Class IX, CHEMISTRY, "Atomic Structure"
Dalton’s Atomic Theory
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The important postulates of Dalton’s atomic theory are:
1. All elements are composed of atoms. Atom is too small so that it could not be divided into further simpler components.
2. Atom cannot be created nor destroyed in any chemical reactions.
3. Atoms of an element are similar in all respects. They have same mass and properties.
4. Atoms of different elements combine in a definite simple ratio to produce compounds.
Discovery of Electron
A discharge tube is a glass tube. It has two electrode, a source of electric current and a vacuum pump.
(Diagram)
Sir William Crooks (1895 performed experiments by passing electric current through gas in the discharge tube at very low pressure. He observed that at 10-4 (-4 is power to 10) atmosphere pressure, shining rays are emitted from cathode. These rays were named cathode rays. Cathode rays are material particles as they have mass and momentum.
Properties of Cathode Rays
The properties of these particles are given below:
1. These particles are emitted from cathode surface and move in straight line.
2. The temperature of the object rises on which they fall.
3. They produce shadow of opaque object placed in their path.
4. These particles are deflected in electric and magnetic fields.
5. These particles are deflected towards positive plate of electric field.
Discovery of Proton
Gold Stein (1886) observed that in addition to the cathode rays, another type of rays were present in the discharge tube. These rays travel in a direction opposite to cathode rays. These rays were named positive rays. By using perforated cathode in the discharge tube the properties of these rays can be studied. Positive rays are also composed of metered particles. The positive rays are not emitted from anode. They are produced by the ionization of residual gas molecules in the discharge tube. When cathode rays strike with gas molecule, electrons are removed and positive particles are produced.
Properties of Positive Rays
1. They are deflected towards negative plate of electric field. Therefore these rays carry positive charge.
2. The mass of positive rays is equal to the mass of the gas enclosed in the discharge tube.
3. The minimum mass of positive particles is equal to the mass of hydrogen ion (H+). These positive ions are called Protons.
4. The charge on proton is equal to +1.602×10-19 Coulomb. (-19 is power of 10)
Natural Radioactivity
The phenomenon in which certain elements emit radiation which can cause fogging of photographic plate is called natural radioactivity. The elements which omit these rays are called radioactive elements like Uranium, Thorium, Radium etc. There are about 40 radioactive elements. Henri Bequrel (1896) discovered radioactivity.Madam Curei also has valuable contribution in this field.
In natural radioactivity nuclei of elements are broken and element converted to other elements. Natural radioactivity is nuclear property of the elements.
Alpha Rays
1. They are helium nuclei. They are doubly positively charged, He2+.
2. They move with speed equal to the 1/10th of the velocity of the light.
3. They cannot pass through thick-metal foil.
4. They are very good ionizer of a gas.
5. They affect the photographic plate.
Beta Rays
1. They are negatively charged.
2. They move with the speed equal to the velocity of light.
3. They can pass through a few millimeter thick metal sheets.
4. They are good ionizer of a gas.
5. They can affect the photographic plate.
Gamma Rays
1. They are electromagnetic radiations.
2. They travel with speed equal to velocity of light.
3. They carry no charge.
4. They have high penetration power than alpha and beta rays.
5. They are weak ionizer of gas.
Rutherford Experiment and Discovery of Nucleus
Lord Rutherford (1911) and his coworkers performed an experiment. They bombarded a very thin, gold fail with Alpha particles from a radioactive source. They observed that most of the particles passed straight through the foil undeflected. But a few particles were deflected at different angles. One out of 4000 Alpha particles was deflected at an angle greater than 150.
(Diagram)
Conclusion
Following conclusions were drawn from the Rutherford’s Alpha Particles scattering experiment.
1. The fact that majority of the particles went through the foil undeflected shows that most of the space occupied by an atom is empty.
2. The deflection of a few particles over a wide angle of 150 degrees shows that these particles strike with heavy body having positive charge.
3. The heavy positively charged central part of the atom is called nucleus.
4. Nearly all of the mass of atom is concentrated in the nucleus.
5. The size of the nucleus is very small as compared with the size of atom.
Defects of Rutherford Model
Rutherford model of an atom resembles our solar system. It has following defects:
1. According to classical electromagnetic theory, electron being charged body will emit energy continuously. Thus the orbit of the revolving electron becomes smaller and smaller until it would fall into the nucleus and atomic structure would collapse.
2. If revolving electron emits energy continuously then there should be a continuous spectrum but a line spectrum is obtained.
(Diagram)
Bohr’s Atomic Model
Neil Bohr (1913) presented a model of atom which has removed the defects of Rutherford Model. This model was developed for hydrogen atom which has only proton in the nucleus and one electron is revolving around it.
Postulates of Bohr’s Atomic Model
The main postulates of Bohr’s Model are given below:
1. Electrons revolve around the nucleus in a fixed orbit.
2. As long as electron revolves in a fixed orbit it does not emit and absorb energy. Hence energy of electron remains constant.
3. The orbit nearest to the nucleus is the first orbit and has lowest energy. When an electron absorbs energy it jumps from lower energy orbit to higher energy orbit. Energy is emitted in the form of radiations, when an electron jumps from higher energy orbit to lower energy orbit. The unit of energy emitted in the form of radiations is called quantum. It explains the formation of atomic spectrum.
4. The change in energy is related with the quantum of radiation by the equation :
E2 – E1 = hv
where
E1 = Energy of first orbit
E2 = Energy of the second orbit
h = Planck’s constant
v = Frequency of radiation
Atomic Number
The number of protons present in the nucleus of an atom is called atomic number or proton number. It is denoted by z. The proton in the nucleus of an atom is equal to number of electrons revolving around its nucleus.
Mass Number
The total number of the protons and neutrons present in the nucleus of an atom is called mass number. The protons and neutrons together are called nucleon. Hence it is also known as nucleon number. It is denoted by A. the number of neutrons present in the nucleus of an atom is rperesented by N.
Mass Number = No of Protons + No of neutrons
A = Z + N
Isotopes
The atoms of same elements which have same atomic number but different mas number are called Isotopes. The number of protons present in the nucleus of an atom remains the same but number of neutrons may differ.
Isotopes of Different Elements
Isotopes of Hydrogen
Hydrogen has three isotopes:
1. Ordinary Hydrogen or Protium, H.
2. Heavy Hydrogen or Deutrium, D.
3. Radioactive Hydrogen or Tritium, T.
Protium
Ordinary naturally occurring hydrogen contains the largest percentage of protium. It is denoted by symbol H. It has one proton in its nucleus and one electron revolve around the nucleus.
Number of Protons = 1
Number of Electrons = 1
Number of Neutrons = 0
Atomic Number = 1
Mass Number = 1
Deutrium
Deutrium is called heavy hydrogen. The percentage of deutrium in naturally occuring hydrogen is about 0.0015%. It has one proton and one neutron in its nucleus. It has one electron revolving around its nucleus. It is denoted by symbol D.
Number of Proton = 1
Number of Electron = 1
Number of Neutrons = 1
Atomic Number = 1
Mass Number = 2
Tritium
Radioactive hydrogen is called tritium. It is denoted by symbol T. The number of tritium isotope is one in ten millions. It has one proton and 2 neutrons in its nucleus. It has one electron revolving around its nucleus.
Number of Proton = 1
Number of Electron = 1
Number of Neutron = 2
Atomic Number = 1
Mass Number = 3
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The important postulates of Dalton’s atomic theory are:
1. All elements are composed of atoms. Atom is too small so that it could not be divided into further simpler components.
2. Atom cannot be created nor destroyed in any chemical reactions.
3. Atoms of an element are similar in all respects. They have same mass and properties.
4. Atoms of different elements combine in a definite simple ratio to produce compounds.
Discovery of Electron
A discharge tube is a glass tube. It has two electrode, a source of electric current and a vacuum pump.
(Diagram)
Sir William Crooks (1895 performed experiments by passing electric current through gas in the discharge tube at very low pressure. He observed that at 10-4 (-4 is power to 10) atmosphere pressure, shining rays are emitted from cathode. These rays were named cathode rays. Cathode rays are material particles as they have mass and momentum.
Properties of Cathode Rays
The properties of these particles are given below:
1. These particles are emitted from cathode surface and move in straight line.
2. The temperature of the object rises on which they fall.
3. They produce shadow of opaque object placed in their path.
4. These particles are deflected in electric and magnetic fields.
5. These particles are deflected towards positive plate of electric field.
Discovery of Proton
Gold Stein (1886) observed that in addition to the cathode rays, another type of rays were present in the discharge tube. These rays travel in a direction opposite to cathode rays. These rays were named positive rays. By using perforated cathode in the discharge tube the properties of these rays can be studied. Positive rays are also composed of metered particles. The positive rays are not emitted from anode. They are produced by the ionization of residual gas molecules in the discharge tube. When cathode rays strike with gas molecule, electrons are removed and positive particles are produced.
Properties of Positive Rays
1. They are deflected towards negative plate of electric field. Therefore these rays carry positive charge.
2. The mass of positive rays is equal to the mass of the gas enclosed in the discharge tube.
3. The minimum mass of positive particles is equal to the mass of hydrogen ion (H+). These positive ions are called Protons.
4. The charge on proton is equal to +1.602×10-19 Coulomb. (-19 is power of 10)
Natural Radioactivity
The phenomenon in which certain elements emit radiation which can cause fogging of photographic plate is called natural radioactivity. The elements which omit these rays are called radioactive elements like Uranium, Thorium, Radium etc. There are about 40 radioactive elements. Henri Bequrel (1896) discovered radioactivity.Madam Curei also has valuable contribution in this field.
In natural radioactivity nuclei of elements are broken and element converted to other elements. Natural radioactivity is nuclear property of the elements.
Alpha Rays
1. They are helium nuclei. They are doubly positively charged, He2+.
2. They move with speed equal to the 1/10th of the velocity of the light.
3. They cannot pass through thick-metal foil.
4. They are very good ionizer of a gas.
5. They affect the photographic plate.
Beta Rays
1. They are negatively charged.
2. They move with the speed equal to the velocity of light.
3. They can pass through a few millimeter thick metal sheets.
4. They are good ionizer of a gas.
5. They can affect the photographic plate.
Gamma Rays
1. They are electromagnetic radiations.
2. They travel with speed equal to velocity of light.
3. They carry no charge.
4. They have high penetration power than alpha and beta rays.
5. They are weak ionizer of gas.
Rutherford Experiment and Discovery of Nucleus
Lord Rutherford (1911) and his coworkers performed an experiment. They bombarded a very thin, gold fail with Alpha particles from a radioactive source. They observed that most of the particles passed straight through the foil undeflected. But a few particles were deflected at different angles. One out of 4000 Alpha particles was deflected at an angle greater than 150.
(Diagram)
Conclusion
Following conclusions were drawn from the Rutherford’s Alpha Particles scattering experiment.
1. The fact that majority of the particles went through the foil undeflected shows that most of the space occupied by an atom is empty.
2. The deflection of a few particles over a wide angle of 150 degrees shows that these particles strike with heavy body having positive charge.
3. The heavy positively charged central part of the atom is called nucleus.
4. Nearly all of the mass of atom is concentrated in the nucleus.
5. The size of the nucleus is very small as compared with the size of atom.
Defects of Rutherford Model
Rutherford model of an atom resembles our solar system. It has following defects:
1. According to classical electromagnetic theory, electron being charged body will emit energy continuously. Thus the orbit of the revolving electron becomes smaller and smaller until it would fall into the nucleus and atomic structure would collapse.
2. If revolving electron emits energy continuously then there should be a continuous spectrum but a line spectrum is obtained.
(Diagram)
Bohr’s Atomic Model
Neil Bohr (1913) presented a model of atom which has removed the defects of Rutherford Model. This model was developed for hydrogen atom which has only proton in the nucleus and one electron is revolving around it.
Postulates of Bohr’s Atomic Model
The main postulates of Bohr’s Model are given below:
1. Electrons revolve around the nucleus in a fixed orbit.
2. As long as electron revolves in a fixed orbit it does not emit and absorb energy. Hence energy of electron remains constant.
3. The orbit nearest to the nucleus is the first orbit and has lowest energy. When an electron absorbs energy it jumps from lower energy orbit to higher energy orbit. Energy is emitted in the form of radiations, when an electron jumps from higher energy orbit to lower energy orbit. The unit of energy emitted in the form of radiations is called quantum. It explains the formation of atomic spectrum.
4. The change in energy is related with the quantum of radiation by the equation :
E2 – E1 = hv
where
E1 = Energy of first orbit
E2 = Energy of the second orbit
h = Planck’s constant
v = Frequency of radiation
Atomic Number
The number of protons present in the nucleus of an atom is called atomic number or proton number. It is denoted by z. The proton in the nucleus of an atom is equal to number of electrons revolving around its nucleus.
Mass Number
The total number of the protons and neutrons present in the nucleus of an atom is called mass number. The protons and neutrons together are called nucleon. Hence it is also known as nucleon number. It is denoted by A. the number of neutrons present in the nucleus of an atom is rperesented by N.
Mass Number = No of Protons + No of neutrons
A = Z + N
Isotopes
The atoms of same elements which have same atomic number but different mas number are called Isotopes. The number of protons present in the nucleus of an atom remains the same but number of neutrons may differ.
Isotopes of Different Elements
Isotopes of Hydrogen
Hydrogen has three isotopes:
1. Ordinary Hydrogen or Protium, H.
2. Heavy Hydrogen or Deutrium, D.
3. Radioactive Hydrogen or Tritium, T.
Protium
Ordinary naturally occurring hydrogen contains the largest percentage of protium. It is denoted by symbol H. It has one proton in its nucleus and one electron revolve around the nucleus.
Number of Protons = 1
Number of Electrons = 1
Number of Neutrons = 0
Atomic Number = 1
Mass Number = 1
Deutrium
Deutrium is called heavy hydrogen. The percentage of deutrium in naturally occuring hydrogen is about 0.0015%. It has one proton and one neutron in its nucleus. It has one electron revolving around its nucleus. It is denoted by symbol D.
Number of Proton = 1
Number of Electron = 1
Number of Neutrons = 1
Atomic Number = 1
Mass Number = 2
Tritium
Radioactive hydrogen is called tritium. It is denoted by symbol T. The number of tritium isotope is one in ten millions. It has one proton and 2 neutrons in its nucleus. It has one electron revolving around its nucleus.
Number of Proton = 1
Number of Electron = 1
Number of Neutron = 2
Atomic Number = 1
Mass Number = 3
Class IX, CHEMISTRY, "Differences"
Metals and Non Metals
Metals
1. Metals have luster shine surface.
2. Metals reflect heat and light.
3. Metals conduct heat and electricity
4. Metals are ductile and can be drawn into wire.
Non-Metals
1. Non-Metals have no luster.
2. Non-Metals usually don’t reflect heat and light.
3. Non-Metals do not conduct heat and electricity.
4. Non-Metals are non ductile and cannot be drawn into wire.
5. Non-Metals are non-malleable and can not form sheets.
Homogeneous and Heterogeneous Mixture
Homogeneous Mixture
1. Those mixtures, which have uniform composition throughout their mass are called homogeneous mixtures.
2. Homogeneous mixture has only one phase through out its mass.
3. Homogeneous mixture are also known as solution.
4. Examples: Salt and water, Sugar and water.
Heterogeneous Mixture
1. Those mixtures, which do not have uniform composition through their mass are called Heterogeneous Mixture.
2. Heterogeneous Mixture has more than one phase through out its mass.
3. Heterogeneous Mixture are not solutions.
4. Examples: Rocks, Soil, Food products.
Molecular and Empirical Formula
Molecular Formula
1. Formula which shows the actual number of atoms of each element present in a molecule is called Molecular Formula.
2. Molecular Formula shows the structure of compound.
3. Two or more compounds cannot have same Molecular Formula.
4. Molecular Formula = n x Empirical Formula.
5. It represents covalent compounds only.
Empirical Formula
1. formula, which shows the relative ratio of atoms of each element present in a molecule, is called Empirical Formula.
2. Empirical Formula can not show the structure of compound.
3. Two or more compounds can have same Empirical Formula.
4. Empirical Formula = Molecular Formula / n
5. It represent an ionic compound as well as a covalent compound.
Symbol and Formula
Symbol
1. A symbol is an abbreviation for the chemical name of an element and represents only one atom of the element.
2. It represents one atom of an element.
3. Symbol is written for elements.
4. Examples: Na, Br, Cl, F etc.
Formula
1. Representation of compound in terms of symbols is called formula. It represents one atom of an element.
2. It represents atoms of same or different elements present in one molecule.
3. It represents an ionic compounds as well as a covalent compound.
4. Examples: H2O, NH3 etc.
Gram and Gram Molecule
Gram
The atomic mass of an element expressed in grams is called gram atomic mass.
2. It is associated with element only.
3. It is the mass of one atomic mole.
4. One gram atom of any substance contains 6.02 x 10(23) atoms. (23 is the power of 10).
Gram Molecule
1. Molecular mass of any element or compound expressed in grams is called gram molecule.
2. It is associated with element and compound.
3. It is the mass of one molecular mole.
4. One gram molecule of any substance contains 6.02 x 10(23) atoms. (23 is the power of 10).
Atom and Molecule
Atom
1. It is the smallest particle of an element which can enter into a chemical reaction.
2. It is represented by a symbol of the element.
3. It shows the properties of the element.
4. It retains its identity in a chemical reaction.
Molecule
1. It is the smallest particle of a substance which can exist and show all the properties of the substance.
2. It is represented by a molecular formula of the substance.
3. It shows the properties of the substance.
4. It does not retain its identity in a chemical reaction.
Exothermic and Endothermic Reactions
Exothermic Reaction
1. Those chemical reactions in which heat energy is evolved are called exothermic reactions.
2. In exothermic reactions the enthalpy of products is lower than the reactants. H is therefore negative for an exothermic reaction.
3. During endothermic reaction, the system becomes colder and net potential energy of substance increases.
4. The energy is absorbed during these reactions.
5. The temperature of reaction therefore decreases.
Endothermic Reactions
1. Those chemical reactions in which heat energy is absorbed are called endothermic reactions.
2. In endothermic reactions the enthalpy of reactants is lower than the products. H is therefore positive in endothermic reaction.
3. During endothermic reaction, the system becomes colder and net potential energy of substance increases.
4. The energy is absorbed during these reactions.
5. The temperature of reaction therefore decreases.
Physical and Chemical Properties
Physical Properties
1. The physical properties of a substance are those characteristics which serve to distinguish it from other substance but do not deal with its ability to undergo chemical changes.
2. These are related to the physical state of matter.
3. Examples: Formation of ice from water, formation of a magnet from ice etc.
Chemical Properties
1. The chemical properties of a substance indicate the ability of a substance to undergo chemical changes.
2. They are related to the chemical change of a substance.
3. Examples: burning of paper, rusting of iron.
Electrolyte and Non-Electrolyte
Electrolytes
1. Electrolytes conduct electricity in molten or in solution form.
2. These form positive and negative ions when dissolved in water e.g. NaCl form Na+ and Cl- ions when dissolved in water.
3. Chemical changes occur when electric current is passed through the electrolyte.
4. Generally these are ionic or polar covalent compounds.
Non-Electrolytes
1. Non-electrolytes do not conduct electric current in molten or in solution form.
2. These do not form positive and negative ions when dissolved in water e.g. Urea, sugar, glucose etc.
2. No chemical change occurs in them on passing current.
3. Generally these are non polar covalent compounds.
4. Generally these are non polar covalent compounds.
Acid and Base
Acid
1. Those compounds which provide hydrogen ion (H+) in aqueous solutions are called Acids.
2. An acid is a substance which produces H+ ions in aqueous solution.
3. Acid is a species (a compound or ion) which donates or tends to donate a proton (H+).
4. An acid is a species (molecule or ion) which can accept a pair of electron. An acid is also called an electrophile (electron loving).
5. They have sour taste.
6. Acid turn blue litmus red methyl orange red.
Base
1. Those compounds, which provides hydroxyl (OH-) ion in aqueous solution, are called bases.
2. A base is a substance, which gives (OH-) in aqueous solution.
3. A base is a species, which accepts or tends to accept a proton.
4. A base is a species (molecule or ion) which can donate a pair of electrons. A base is also called a nucleophile (Nucleus loving).
5. Bases have bitter taste.
6. Bases turn red litmus to blue, colorless phenolphthalein to pink and methyl orange to yellow.
Ionic and Covalent Bond
Ionic Bond
1. Ionic bond is formed by complete transfer of electrons from one atom to another atom.
2. Ionic bond is always formed between different atoms. E.g. NaCl, CaCl2.
3. In ionic bond atoms have very large electro-negativity and ionization energy difference.
4. This bond is usually formed between metals and non-metals.
5. This bond is very strong.
6. As a result of this bond ionic compounds are formed.
7. It is always formed between two different atoms.
8. It is formed when difference of electro-negativity of combining atoms is 1.7 or more.
Covalent Bond
1. Covalent bond is formed by the mutual sharing of electrons between two atoms.
2. Covalent bond may be formed between similar or dissimilar atoms e.g. H2, O2, HCl etc.
3. In covalent bond atoms have very small electro-negativity or ionization energy difference.
4. This bond is usually formed between non-metals only.
5. This bond is comparatively less strong.
6. As a result of this bond covalent compounds are formed.
7. It is formed between similar and different types of atoms.
8. It is formed when difference of electro-negativity of combining atoms is less than 1.7.
Ionic and Covalent Compounds
Ionic Compounds
1. The ionic compounds are usually solid, hard and brittle.
2. The ionic compounds are good conductors of electricity either in fused state or in the form of aqueous solution.
3. Ionic Compounds have high melting points and boiling points.
4. Ionic compounds have high melting points and boiling points.
5. Covalent compounds are mostly volatile.
Covalent Compounds
1. Covalent compounds exist in all the three states i.e. gas, liquid and solid.
2. A pure covalent compound does not conduct electricity.
3. These have usually low melting and boiling points.
4. These are soluble in water.
5. These are insoluble in water but soluble in organic solvents.
Co-Ordinate Covalent and Covalent Bond
Co-Ordinate Covalent Bond
1. It is a bond in which the shared electron pair is denoted by one atom only.
2. One atom donates electrons but other has no contribution.
3. Lewis acids and bases always from this bond.
4. It is represented by ->.
5. It is formed by the donation of an electron apir by one of the two bonded atoms.
6. It is formed by the completely filled atomic orbital.
Covalent Bond
1. It is a bond formed by the mutual sharing of electrons.
2. In the shared electron pair both atoms have equal contribution.
3. Lewis acids and bases do not form this bond.
4. It is represented by _.
5. It is formed by the mutual sharing of electrons between atoms.
6. It is formed by the overlap of partially filled atomic orbital.
Polar and Non-Polar Covalent Bond
Polar Covalent Bond
1. The covalent bond between two atoms having different electro-negativity is called a polar covalent bond.
2. In a polar bond, the shared electron pair is not equally attracted by the bonded atoms.
3. Bonded atoms become slightly charged and acquire partial =ve and -ve charges.
4. It has an ionic character.
5. The bond energy is greater.
Non-Polar Covalent Bond
1. The covalent bond between two atoms having same electro-negativity is called a non-polar covalent bond.
2. In a non polar bond, the shared electron pair is equally attracted by the bonded atoms.
3. Bonded atoms remain electrically neutral and do not acquire partial charges.
4. It has no ionic character.
5. The bond energy is lesser.
Electrolytic and Galvanic or Voltaic Cell
Electrolytic Cell
1. It is a device for converting electrical energy into chemical energy. It means by passing current through an electrolyte, chemical reaction takes place.
2. It consists of a vessel containing an electrodes and a source of direct current (battery).
3. Example: Electrolysis of aqueous solution of NaCl.
Galvanic or Voltaic Cell
1. It is a device for converting chemical energy into electrical energy. It means spontaneous redox reaction is used for the production of electric current. This cell was prepared by L.Galvani and A.Volts, hence named as Galvanic or Voltaic Cell.
2. It consists of two half-cells. Each half cell consists of an electrodes and the solution with which it is in contact.
3. Example: Daniel Cell-Zn/ZnSO4 and Cu/CuSO4 cell.
Solution and Suspension
Solution
The size of particles is between 0.1 to 1nm.
2. Particles cannot be seen with low power microscope.
3. It is homogeneous.
4. Particles do not settle down.
5. It is transparent.
6. Components cannot be separated by filtration.
Suspension
1. The size of particles is larger than 1000nm.
2. Particles can be seen by low power microscope.
3. It is heterogeneous.
4. Particles settle down.
5. It is not transparent.
6. Components can be separated by filtration.
Metals
1. Metals have luster shine surface.
2. Metals reflect heat and light.
3. Metals conduct heat and electricity
4. Metals are ductile and can be drawn into wire.
Non-Metals
1. Non-Metals have no luster.
2. Non-Metals usually don’t reflect heat and light.
3. Non-Metals do not conduct heat and electricity.
4. Non-Metals are non ductile and cannot be drawn into wire.
5. Non-Metals are non-malleable and can not form sheets.
Homogeneous and Heterogeneous Mixture
Homogeneous Mixture
1. Those mixtures, which have uniform composition throughout their mass are called homogeneous mixtures.
2. Homogeneous mixture has only one phase through out its mass.
3. Homogeneous mixture are also known as solution.
4. Examples: Salt and water, Sugar and water.
Heterogeneous Mixture
1. Those mixtures, which do not have uniform composition through their mass are called Heterogeneous Mixture.
2. Heterogeneous Mixture has more than one phase through out its mass.
3. Heterogeneous Mixture are not solutions.
4. Examples: Rocks, Soil, Food products.
Molecular and Empirical Formula
Molecular Formula
1. Formula which shows the actual number of atoms of each element present in a molecule is called Molecular Formula.
2. Molecular Formula shows the structure of compound.
3. Two or more compounds cannot have same Molecular Formula.
4. Molecular Formula = n x Empirical Formula.
5. It represents covalent compounds only.
Empirical Formula
1. formula, which shows the relative ratio of atoms of each element present in a molecule, is called Empirical Formula.
2. Empirical Formula can not show the structure of compound.
3. Two or more compounds can have same Empirical Formula.
4. Empirical Formula = Molecular Formula / n
5. It represent an ionic compound as well as a covalent compound.
Symbol and Formula
Symbol
1. A symbol is an abbreviation for the chemical name of an element and represents only one atom of the element.
2. It represents one atom of an element.
3. Symbol is written for elements.
4. Examples: Na, Br, Cl, F etc.
Formula
1. Representation of compound in terms of symbols is called formula. It represents one atom of an element.
2. It represents atoms of same or different elements present in one molecule.
3. It represents an ionic compounds as well as a covalent compound.
4. Examples: H2O, NH3 etc.
Gram and Gram Molecule
Gram
The atomic mass of an element expressed in grams is called gram atomic mass.
2. It is associated with element only.
3. It is the mass of one atomic mole.
4. One gram atom of any substance contains 6.02 x 10(23) atoms. (23 is the power of 10).
Gram Molecule
1. Molecular mass of any element or compound expressed in grams is called gram molecule.
2. It is associated with element and compound.
3. It is the mass of one molecular mole.
4. One gram molecule of any substance contains 6.02 x 10(23) atoms. (23 is the power of 10).
Atom and Molecule
Atom
1. It is the smallest particle of an element which can enter into a chemical reaction.
2. It is represented by a symbol of the element.
3. It shows the properties of the element.
4. It retains its identity in a chemical reaction.
Molecule
1. It is the smallest particle of a substance which can exist and show all the properties of the substance.
2. It is represented by a molecular formula of the substance.
3. It shows the properties of the substance.
4. It does not retain its identity in a chemical reaction.
Exothermic and Endothermic Reactions
Exothermic Reaction
1. Those chemical reactions in which heat energy is evolved are called exothermic reactions.
2. In exothermic reactions the enthalpy of products is lower than the reactants. H is therefore negative for an exothermic reaction.
3. During endothermic reaction, the system becomes colder and net potential energy of substance increases.
4. The energy is absorbed during these reactions.
5. The temperature of reaction therefore decreases.
Endothermic Reactions
1. Those chemical reactions in which heat energy is absorbed are called endothermic reactions.
2. In endothermic reactions the enthalpy of reactants is lower than the products. H is therefore positive in endothermic reaction.
3. During endothermic reaction, the system becomes colder and net potential energy of substance increases.
4. The energy is absorbed during these reactions.
5. The temperature of reaction therefore decreases.
Physical and Chemical Properties
Physical Properties
1. The physical properties of a substance are those characteristics which serve to distinguish it from other substance but do not deal with its ability to undergo chemical changes.
2. These are related to the physical state of matter.
3. Examples: Formation of ice from water, formation of a magnet from ice etc.
Chemical Properties
1. The chemical properties of a substance indicate the ability of a substance to undergo chemical changes.
2. They are related to the chemical change of a substance.
3. Examples: burning of paper, rusting of iron.
Electrolyte and Non-Electrolyte
Electrolytes
1. Electrolytes conduct electricity in molten or in solution form.
2. These form positive and negative ions when dissolved in water e.g. NaCl form Na+ and Cl- ions when dissolved in water.
3. Chemical changes occur when electric current is passed through the electrolyte.
4. Generally these are ionic or polar covalent compounds.
Non-Electrolytes
1. Non-electrolytes do not conduct electric current in molten or in solution form.
2. These do not form positive and negative ions when dissolved in water e.g. Urea, sugar, glucose etc.
2. No chemical change occurs in them on passing current.
3. Generally these are non polar covalent compounds.
4. Generally these are non polar covalent compounds.
Acid and Base
Acid
1. Those compounds which provide hydrogen ion (H+) in aqueous solutions are called Acids.
2. An acid is a substance which produces H+ ions in aqueous solution.
3. Acid is a species (a compound or ion) which donates or tends to donate a proton (H+).
4. An acid is a species (molecule or ion) which can accept a pair of electron. An acid is also called an electrophile (electron loving).
5. They have sour taste.
6. Acid turn blue litmus red methyl orange red.
Base
1. Those compounds, which provides hydroxyl (OH-) ion in aqueous solution, are called bases.
2. A base is a substance, which gives (OH-) in aqueous solution.
3. A base is a species, which accepts or tends to accept a proton.
4. A base is a species (molecule or ion) which can donate a pair of electrons. A base is also called a nucleophile (Nucleus loving).
5. Bases have bitter taste.
6. Bases turn red litmus to blue, colorless phenolphthalein to pink and methyl orange to yellow.
Ionic and Covalent Bond
Ionic Bond
1. Ionic bond is formed by complete transfer of electrons from one atom to another atom.
2. Ionic bond is always formed between different atoms. E.g. NaCl, CaCl2.
3. In ionic bond atoms have very large electro-negativity and ionization energy difference.
4. This bond is usually formed between metals and non-metals.
5. This bond is very strong.
6. As a result of this bond ionic compounds are formed.
7. It is always formed between two different atoms.
8. It is formed when difference of electro-negativity of combining atoms is 1.7 or more.
Covalent Bond
1. Covalent bond is formed by the mutual sharing of electrons between two atoms.
2. Covalent bond may be formed between similar or dissimilar atoms e.g. H2, O2, HCl etc.
3. In covalent bond atoms have very small electro-negativity or ionization energy difference.
4. This bond is usually formed between non-metals only.
5. This bond is comparatively less strong.
6. As a result of this bond covalent compounds are formed.
7. It is formed between similar and different types of atoms.
8. It is formed when difference of electro-negativity of combining atoms is less than 1.7.
Ionic and Covalent Compounds
Ionic Compounds
1. The ionic compounds are usually solid, hard and brittle.
2. The ionic compounds are good conductors of electricity either in fused state or in the form of aqueous solution.
3. Ionic Compounds have high melting points and boiling points.
4. Ionic compounds have high melting points and boiling points.
5. Covalent compounds are mostly volatile.
Covalent Compounds
1. Covalent compounds exist in all the three states i.e. gas, liquid and solid.
2. A pure covalent compound does not conduct electricity.
3. These have usually low melting and boiling points.
4. These are soluble in water.
5. These are insoluble in water but soluble in organic solvents.
Co-Ordinate Covalent and Covalent Bond
Co-Ordinate Covalent Bond
1. It is a bond in which the shared electron pair is denoted by one atom only.
2. One atom donates electrons but other has no contribution.
3. Lewis acids and bases always from this bond.
4. It is represented by ->.
5. It is formed by the donation of an electron apir by one of the two bonded atoms.
6. It is formed by the completely filled atomic orbital.
Covalent Bond
1. It is a bond formed by the mutual sharing of electrons.
2. In the shared electron pair both atoms have equal contribution.
3. Lewis acids and bases do not form this bond.
4. It is represented by _.
5. It is formed by the mutual sharing of electrons between atoms.
6. It is formed by the overlap of partially filled atomic orbital.
Polar and Non-Polar Covalent Bond
Polar Covalent Bond
1. The covalent bond between two atoms having different electro-negativity is called a polar covalent bond.
2. In a polar bond, the shared electron pair is not equally attracted by the bonded atoms.
3. Bonded atoms become slightly charged and acquire partial =ve and -ve charges.
4. It has an ionic character.
5. The bond energy is greater.
Non-Polar Covalent Bond
1. The covalent bond between two atoms having same electro-negativity is called a non-polar covalent bond.
2. In a non polar bond, the shared electron pair is equally attracted by the bonded atoms.
3. Bonded atoms remain electrically neutral and do not acquire partial charges.
4. It has no ionic character.
5. The bond energy is lesser.
Electrolytic and Galvanic or Voltaic Cell
Electrolytic Cell
1. It is a device for converting electrical energy into chemical energy. It means by passing current through an electrolyte, chemical reaction takes place.
2. It consists of a vessel containing an electrodes and a source of direct current (battery).
3. Example: Electrolysis of aqueous solution of NaCl.
Galvanic or Voltaic Cell
1. It is a device for converting chemical energy into electrical energy. It means spontaneous redox reaction is used for the production of electric current. This cell was prepared by L.Galvani and A.Volts, hence named as Galvanic or Voltaic Cell.
2. It consists of two half-cells. Each half cell consists of an electrodes and the solution with which it is in contact.
3. Example: Daniel Cell-Zn/ZnSO4 and Cu/CuSO4 cell.
Solution and Suspension
Solution
The size of particles is between 0.1 to 1nm.
2. Particles cannot be seen with low power microscope.
3. It is homogeneous.
4. Particles do not settle down.
5. It is transparent.
6. Components cannot be separated by filtration.
Suspension
1. The size of particles is larger than 1000nm.
2. Particles can be seen by low power microscope.
3. It is heterogeneous.
4. Particles settle down.
5. It is not transparent.
6. Components can be separated by filtration.
Class IX, CHEMISTRY, "Solution and Suspension"
Solution
A homogeneous mixture of different chemical substances which has uniform chemical composition through out and shows uniform physical properties is called solution. For example dissolve a small amount of copper sulphate in water the water will become blue. If this blue liquid is filtered, it will pass through the filter paper without leaving any solid. The mixture thus prepared is called a solution.
Binary Solution
A solution which is formed by mixing two substances is called binary solution. For example solution of glucose and water.
Solute
The component of a binary solution which is in lesser amount is called solute. For example in copper sulphate solution, copper sulphate is solute.
Solvent
The component of a binary solution which is in greater amount is called solvent. For example in copper sulphate solution, water is solvent.
Saturated solution
A solution in which maximum amount of a solute has been dissolved at a particular temperature and in which the dissolved form of solute is at equilibrium with its undissolved form is called saturated solution.
Unsaturated Solution
Solution which can dissolve further amount of a solute at a [particular temperature is called an unsaturated solution.
Supersaturated Solution
The solution which contains even more amount of solute required to prepare saturated solution is called super saturated solution. The hot saturated solution of compound like sodium thiosulphate does not crystallize its solute if cooled slowly without disturbance. Such a solution is called supersaturated solution.
Dilute Solution
A solution which contains small amount of a solute as compared to the solvent is called dilute solution.
Concentrated Solution
A solution which contains excess amount of a solute as compared to that of a solvent is called a concentrated solution.
Concentrated Solution
The amount of solute present in given quantity of solvent is called concentration of solution. The concentration of a solution can be expressed in many ways depending upon the amount o solute and solvent present in it.
Concentration of Solution
The amount of solute present in given quantity of solvent is called concentration of solution. The concentration of a solution can be expressed in many ways depending upon the amount of solute and solvent present in it.
Percentage by Mass
The percentage of solute by mass is the mass of solute present in hundred part of the solution. For example 5% hydrogen peroxide solution by mass means that 5g hydrogen peroxide are dissolved in 95g of water to give 100g of solution.
Percentage of Mass = (Mass of Solute/Mass of Solution) x 100
Percentage by Volume
The concentration unit expresses the volume of solute present in 100cm3 of solution. For example 15% solution of alcohol by volume will mean that 15cm3 alcohols are present in 100cm3 of solution. (Here 3 represents cube)
Percentage by Volume = (Volume of Solute/Volume of Solution) x 100
Molar Solution
The solution that contains one mole of solute in 1dm3 of solution is called a molar solution. The concentration of this solution is expressed as M.
Molarity
Molarity of a solution is the number of moles of solute present in 1dm3 of the solution. It is expressed as M.
M = Number of Moles of Solute/Volume of Solution in dm3
or
M = (Mass of solute/Molecular Mass) x (1/ Volume of Solution in dm3)
Crystallization
The process in which crystal separates from saturated solution on cooling is called crystallization. It is a useful process because it can be used to purify the impure solid compounds. It can also be used to separate a mixture of solids.
Hydration
The ions surrounded by solvent molecules in solution are called solvated ions. If water is a solvent these ions are called hydrated ions.
Suspension
A suspension in such a mixture in which solute particles do not dissolved in solvent and if filtrated its particles do not pass through the pores of filter paper.
Colloidal Solution
In a colloidal solution the solute particles are slightly bigger than those present in a true solution but not big enough to seen with naked eye.
Standard Solution
A solution whose molarity (strength) is known is called Standard Solution.
True Solution
A True Solution is such a mixture in which solute particles are completely homogenized in the solvent for example solution of sodium chloride or copper sulphate in water.
Solubility
Solubility o a solute in a particular solvent is defined as the amount of solute in grams, which can dissolve in 100g of the solvent at a particular temperature to give a saturated solution.
or
The amount of a solute in gram moles, which can dissolve in one kilogram of the solvent at a particular temperature, to give a saturated solution.
Factors Affecting the Solubility
Effect of Solvent
Similar solvents dissolve similar solutes, i.e. if the chemical structure and the electrical properties such as dipole moment of solute and solvent are similar, the solubility will increase. If there is dissimilarity in properties, then either the solute will not dissolve or there will be very little solubility.
Effect of Solute
Different solutes have different solubility’s in a particular solvent e.g. if the saturated solutions of table sugar and sodium chloride are prepared, it is found that the concentration of sodium chloride solution is 5.3 molar while that of sugar solution is 3.8 molar. In other words, the solubility of sodium chloride in water is far greater than that of sugar. This is due to the fact that the attraction of sodium (Na+ and chloride (Cl-) ions with water is greater than that of sugar molecules with water.
Effect of Temperature
Change in temperature has different effects on the solubility of different compounds. Usually the solubility increase with the increase in temperature but it cannot be taken as a general rule. The solubility of compounds like lithium carbonate, calcium chromate decreases with the increase in temperature. The solubility of gases in water also decreases with the increase in temperature. On the other hand, there are a large number of compounds whose solubility in water increase with the increase in temperature e.g. sodium nitrate, silver nitrate, Potassium chloride etc. the solubility of sodium chloride in water does not increase appreciably with the increase in temperature.
A homogeneous mixture of different chemical substances which has uniform chemical composition through out and shows uniform physical properties is called solution. For example dissolve a small amount of copper sulphate in water the water will become blue. If this blue liquid is filtered, it will pass through the filter paper without leaving any solid. The mixture thus prepared is called a solution.
Binary Solution
A solution which is formed by mixing two substances is called binary solution. For example solution of glucose and water.
Solute
The component of a binary solution which is in lesser amount is called solute. For example in copper sulphate solution, copper sulphate is solute.
Solvent
The component of a binary solution which is in greater amount is called solvent. For example in copper sulphate solution, water is solvent.
Saturated solution
A solution in which maximum amount of a solute has been dissolved at a particular temperature and in which the dissolved form of solute is at equilibrium with its undissolved form is called saturated solution.
Unsaturated Solution
Solution which can dissolve further amount of a solute at a [particular temperature is called an unsaturated solution.
Supersaturated Solution
The solution which contains even more amount of solute required to prepare saturated solution is called super saturated solution. The hot saturated solution of compound like sodium thiosulphate does not crystallize its solute if cooled slowly without disturbance. Such a solution is called supersaturated solution.
Dilute Solution
A solution which contains small amount of a solute as compared to the solvent is called dilute solution.
Concentrated Solution
A solution which contains excess amount of a solute as compared to that of a solvent is called a concentrated solution.
Concentrated Solution
The amount of solute present in given quantity of solvent is called concentration of solution. The concentration of a solution can be expressed in many ways depending upon the amount o solute and solvent present in it.
Concentration of Solution
The amount of solute present in given quantity of solvent is called concentration of solution. The concentration of a solution can be expressed in many ways depending upon the amount of solute and solvent present in it.
Percentage by Mass
The percentage of solute by mass is the mass of solute present in hundred part of the solution. For example 5% hydrogen peroxide solution by mass means that 5g hydrogen peroxide are dissolved in 95g of water to give 100g of solution.
Percentage of Mass = (Mass of Solute/Mass of Solution) x 100
Percentage by Volume
The concentration unit expresses the volume of solute present in 100cm3 of solution. For example 15% solution of alcohol by volume will mean that 15cm3 alcohols are present in 100cm3 of solution. (Here 3 represents cube)
Percentage by Volume = (Volume of Solute/Volume of Solution) x 100
Molar Solution
The solution that contains one mole of solute in 1dm3 of solution is called a molar solution. The concentration of this solution is expressed as M.
Molarity
Molarity of a solution is the number of moles of solute present in 1dm3 of the solution. It is expressed as M.
M = Number of Moles of Solute/Volume of Solution in dm3
or
M = (Mass of solute/Molecular Mass) x (1/ Volume of Solution in dm3)
Crystallization
The process in which crystal separates from saturated solution on cooling is called crystallization. It is a useful process because it can be used to purify the impure solid compounds. It can also be used to separate a mixture of solids.
Hydration
The ions surrounded by solvent molecules in solution are called solvated ions. If water is a solvent these ions are called hydrated ions.
Suspension
A suspension in such a mixture in which solute particles do not dissolved in solvent and if filtrated its particles do not pass through the pores of filter paper.
Colloidal Solution
In a colloidal solution the solute particles are slightly bigger than those present in a true solution but not big enough to seen with naked eye.
Standard Solution
A solution whose molarity (strength) is known is called Standard Solution.
True Solution
A True Solution is such a mixture in which solute particles are completely homogenized in the solvent for example solution of sodium chloride or copper sulphate in water.
Solubility
Solubility o a solute in a particular solvent is defined as the amount of solute in grams, which can dissolve in 100g of the solvent at a particular temperature to give a saturated solution.
or
The amount of a solute in gram moles, which can dissolve in one kilogram of the solvent at a particular temperature, to give a saturated solution.
Factors Affecting the Solubility
Effect of Solvent
Similar solvents dissolve similar solutes, i.e. if the chemical structure and the electrical properties such as dipole moment of solute and solvent are similar, the solubility will increase. If there is dissimilarity in properties, then either the solute will not dissolve or there will be very little solubility.
Effect of Solute
Different solutes have different solubility’s in a particular solvent e.g. if the saturated solutions of table sugar and sodium chloride are prepared, it is found that the concentration of sodium chloride solution is 5.3 molar while that of sugar solution is 3.8 molar. In other words, the solubility of sodium chloride in water is far greater than that of sugar. This is due to the fact that the attraction of sodium (Na+ and chloride (Cl-) ions with water is greater than that of sugar molecules with water.
Effect of Temperature
Change in temperature has different effects on the solubility of different compounds. Usually the solubility increase with the increase in temperature but it cannot be taken as a general rule. The solubility of compounds like lithium carbonate, calcium chromate decreases with the increase in temperature. The solubility of gases in water also decreases with the increase in temperature. On the other hand, there are a large number of compounds whose solubility in water increase with the increase in temperature e.g. sodium nitrate, silver nitrate, Potassium chloride etc. the solubility of sodium chloride in water does not increase appreciably with the increase in temperature.
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