Comprehensive WBBSE Class 9 Physical Science Notes Chapter 4 Matter: Atomic Structure, Physical and Chemical Properties of Matter can help students make connections between concepts.
Matter: Atomic Structure, Physical and Chemical Properties of Matter Class 9 WBBSE Notes
Indian Philosopher Kanad named the basic smallest particles of matter as Kana.
Greek Philosopher Demokritos called the smallest particle as Adornos’ meaning indivisible.
John Dalton developed the idea of ultimate unit of matter which is known as atom.
Cathode Rays : It consists of negatively charged material particles called electrons. Cathode rays are produced from cathode surface when a gas at low pressure is subjected to electric discharge.
Electrons : Fundamental sub-atomic particles carrying negative charge (1.602 × 10-19 coulombs) and having mass 91 × 10-31 kg, discovered by J. J. Thomson.
Proton : A sub-atomic positively charged particle, having charge 1.602 × 10-19 coulombs and mass 1.67 × 10-27 kg. Mass of a proton is nearly 1837 times higher than the mass of an electron.
Alpha particle : He2+ ions or helium nuclei.
Rutherford’s Experiment : Led to the discovery of nucleus. Radius of nucleus (~ 10-15 m) is very small as compared with radius of an atom (~ 100-10m).
Neutrons : Sub-atomic neutral particles having mass 1.675 x 10-27 kg, discovered by James Chadwick.
Nucleons : The constituent particles of the nucleus i.e., protons and neutrons together are called nucleons.
The mass of an electron is \(\frac{1}{1837}\) times the mass of a proton.
The mass of an electron is \(\frac{1}{1839}\) times the mass of a neutron.
Rutherford’s model of the atom proposed that a very tiny nucleus is present inside the atom and electrons revolve round this nucleus. The stability of the atom could not be explained by this model.
Neils Bohr proposed that electrons revolve round the nucleus in a limited number of orbits, called permissible orbits. The orbits are also called energy level.
Energy levels are designated by letters K, L, M, N, O, P, Q.
Elements with their outermost shell completely filled are chemically inert.
The outermost shell of an atom is called valence shell and the electrons present in the valence shell are called valence electrons.
The maximum number of electrons that can be accommodated in a shell is given by 2n2, where
n = shell number.
Atomic number (Z)= Number of protons
= Number of electrons
– Mass number (A) = Number of protons + Number of neutrons
Isotopes : Atoms of same element having same atomic number but different mass number.
Isobars : Atoms of different elements having same mass number but different atomic number.
Isotones: Atoms of different elements having same neutron number but different atomic and mass number.
Nuclear force : In the nucleus, continuous transformation of the closely packed neutrons and protons take place i.e. protons transform into neutrons and vice versa. As a result of this transformation a strong attractive force acts within the nucleus. This force is known as nuclear force.
Ion : A charged atom or radical is called an ion.
Cation : Positively charged atom or radical is called a cation.
Anion : Negatively charged atom or radical is called an anion.
Frequency (v) of the radiation emitted when an electron jumps from an orbit having energy E1 to an orbit having energy E1 is given as \(v=\frac{E_2-E_1}{h}\)
Electronic arrangement of elements:
Each element has its own characteristic emission spectrum.
Series of lines in the line spectrum of hydrogen atom :
Orbital: Region of space around the nucleus where probability of finding an electron is maximum.
Quantum Mechanics: It is a theoretical science that takes into account the dual nature of electron.
Degenerate orbitals: The orbitals having equal energy.
Ground state: The lowest energy state of an atom.
Excited state: An atom is said to be in excited state if some lower energy orbital is vacant and the electron is present in higher energy orbital.
Electronic configuration : Distribution of electrons among various orbitals in an atom.
Quantum numbers : A set of four number used to specify energy, size, shape, orientation of the electron orbital and spin of the electron.
Principal Quantum number (n) : Tells us about energy of the electron and size of the orbital.
Azimuthai Quantum number (1): Tells us about the shape of the orbital and orbital angular momentum.
Magnetic Quantum number (m): Tells us about orientations of the electron cloud in a sub-shell.
Spin Quantum number (s) : Tells us about spin of the electron.
s-orbitais are spherically symmetrical.
Shape of s-orbitais:
p-orbitals are dumb-bell shaped:
The five d-orbitais are : dxy dyz dxz dx2 – y2 and dz2
Angular momentum \(m v r=\frac{n h}{2 \pi}\)
[m mass of the electron, u = velocity of the electron, r = radius of the orbit,
h = Planck’s constant, n = Principal Quantum number]
de-Brogue equation:
\(\lambda=\frac{h}{m v}\) A = [λ = wavelength, m = mass of the electron, υ = velocity of the electron, h = Planck’s constant]
Heisenberg’s uncertainty principle:
\(\Delta x \cdot \Delta p \geq \frac{h}{4 \pi}\)
[Δx = uncertainty in position,
Δp = uncertainty in momentum
h = Plancks constant]
Schrodinger’s wave equation:
[m = mass of an electron, V potential energy, E = total energy, Ψ = wave function of electron (eigen function
Aufbau Prinuiple : The orbitais are filled in the order of increasing energy,starting with the orbital of lowest energy.
Pauli’s Exclusion Principle: No two electrons in an atom can have same set of ail the four quantum numbers.
Hund’s Rule of Maximum Spin Multiplicity: The pairing of electrons in orbitais of a particular sub-shell cannot take place until all the orbitais of the sub-shell are singly occupied.
Nuclear isomer : Atoms (nuclides) which have the same atomic number and mass number but have different radioactive properties are called nuclear isomers.
Example:
Isodiapher : Atoms in which the difference between the number of neutrons and the number of protons is saine are termed as isodiapher.
Example:
Isoster: Molecules or ions with same number of atoms and also the same number of electrons are said to form isosteric group or more simply isosters.
Example : in N2O and CO2, number of atoms 3 and numbers of electrons = 20. So, these are isosters.
Gay Lussac’. Law : Under the same conditions of temperature and pressure, when two or more gases combine together. they do so in simple ratio by volumes, and the volumes of the products if gaseous also bear a simple ratio to the volume of the reacting gases.
Berzelius hypothesis: Under the same conditions of temperature and pressure, equal volume of all gases contain the same number of atoms.
Avogadros hypothesis : Under the same conditions of temperature and pressure equal volume of all gases (both elementary and compound) contain equal number of molecules.
Molecule : The smallest particle of an element or compound which can exist in the free state is known as molecule.
Elementary molecule : The molecule which is formed by one or more atoms of an element is known as elementary molecule. Example : Hydrogen (H2), Oxygen (O2) etc.
Compound molecule : The molecule which is formed by atoms of more than one element is known as compound molecule. Example: Water (H2O), Carbon dioxide (CO2) etc.
Vapour density: Vapour density of a gas is the ratio of the weight of a certain volume of the gas to the weight of same volume of hydrogen under similar conditions of temperature and pressure.
Molecular mass =
Atomic mass = \(\frac{\text { mass of } 1 \text { atom of an element }}{\frac{1}{12} \times \text { mass of } 1 \text { carbon }\left({ }^{12} \mathrm{C}\right) \text { atom }}\)
Gram atomic mass : When the atomic mass of an element is expressed in gram then the amount of element in gram is known as gram atomic mass of the element.
Gram molecular mass : When the molecular mass of an element or compound is expressed in gram then that amount of element or the compound in gram is known as gram molecular mass of the element or the compound.
Molar volume : The volume of a gaseous substance (element or compound) under a fixed temperature and pressure is known as molar volume of the gas.
Avogadro’s number: One gram-molecule of any element or compound containing equal number of molecules (NA 6.022 × 1023) is known as Avogadro’s number.
R. A. Millikan determined the value of Avogadro’s number by ‘oil drop experiment’ in 1913.
Modern definition of Avogadro’s number : The number of atoms present in exactly 12g of carbon (12C) is designated as Avogadro’s number.
Avogadro’s constant : ‘Avogadro’s number/mole’ is called Avogadro’s constant
i.e. 6.022 × 1023 Avogadro’s number has no unit. But Avogadro’s constant has the unit mol’. It is a universal constant.
Atomicity : It is the number of atoms by which an elementary molecule is composed of.
Atomic mass unit (amu) : It is an unit used for measusing the atomic mass of an atom.
1 amu = 16606 ×10-24g
Average atomic mass
\(\frac{\sum \text { natural abundance of isotope }(\%) \times \text { its atomic mass }}{100}\)
Mole (SI System) : A mole is the amount of a substance which contains as many entities (atoms, molecules, ion or any other particle) as there are atoms in exactly 0012 kg (or 12g) of carbon-12 isotope.
Solution : A solution (or true solution) is a uniformly homogeneous mixture of two or more substances whose relative proportions may be varied upto a certain limit.
Solute : The substances which are present in smaller quantities and gets dissolved are called solutes.
Solvent : The medium in which the solutes are uniformly dispersed through dissolution is known as the solvent. solute + solvent = solution
Homogeneous mixture : A homogeneous mixture is one in which the constituent substances are so intimately mixed that even a very close examination cannot distinguish any surface of separation between them.
Heterogeneous mixture : A heterogenous mixture on the other hand is one in which the particles of the constituent substances of a mixture are distinctly distinguishable.
Some useful solvents : Water is a versatile liquid solvent of organic and inorganic substances, carbon disuiphide is the solvent of sulphur, oxalic acid is the solvent of rust, benzene is the solvent of rubber etc.
Types of solution:
(a) Solution of solid in solids : Solution of one metal in another gives the alloys like brass, bronze etc.
(b) Solution of liquids in solids : Mercury (liquid metal) can remain dissolved in solid metal like gold, silver, sodium etc. to form amalgams.
(c) Solution of gases in solids : Hydrogen gets dissolved in spongy palladium and in some similar metals. This phenomenon is called occlusion.
(d) Solution of solids less in liquids : Aqueous solution of common salt (NaCl), sugar, copper sulphate are examples.
(e) Solution of liquids in liquids : A mixture of water and alcohol, a mixture of benzene and xylene are examples.
(f) Solution of gases in liquids: Ordinary water contains air dissolved in it. Soda water contains carbon dioxide under pressure. Dissolved oxygen in water allows aquatic life to survive.
(g) Solution of solids in gases : e.g. smoke, where solvent is air and solute is carbon and dust particles dissolved in air.
(h) Solution of liquids in gases: e.g. fog.
(j) Solution of gases in gases : e.g. air, mainly a mixture of nitrogen and oxygen.
Classification based on concentration:
- saturated solution
- unsaturated solution
- supersaturated solution
Suspension : The particle in suspension state is 10-4cm in diameter.
True solution : The particle in true solution is 10-8 cm in diameter.
Sol : It is a colloidal solution in which a solid is dispersed in a liquid (e.g.paints).
Emulsion : It is a colloidal solution in which a liquid is dispersed in another liquid (e.g. milk).
Gel : It is a colloidal solution in which a liquid is dispersed in a solid (e.g.fruit jellies, cheese etc.)
Aerosol : It is a colloidal solution in which a solid or a liquid is dispersed in a gas (e.g. fog. smoke, clouds etc.)
Lyophilic colloids : These are the substances which pass readily into the colloidal state whenever mixed with a suitable solvent (e.g. protein, starch etc.)
Lyophobic colloids : These are the substances which do not yield colloidal solutions on mere shaking with a liquid (e.g. gold, silver. Fe(OH)4, As).
Positive sols : These are the sols which carry positive charge on the dispersed phase particles. (e.g. sols of Fe(OH), Al(OH)3 etc.)
Negative sols : These are the sols which carry negative charge on their particles (e.g. sols of Cu, Ag, Au, As2S3 etc.)
Multi-molecular colloids :These are the colloids in which the individual particles consist of aggregates of atoms of small molecules having molecular sizeless than 10-7cm in diameter. (e.g. sols. of gold atoms, platinum sol.)
Macro-molecular colloids : These are the colloids in which the size of the particles of the dispersed phase are of the order of colloidal dimensions (e.g. sol of starch, cellulose)
Peptisation : It is the process of converting precipitates into colloidal state by adding small amount of a suitable electrolyte.
Dialysis : It is the process of separating substances in colloidal state from those present in ionic states with the help of a semipermeable membrane
Tyndal I effect : It is the process of scattering of light from the surface of colloidal particles and in this process a beam of light passed through a colloidalsolution becomes visible as a bright streak.
Brownian motion : This is ceaaeless erratic irregular and random motion of colloidal particles suspended in a dispersion medium.
Elcctrophoresis : It is the movement of charged colloidal particles, under the influence of an electric field, towards the oppositely charged electrodes.
Coagulation : It is the phenomenon of change of colloidal state to suspended state. It can be brought about by adding an electrolyte to a colloidal solution.
Hardy schuize rule : It states that greater the valency of the coagulating ion added, the greater is its power to coagulate.
Gold number : ¡t is the weight in milligrams of a protective colloid which prevents the coagulation of 10 mL of a gven gold sol on adding 1 mL of 10% solution of sodium chloride.
Smaller the value of gold number, of a lyophilic colloid, the greater is its protective power.
Colloids in everyday life :
(i) Foods, medicines and pharmaceutical preparation.
(ii) Soaps, synthetic detergents,
(iii) Paints, varnishes, enamels, resins, gums, glues, rubber.
(iv) Industrial processes such as tanning, dyeing, rubber plating and removal of smoke from air are based on the colloidal nature of particles.
Suspended particulate matter (SPM) : Many partículate matter, such as fine particles of dust, sand, coal dust, pollen grains, cement dust, metal dust, fly-ash etc. remain suspended in normal air; these are termed as pollutants which cause air pollution.
Solubility : Solubility of a given solute in a solvent is defined as the weight in grams of the solute dissolved in 100 grams of the solvent so as to saturate the solution at a given temperature.
Solubility is just a number, it has no unit,
Effect of temperature and pressure on soluhility of gas in liquid:
(a) Effect of tempe rature : The solubility of gases in liquids decreases with the increase of temperature and at lower temperature, solubility increases.
(b) Effect of pressure : At a definite temperature the solubility of a gas in a particular solvent is proportional to the pressure applied on the gas. This is known as Henry’s law.
Effect of temperature Of solubility of solid in liquid :
- The solubility of KNO3 in water, increases tremendously with the increase of temperature.
- In case of NaCl, the increase of solubility in water is much less, almost same.
- The solubility of CaSO4in water decreases with the increase of temperature.
- The solubility of Glauber’s salt (Na2SO4.104O) in water first increases upto a certain temperature (324°C) and then sharply decreases with temperature.
Solubility curve :
Crystals : Crystals are homogeneous solid particles (big or small) with definite geometric shape, and are bounded by plane surfaces which meet at sharp edges.
Crystallisation : Crystallisation is a process by which crystals of a substance are obtained from its solution.
Water of crystallisation : When crystal is formed from the aqueous solution of it then one or more fixed number of water molecules associated with each molecules of that substance by chemical bonding are called water of crystallisation.
Hydrated crystals: Crystals having the water molecules associated with them are called hydrated crystals.
Example:
MgSO4. 7H2O (Epsome salt)
FeSO4. 7H2O (Green vitriol)
CuSO4. 5H2O(Blue vitriol)
H2C2O4. 2H<2O (Oxalic acid)
Efflorescent substance and Efflorescence: Certain hydrated crystals when exposed to air at ordinary temperature lose their water of crystallisation partially or fully and are transformed into amorphous varieties. Such a substance is called efflorescent substance and the phenomenon is known as efflorescence.
Example:
Cause of efflorescence: Efflorescence is happened only when the vapour pressure within the hydrated crystal at ordinary temperature is greater than the vapour pressure of the atmosphere.
Deliquescent substance and Deliquescence: Certain water soluble substances when exposed to the atmosphere at ordinary temperature absorb water vapour from the air and are soluble in it to make saturated solution, such substances are called deliquescent substances and the phenomenon is known as deliquescence.
Example:
(i) MgCl2 (Magnesium chloride)
(ii) CaCl2(Calcium chloride)
(iii) FeCl3 (Ferric chloride)
(iv) NaOH (Caustic soda)
N.B. : Edible Salt (NaCl) becomes moist in the rainy season due to the presence of deliquescent MgCl2. 6H2O as impurity in it.
Cause of deliquescence: Deliquescence occurs when the vapour pressure of water in the delequescent substance is less than the vapour pressure of the atirosphere of ordinary temperature.
Hygroscopic substances: There are certain substances which when exposed to air absorb moisture from air but are not dissolved in the absorbed water. These substances are called hygroscopic substances. Example: Conc. HSO4, CaO etc.
Drying agent: There are certain substances which absorb water or water vapour molecule from other compounds without forming any chemical bond. These substances are called drying agent.
Example: Anhydrous CaCl2, Anhydrous MgSO4, Conc. HSO4 etc.
The concentration of a solution can be expressed in several different ways:
(a)Percentage by weight (% w/w): It is the amount of solute in grams present in 100 grams of the solution.
(b) Weightivolume percentage (% w/V): It is the amount of solute in grams present in 100 mL of the solution.
(c) Volume/volume percentage (% V/V): It is the volume in mL of the solute present in 100 mL of the solution.
(d) Strength: It is the number of grams of the solute dissolved per litre of the solution.
(e) Normality (N): It is the number of gram equivalent of solute dissolved per litre of a solution.
(f) Molarity (M): It is the number of moles of solute dissolved per litre of the solution.
(g) Molality (m): It is the number of moles of solute dissloved in 1000 gram of the solvent.
(h) Formality (F): Formality of a solution is the number of gram formula weight of the ionic solute (e.g. NaCl) dissolved per litre of the solution.
(j) Parts per million (ppm) : It is defined as the number of parts of a component per million parts of the solution.
Vapour pressure of liquid : The pressure developed above the liquid at a given temperature at the equilibrium point.
Lowering of vapour pressure : It is the difference in the vapour pressure of the pure solvent and that of the solution.
Raoult’s Law : The vapour pressure of a solution is equal to the product of the mole fraction of the solvent and its vapour pressure in pure state.
Ideal solution: The solution which obey Raoult’s law at all concentration and follow the conditions
\(\Delta \mathrm{H}_{\text {mix }}=0 ; \Delta \mathrm{V}_{\text {mix }}=0 \)
Non-ideal solution : The solution which show positive or negative deviations from Raoults law.
Azeotrope : The mixture of liquids which boils at constant temperature like pure liquid and has same composition of components in liquid as well as vapour phase.
Colligative properties : The properties of the solution which are independent of nature of solute but depend upon the concentration of solute particles.
Osmosis : The passage of solvent from pure solvent or solution of low concentration to the solution of higher concentration through the semi permeable membrane.
Isotonic solutions : The solutions of same molar concentration and same osmotic pressure at given temperature.
Acids:
(j) Arrhenius concept : A substance which furnishes H ions in aqueous solution.
(ii) Bronsted concept : A substance which is proton donor.
(iii) Lewis concept : A substance which is acceptor of electron pair.
Types of acids : According to their structure:
(i) Hydracids : Those acids which have hydrogen atoms and other non-metal or radicals but no oxygen atom are called hydracids. Example : Hydrochloric acid (HCl), Hydrocyanic acid (HCN).
(ii) Oxyacids : Those acids which have one or more oxygen atoms with hydrogen and other non-metal atoms are called oxyacids. Example : Nitric acid (HNO3), Sulphuric acid (H2SO4).
Types of oxyacids:
(a) Ous-acids : Those oxiacids which have lesser number of oxygen atom and the valency of the main element naming the acid is less are called otis-acids.
Example: HNO2 (Nitrous acid)
[valency of N = 3]
H2SO3 (Sulphurous acid)
[valency of S = 4]
(b) Ic-acids : Those oxyacids which have greater number of oxygen atom and the valency of the main element naming the acid is high. are called ic-acids.
Example: HNO3 (Nitric acid)
[valency of N= 5]
H2SO4 (Sulphuric acid)
[valency of S = 6]
(c) Hypoacids : Those acids which have less oxygen atom than ous-acids are called hypoacids.
Example : HOCl (Hypochiorous acid)
(d) Per-acids : Those acids which have greater number of oxygen atoms than ic-acids are called per-acids. Example : HClO4 Perchloric acid)
Basicity of acid : The number of H⊕ ions produced in aqueous solution of a molecule of acid, is called basicity of acid.
Types of acids according to basicity of acid:
(a) Monobasic acids : Those acids which ionise in aqueous solution to produce one He ion are called monobasic acids. Example : HCl, HNO3, HBr etc.
(b) Diba sic acids : Those acids which ionise in aqueous solution to produce inherit; two He ions are called dihasic acids. Example : H2SO4, H2CO3, H2SO3 etc.
(c) Tribasic acids : Those acids which ionise in aqueous solution to produce three He ions are called trihasic acids. Example : H3PO4
• Types of acids according to their ionisation :
Strong acids : The acids which ionise mostly in aqueous solution and small number of molecules exist in molecular state, are called strong acids.
Example : HCl, HNO3, H2SO4 etc.
Weak acids : The acids whose most of the molecules exist in molecular state and ionise a little number of molecules, are called weak acids.
Example : CH3COOH, H2CO3, H2S etc.
Types of acids according to their source :
Mineral acids : The acids which are produced from minerals are called mineral acids.
Example :
(i) Hydrochloric acid (HCl)
[Prepared from NaCl]
(ii) Nitric acid (HNO3)
[Prepared from KNO3]
Organic acids : The acids containing carbon atom and prepared from animals or plants are known as organic acids.
Example :
(i) Formic acid (HCOOH) [Source : Ant]
(ii) Lactic acid [CH3CH(OH)COOH] [Source : milk]
Properties of acids :
(i) An acid produces H ⊕ ion in aqueous solution.
(ii) Acids are generally soluble in water and the solution is acidic or sour in taste.
(iii) The aqueous solution of acids turns blue litmus into red, orange coloured methyl orange into pinkish red. On the other hand, phenol phthalein remains colourless in aqueous solution of an acid.
(iv) Electropositive metals like Zn, Mg, Fe (which belong above hydrogen in electrochemical series) react with dilute acid to produce hydrogen.
Example :
Zn + H2SO4 = ZnSO4 + H2↑
2Al + 6HCl – 2AlCl3 + 3H2↑
Acids form salt and water reacting with metallic oxides and hydroxides.
Example :
2NaOH + H2SO4 = Ha2SO4 + 2H2O
CaO + 2HCl = CaCl2 + H2O
Acids form carbon dioxide reacting with metallic carbonate and bicarbonate.
Example :
Na2CO3 + 2HCl = 2NaCl + CO2 + H2O
NaHCO3 + HCl = NaCl + CO2 + H2O
Fuming nitric acid : It is a product produced by dissolving excess amount of nitrogen dioxide in nitric acid.
Acid rain : Dissolved NO2 and SO2 of atmosphere produce nitric acid (HNO3) and sulphuric acid
(H2SO4)which come down along with rain. This phenomenon is known as acid rain.
Stone cancer : Corrosion of marble or stone due to acid rain is called stone cancer.
Aqua regia : It is a mixture of 3 volumes concentrated hydrochloric acid and 1 volume concentrated nitric acid which can dissolve metals like gold and silver.
Royal water : Aqua regia is called royal water.
Fuming sulphuric acid : Fuming sulphuric acid or oleum (H2S2O7) is obtained when sulphur trioxide is passed over 98% Sulphuric acid.
Muriatic acid : Hydrochloric acid.
Aqua fortis : Nitric acid.
King of chemicals : Sulphuric acid.
Laughing gas : Nitrous oxide (N2O).
Preparation of hydrogen chloride in laboratory :
Preparation of nitric acid in laboratory :
Manufacture of sulphuric acid by contact process :
Identification of acids :
(a) Identification of Hydrochloric acid :
(b) Identification of Sulphuric acid :
(c) Identification of Nitric acid (Ring test) :
Formation of Hydrogen from nitric acid : Magnesium liberates hydrogen from cold and very dilute nitric acid.
Mg + 2HNO3 = Mg(NO3)2+ H2
Passive iron : Cold and concentrated nitric acid or fuming nitric acid when comes in contact with iron produces passive iron. Passive iron is chemically inactive.
Uses of hydrochloric acid :
- It is used in dyeing and calicoprinting.
- It is used as a cleaning agent in galvanising and tin plating.
- It is used in the manufacture of glucose, glue and many useful metallic chlorides.
Uses of nitric acid :
- It is used in the production of explosives such as dynamite, gun cotton, nitro-glycerine, trinitro toluene (TNT).
- It is used to prepare rayon, artificial silk and dyes.
Uses of sulphuric acid :
- It is used for the preparation of different chemical compounds like HCl, HNO3, ether, alcohol etc.
- It is used as an important raw material for dyes, medicine, plastic etc.
- Oxides : Binary compound of oxygen with any element (metallic and non-metallic) are called oxides.
Types of Oxides :
(a) Acidic oxide : An acidic oxide is an oxide which reacts with a base to form salt and water.
- Non-metallic oxide : CO2, SO2, P2O5, SiO2
- Metallic oxide : CrO3 (Chromium trioxide), Mn2O7 (Manganese heptoxide) etc.
(b) Basic oxide: A basic oxide is an oxide which reacts with an acid to form salt and water. These are generally the oxides of metals.
Example :
CaO (Calcium oxide)
MgO (Magnesium oxide)
FeO, CuO etc.
(c) Neutral oxide : The oxides which neither react with acid nor with base are called neutral oxide.
Example : CO (Carbon monoxide)
H2O (Water), N2O (Nitrous oxide)
NO (Nitric oxide)
(d) Amphoteric oxide : The oxides which react with both acids and bases to produce salt and water are called amphoteric oxide.
Example :
Al2O3 (Aluminium oxide)
ZnO (Zinc oxide)
SnO (Stanous oxide)
PbO (Lead monoxide) etc.
Peroxide : The oxides which react with cold and dilute mineral acid to produce hydrogen peroxide and have peroxylinkage (— O — O —) are called peroxide.
Example :
Na2O2 (Sodium peroxide)
BaO2 (Barium peroxide)
Mixed oxide : An oxide which is formed by oxides of more than one oxide of a metal having variable valency is called mixed oxide.
Example :
Fe3O4 [FeO + Fe2O3] ; Mn3O4 [2MnO + MnO];
Pb3O4 [PbO2 + 2PbO] etc.
Poly-oxide : The oxides which have oxygen atoms more than an ordinary oxide and also do not react with acid to produce hydrogen peroxide are called poly-oxides.
Example :
Mn2O7 (Manganese heptoxide)
PbO2 (Lead dioxide)
Sub-oxide : An oxide which has less oxygen atoms rather than required for oxidation state of the element is called sub-oxide. Example : Carbon sub-oxide (C3O2)
Super-oxide : Metals like Na, Li, Ca form super-oxide and have negative ion (O-O) in it.
Example : LiO2, KO2 etc.
Alkali: Those metallic oxides and hydroxides react with acids to produce salts and water are known as alkali. Example : CaO, Na2O, NaOH, Al(OH)a etc.
Base (According to Arrhenius concept): The metallic oxides or hydroxides after dissolving in water produce negatively charged hydroxyl ions are known as base.
Properties of bases :
- The aqueous solution of bases turn red litmus into blue, orange coloured methyl orange into yellow.
- Bases produce salts and water when react with acids.
- Concentrated solutions of some strong bases are slippery to touch.
- The aqueous solution of bases can conduct electricity.
- Strong bases when heated with electro-positive elements (like Al, Zn etc.) produce hydrogen.
Example : 2NaOH + Al + 2H2O = 2NaAlO2 + 3H2↑
Strong bases absorb carbon dioxide from atmosphere to form carbonate salt and water.
Example : 2NaOH + CO2 = Na2CO3 + H2O
Difference between alkali and bases : The oxides and hydroxides of metal and metal like radicals are called alkalis. But those alkalis which dissolve in water to form OH are called bases.
Example :
- Acidity of bases : The number of OH ions produced in aqueous solution of one molecule of bases are called the acidity of bases.
Types of bases according to acidity of bases :
(a) Monoacidic bases : The bases which produce one OK– ion in aqueous solution from one molecule are called monoacidic bases.
Example:
(b) Diacidic bases : The bases which produce two OH” ion in aqueous solution from one molecule are called diacidic bases.
Example :
Types of bases according to their ionisation :
Strong bases : The bases which ionise mostly in aqueous solution and small number of molelcules exist in molecular state are called strong bases.
Example : NaOH, KOH, Ca(OH)2 etc.
Weak bases : The bases whose most of the molecules exist in molecular state and ionise a little number of molecules are called weak bases.
Example : NH4OH, Cu(OH)2 etc.
Types of bases according to their source :
Mineral bases : The bases which are produced from minerals are called mineral bases.
Example : NaOH, KOH etc.
Organic bases : The bases cortaining nitrogen atom and whose sources are animals and plants are called Organic bases.
Example:
Salts : The replaceable hydrogen atoms in an acid when replaced by metal or basic radical partially or fully then the compound so produced is called salt.
Definition of salts according to Arrhenius concept: The compound formed by cation other than hydrogen ion (H)+ and anion other than hydroxyl ion (OH~) is known as salt.
Types of salt :
- Normal salts : Normal salts are formed by the complete replacement of the hydrogen atoms in an acid by metals or basic redicals.
Example :
(i) NaCl [Sodium chloride] [NaOH + HCl = NaCl + H2O]
(ii) Na2SO4 [Sodium Sulphate] [2NaOH + H2SO4 = NA,SO4 + H2O]
Acid salt: An acid salt is one which is formed only by partial replacement of hydrogen atom in a molecule of an acid (di-basic, tri-basic) by metal or basic radical.
Example :
NaHSO4 (Sodium bisulphate)
NaHCO3 (Sodium bicarbonate)
Na2HPO4 (Di-sodium hydrogen phosphate)
Basic Salt: When the hydroxyl (OH) or, oxide (O) radicals are partially neutralised by acid then basic salts are produced.
Example :
(i) Pb(OH)Cl [Basic lead chloride]
(ii) Pb(OH)NO3 [Basic lead nitrate]
Double salts : When two normal salts are mixed with each other with their molecular weight ratio to form a solution and cooled the solution to make a combined crystal which is stable in solid state but ionises into different constituent ions in solution are called double salts.
Example :
K2SO4 Al2(SO4)3 . 24H2O
(NH4)2SO4FeSO4 . 6H2O
Complex salts : When the solution of two mixed salts is concentrated and cooled then a crystal of new salt is produced. The constituent salts lose their identity in new salt. In aqueous solution they produce new complex ions. This type of salts are called complex salts.
Example : K4[Fe(CN)6] ; [Cu(NH3)] SO4
Neutralisation : The reaction in which equivalent amount of an acid reacts with equivalent amount of a base and thus the properties of acid and bases are completely lost forming salt and water, is called neutralisation reaction.
Example :
Titration : The process by which bases are neutralised with the help of acids or vice-versa is called titration.
Indicators : The chemicals which are able to determine the end point of neutralisation reaction by changing their colours are called indicators.
Indicators used in neutralisation reaction :
| Name of the indicators and their actual colour | In acid solution | In alkali solution | In neutral solution |
| 1. Litmus (violet) | red | blue | violet |
| 2. Phenolphthalein | colourless | pink | colourless |
| 3. Methyl-orange (orange) | red | yellow | orange |
Choice of indicators :
(a) Strong acid and weak base — Methyl orange
(b) Weak acid and strong base — Phenolphthalein
(c) Strong acid and strong base — Any indicator
(d) Weak acid and weak base — No such indicator.
Vanishing colour : Vanishing colour is a coloured solution which becomes colourless on exposure to air after a lapse of time.
Example :
The aqueous solution of ammonium hydroxide (NH4OH) mixed with phenolphthalein turns pink colour. Now if this solution is allowed to spray on cloth, at once a pink colour is developed upon cloth. But since ammonia in NH4OH is highly volatile, after lapse of time NH3 disappears from the cloth and the solution becomes neutral. As a result pink colour of phenolphthalein vanishes.
Vanishing colour cannot be prepared with dilute NaOH solution, due to the fact that NaOH does not evaporate.
pH : The symbol pH is derived from ‘Poteuz,’ the Danish word for power, pH refers to potency of hydronium ion. pH = -log [H] = – log[OH–]
pH range :
- The pH of acidic solution is less than 7. [pH < 7, the solution will be acidic]
- The pH of basic solution is greater than 7. [ pH > 7, the solution will be basic]
- The pH of neutral solution is 7. [pH = 7, the solution will be neutral] So, the range of pH is from 0 to 14.
Colour changes for universal indicator at different pH values.
- The approximate pH of a solution can be determined with the help of pH papers.
- pH -papers have different colours in solutions of different pH.
A pH paper can determine pH of a solution with an accuracy of about 0 -5.
pH-metres : For accurate measurement of pH (upto accuracy 0.001 units), pH metres are used.
Universal indicator : A mixture of organic dyes which gives different colours with solutions of different
pH values is called a universal indicator.
| Name of the Fluid | pH | |
| 1. | 0 1(M) NaOH solution | 13 |
| 2. | Milk of magnesia | 10 |
| 3. | Egg white, sea water | 7-8 |
| 4. | Human blood | 7-4 |
| 5. | Tears | 7-4 |
| 6. | Milk | 6-8 |
| 7. | Human Saliva | 64 |
| 8. | Tomato juice | 4-2 |
| 9. | Soft drinks and Vinegar | 30 |
| 10. | Lemon juice | 2-2 |
| 11. | Gastric juice | 1.2 |
| 12. | 1(M) HCl solution | 0 |
Importance of pH :
| Field | Utility |
| 1. Biotechnology | Biochemical and organic reactions at controlled pH value give best results. |
| 2. Agriculture | Some crops such as citrus fruits grow better in alkaline soils. Sugarcane grows better in neutral soil, whereas rice grow better in slightly acidic soil. The pH of soil is tested before growing particular crop. |
| 3. Medicine | Determination of pH value of blood and urine helps to diagnose certain disease. |
| 4. Cosmetics | Soaps, shampoos, face creams are prepared for the consumers having different pH value of skin secretions. |
| 5. Milk plants | pH of milk is rigorously controlled at pH 6 8 as otherwise it turns sour. |
| 6. Breweries | Controlling pH value of wine to obtain a desired flavour. |
Antacids : Medicines which can remove the excess acid from the stomach and raise the pH to appropriate level are called antacids.
Types of antacids:
Cimetidine (Tegamet): It binds to the receptors that trigger the release of hydrochloric acid into the stomach. This results in release of lesser amount of acid.
Cimetidine remained the largest selling drug in the world for a long time.
(ii) Ranitidine (Zantac): After cimetidine, ranitidine (Zantac) was introduced for treatment of hyperacidity.
Separation of mixtures:
Fractional distillation : The process by which two or more miscible liquids are separated by distillation using the difference in their boiling point is called fractional distillation.
When the boiling points of two or more miscible liquids differ by about 15°C — 20°C these are separated by simple fraction distillation.
fractionating column : If the boiling points of two liquids are very close i.e. difference is small in that case a fractionating column is used to separate the liquid mixtures.
Separation of immiscible liquids using separating funnel : Immiscible liquids are easily separated by using a separating funnel. Each liquid forms its own liquid layer due to difference in density.
Chromatography : It is a method by which mixture of different substances are separated. In this method due to difference in adsorption of different substances in solid phase (adsorbent) and difference in migration of the substances in the mobile phase (liquid or gas), various substances are separated.
Paper Chromatography : It is a very easy technique to separate the various organic dyes present in the writing ink or printing ink.
Advantages of chromatography :
- A small amount of the compound present in the mixture can be separated.
- The properties of the individuals present in the mixture do not alter.
- It is a very easier method to separate the different components present in the mixture.
Sublimation : There are some solids which on heating directly transform into vapour phase without transforming through the intermediate liquid state. This phenomenon is called sublimation.
Example : NH4Cl (ammonium chloride) and SiO2 (sand) are separated by this method.
Natural water : Natural water, except rain water, is generally impure and contains dissolved metals, apart from other impurities.
Sources of natural water :
- Rain water
- River water
- Spring and well-water
- Mineral water
- Sea water
Soft water : Water in which lather is easily formed with soap is called soft water.
Hard water : The class of water that does not easily form lather with soap, is called hard water.
Hardness of water : The property of natural water which does not allow it to form lather is known as hardness of water.
Cause of hardness of water : The hardness of water is caused by the presence of bicarbonate, chloride and sulphate of calcium, magnesium and iron in water.
Types of hardness :
- Temporary hardness : It is due to the presence of soluble bicarbonate of calcium, magnesium and iron in water.
- Permanent hardness : This is caused by the presence of chloride and sulphate of calcium, magnesium and iron in water.
Removal of temporary hardness :
Removal of permanent hardness :
Lime-soda process :
CaSO4 + Na2CO3 = CaCO3↓ + Na2SO4
CaCl2 + Na2CO3 = CaCO3↓ + 2NaCl
Permutit process :
2Na (Permutit) + Ca(HCO3)2 = Ca(Permutit)2 ↓+ 2NaHCO3
[Zeolite, a natural mineral is a giant molecule of sodium aluminium silicate (NaAlSO4 3H2O). Artificially prepared mineral called permutit is used to remove permanent hardness]
Calgon process :
Ion exchange resin process :
De-ionised water : Very soft and pure water which is obtained by treating hard water through the ion-exchange process, is free from all types of ions i.e., cations and anions. Such type of water is known as de-ionised water.
Water is an universal and versatile solvent :
- Water is a polar compound. Most of the ionic compounds are soluble in polar solution and thus water is, therefore, a good ionising solvent for acids, bases and salts.
- Water is an amphiprotic solvent i.e. it can act both as an acid (a proton donar) and a base (a proton acceptor) in the reaction.
- Many covalent compounds like sugar, urea, alcohol, organic acid, organic bases are soluble in water through hydrogen bonding. Hence, we may conclude that water is an universal as well as versatile solvent.
Water pollution : Some soluble and insoluble matters in water which produce some harmful effect to men and aquatic animals are known as water pollution.
Pollution due to natural source :
Pollution due to arsenic : Ground water is sometime contaminated with arsenic acid and arsenius acid as a deprotonated form. Consumption of these arsenic contaminated water for a long period can cause dangerous poisonous effects in the form of black spots on palms and feet and skin become rough. This is known as black foot disease. From scientific researches it can be assumed that arsenic poison may be a cause of skin cancer.
Pollution due to fluoride : Sea-water, river-water contain fluoride from natural source. Excess fluoride in water is responsible for fluorosis disease which is a cause of dental decay and bone decay.
Pollution due to artificial or human activity : Fertilizers, insectisides, pestisides, are regularly used for increasing the production of food grains. These substances when washed with water are mixed in ponds, river etc. As a result of this, water is polluted.
Removal of arsenic from water :
- Arsenic is removed from water if water is allowed to pass through a bag made of cloth containing alumina or ferric hydroxide replacing the candle used in the filter. Arsenic or arsenius ions are absorbed in the alumina bed or ferric hydroxide bed and thus are removed.
- Take 20 litre of water, \(\frac{1}{4}\)th amount of tea spoon of bleaching powder and the same amount of ferrous sulphate in a bucket. After stirring well, arsenic will be removed.