 (g) → CO2 (g) + H2O (water vapour) (g) + heat

main oxides formed by iron

• ferrous oxide [iron (ii) oxide]: FeO

• ferric oxide [iron (iii) oxide]: Fe2O3

• magnetic iron oxide [mixture of iron (ii) and (iii) oxides]: Fe3O4

reason why respiration is considered an exothermic process

because the CO2 in our food break down into glucose, which reacts with oxygen in our cells and releases a large amount of energy

conditions under which methanol [aka methyl alcohol] is manufactured from carbon monoxide and hydrogen

• at 300 atm (atmospheric pressure)

• catalysts: zinc oxide + chromium oxide (ZnO + CrO3)

physical state of HCl

HCl can be read in two ways, which differ in states;

• hydrochloric acid: aqueous

• hydrogen chloride: gaseous

process of electrolysis of water

it is in which heat is given to the water which gets absorbed to give oxygen and hydrogen

H2O + heat → H2 + O2

substance used for testing carbon dioxide gas

lime water [Ca(OH)2 (aq)]

uses of slaked lime/ calcium hydroxide

• aqueous solutions of this are used for whitewashing walls. it reacts with carbon dioxide in air to form a thin layer of CaCO3 on the walls

• it is also the formula of marble, which is used for construction of buildings

use of magnesium hydroxide

used as antacid to relieve indigestion

list of types of chemical reactions

• combination reaction

• decomposition reaction

• displacement reaction

• double displacement reaction

• oxidation and reduction reaction

combination reaction

those reactions in which two or more substances combine to form a single substance are called combination reaction

A + B → AB

examples of combination reaction

• H2 (g) + O2 (g) → H2O (l)

• Mg (s) + O2 (g) → MgO (s)

• C (g) + O2 (g) → CO2 (g)

• H2 (g) + Cl2 (g) → HCl (g)

• Na (s) + Cl2 (g) → NaCl (s)

• Fe (s) + S (s) → FeS (s)

• CaO (s) + H2O (l) → Ca(OH)2 (s)

• NH3 (g) + HCl (g) → NH4Cl (s)

uses and connection between calcium oxide, calcium hydroxide and calcium carbonate

• calcium oxide = quicklime = CaO

• calcium hydroxide = slaked lime = Ca(OH)2

• calcium carbonate = limestone/chalk/marble = CaCO3

the calcium oxide is mixed with water to make calcium hydroxide (s), which is made into a solution and used to whitewash walls. after its applied to the walls, it reacts with CO2 in air to form a thin layer of calcium carbonate on the walls (this is formed after 2-3 days, and gives the walls a shiny finish)

decomposition reaction

those reactions in which a compound splits up into two or more simpler substances are known as decomposition reaction

AB → A + B

thermal decomposition

they are displacement reactions in which the catalyst is heat

examples of thermal decomposition reaction

• CaCO3 (s) → (heat) CaO (s) + CO2 (g)

• 2KClO3 (s) → (heat) 2KCl (s) + 3O2 (g) [this reaction is used for making oxygen in labs]

• 2FeSO4 (green) (s) → (heat) Fe2O3 (brown) (s) + SO2 (g) + SO3 (g)

• 2Pb(NO3)2 (colorless) → (heat) 2PbO (yellow) + NO2 (brown fumes) + O2

ferrous sulphate heptahydrate (FeSO4.7H2O)

• these are the ferrous sulphate crystals that we actually find in labs

• they contain 7 waters of crystallisation, which are green in color. actual FeSO4 is white in color

electrolytic decomposition

they are decomposition reactions in which electricity is the catalyst

electrolysis of water

-

examples of electrolytic decomposition

• 2NaCl (molten) → (electricity) 2Na + Cl2

• 2Al2O3 (molten) → (electricity) 4Al (l) + 3O2

photolytic decomposition

they are decomposition reactions in which photo (light) is the catalyst

examples of photolytic decomposition

• 2AgCl (white) (s) → (light) 2Ag (greyish white) + Cl2 (yellowish green)

• 2AgBr (pale yellow) → (light) 2Ag (greyish white) + Br2 (red-brown)

NOTE: both these reactions are used in black and white photography

uses of decomposition reaction

• electrolytic decomposition is used to extract metals from oxygen or chloride compounds

• when the molten compound is decomposed by passing electricity, the metal is produced at cathode

decomposition reaction in our body

• digestion of food in our body is a decomposition reaction

• when we eat food like wheat, rice or potatoes, the starch present in it decomposes to form simple sugars like glucose. similarly, the proteins decompose to form amino acids

reactivity series

• it is an arrangement of metals from most reactive to least reactive;

K Na Ca Mg A Zb Fe Sn Pb [H] Cu Hg Ag Au

displacement reactions

• they are reactions in which one element takes the place of another element in a compound. it occurs because more reactive elements take the place of less reactive ones

• AB + C → BC + A

examples of displacement reactions

• CuSO4 (blue) (aq) + Zn/ Mg (silvery white) → Zn/MgSO4 (aq) (colorless) + Cu (red-brown)

• CuSO4 (blue) (aq) + Fe (grey) → FeSO4 (greenish) (aq) + Cu (red-brown)

• CuCl2 (green) (aq) + Pb (blueish gray) → PbCl2 (colorless) + Cu (red-brown)

• 2AgNO3 (colorless) (aq) + Cu (red-brown) → Cu(NO3)2 (blue) (aq) + 2Ag (greyish white)

• 2HCl (aq) + Fe/ Mg → Fe/MgCl2 (aq) + H2

• 2Na + 2H2O → 2NaOH (aq) + H2

• Cl2 + 2KI (aq) → 2KCl (aq) + I2

• CuO + Mg → MgO + Cu (all s)

• Fe2O3 + 2Al → Al2O3 + 2Fe (molten)

double displacement reaction

• they are reactions in which two compounds exchange elements to form new compounds

• it usually occurs in solutions and one of the products, being insoluble, separates out

• AB + CD → AC + BD

examples of double displacement reaction

• AgNO3 + NaCl → NaNO3 (aq) + AgCl (white ppt.)

• BaCl2 + Na/CuSO4 → 2Na/CuCl (aq) + BaSO4 (white ppt.)

• CuSO4 + H2S (g) → H2SO4 + CuS (black ppt.)

• AlCl3 + 3NH4OH → 3NH4Cl + Al(OH)3 (white ppt.)

• Pb(NO3)2 + 2KI → 2KNO3 + PbI2 (yellow ppt.)

oxidation

• the addition of oxygen/ non metallic element is called oxidation

• the removal of hydrogen/ metallic element is also called oxidation

reduction

• the addition of hydrogen/ metallic element is called reduction

• the removal of oxygen/ non-metallic element is also called reduction

oxidising agent

• substances which give oxygen are called oxidising agents

• substances which remove hydrogen are also called oxidising agent

• the substance which gets reduced is the oxidising agent

reducing agent

• substances which give hydrogen are called reducing agents

• substances which remove oxygen are also called reducing agents

• the substance which gets oxidised is the reducing agent

examples of redox reaction

• CuO + H2 → (heat) Cu + H2O

  • oxidised substance: H2

  • reduced substance: CuO

  • oxidising agent: CuO

  • reducing agent: H2

• H2S + Cl2 → S + 2HCl

  • oxidised substance: H2S

  • reduced substance: Cl2

  • oxidising agent: Cl2

  • reducing agent: H2S

• MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O

  • oxidised substance: HCl

  • reduced substance: MnO2

  • oxidising agent: MnO2

  • reducing agent: HCl

• 2Cu + O2 → (heat) 2CuO

  • oxidised substance: Cu

  • reduced substance: O2

  • oxidising agent: O2

  • reducing agent: Cu

reaction used in production of zinc metal in industry

ZnO + C → Zn + CO

corrosion

• it is the process in which metals are eaten up gradually by the action of air, moisture, acids etc

• it weakens steel objects and structures such as railings, bridges etc., and cuts short their life

rusting

• it is action of corrosion on iron

• it is the most common form of corrosion

• during the corrosion of iron, iron metal is oxidised by the oxygen in air in the presence of moisture (water) to form hydrated ferric oxide called rust

  • 4Fe + 3O2 + 2xH2O → 2Fe2O3.xH2O

• rust is a soft, porous substance that gradually falls off the iron object

rancidity

• it is the condition produced by aerial oxidation of fats and oils in food marked by unpleasant smell and taste

• when the fats and oils present in food gets oxidised by air, the oxidation products have an unpleasant smell and taste

prevention of rancidity

• rancidity can be prevented by adding anti-oxidations to food: they prevent food from getting oxidised easily

• it can be prevented by packing food in nitrogen gas: as nitrogen is an unreactive gas, oxygen will not be able to enter

• it can be prevented by keeping food in a refrigerator: the low temperature slows oxidation

• it can be prevented by storing food in air-tight containers: it prevents exposure to oxygen

• it can be prevented by storing foods away from light: oxidation is slowed down in the absence of light