Chapter Some Basic Concepts of Chemistry

SOME BASIC CONCEPT OF CHEMISTRY

Classification of matter on the basis of chemical properties:

Matter classified into two groups pure substances and mixtures.Pure substances further divided into elements and compounds.

1. Elements : it is a simple individual which has a defined atomic number and has a definite position in periodic table.

All the elements are classified into three groups metal,non-metal,and semi metals(metalloids)

METALS

Metals are regarded as those elements which have following properties:

(i)They are generally solids at ordinary conditions .i.e. temperature ,pressure etc .Except mercury which is in liquid state at room temperature.
(ii)They are lustrous in nature.
(iii)They posses high density.
(iv)They are good conductors of heat and electricity.
(v)They are malleable and ductile.
(vi)They posses high melting and boiling point.
(vii)They react with mineral acids liberating hydrogen.
(viii)They form basic oxides.
(ix)They form non-volatile hydrides if combine with hydrogen.
(x)They have molecules usually mono-atomic in the vapour state.

 

NON-METALS

Non-metals do not show properties of metal.

Six of non-metals carbon,boron,phosphorus,sulphur, selenium and iodine are solid.
Bromine is the only non-metal at room temperature it is in liquid form.
The remaining non-metals are in gaseous form i.e. nitrogen, oxygen,fluorine, chlorine,hydrogen, helium,argon,neon,krypton,xenon and radon are gases.
Non-metals are generally
(i)Brittle.
(ii)Non-lustrous.
(iii)Having low melting and boiling points.
(iv)Bad conductors of heat and electricity.
(v)Capable of forming acid oxides and nuetral oxides.
(vi)Do not evolve hydrogen from acids.
(vii)Capable of forming volatile hydrides.

 

Semi-metals(metalloids)

Semi-metals(metalloids) : There are many elements which do not fit completely into either metal or non-metal class.
Elements which have some properties of metal and some properties of non-metal are called semi-metals or metalloids.
i.e. silicon,germanium,arsenic,antimony and tellurium.

 

2.Compounds: are also pure substances that are composed of two or more different elements in fixed proportion by mass.The properties of compounds are altogether are different from its constituent elements.
i.e. Water has definite composition of oxygen(88.8%) and hydrogen(11.2%) in ratio (1:8) respectively by mass.Properties of water is completely different from hydrogen and oxygen.Hydrogen and oxygen is in gaseous state where as water is in liquid state.Oxygen and hydrogen both are combustible but water is used in extinguishing fire.
Component of compound elements are separated by chemical means only not by physical means.

Components are further classified into two parts :

(i)Organic compounds: The compounds which are obtained from living sources are termed as organic compounds.Term organic refers to hydrocarbons(compounds of hydrogens and carbons) and their derivatives.

(ii)Inorganic compounds: The compounds which are obtained from non living sources are termed as inorganic compounds.

 

MIXTURES

Mixtures are the combination of two or more substances either elements or compounds or both in any proportion.Substance which form mixture are called components.Components are present in the mixture without loss of their identity.

Mixtures are of two types :

HOMOGENEOUS MIXTURE

1.In Homogeneous mixture components are mixed uniformly to microscopic level and in this components can not be seen with naked eye or microscope.

2.Mixture is uniform throughout having a single phase.

3.Homogeneous mixture is isotropic in nature i.e. having same composition and properties in every proportion.

EXAMPLE : Sugar dissolved in water, methyl alcohol in water,  iodine in carbon tetrachloride, benzene in toluene etc

HETEROGENEOUS MIXTURE

1. Heterogeneous mixture components are not uniformly distributed.Due to this components can be seen with the help of microscope or naked eye.

2.Mixture can have two or more phase.

3.Heterogeneous mixture are anisotropic in nature .i.e properties are not uniform throughout the mixture.

EXAMPLE : Mixture of sulphur and sand, mixture of iron fillings and sand ,muddy water etc.

 

LAW OF CHEMICAL COMBINATION 

(i) Law of conservation of mass

This law was first stated by Lavoiser in 1774.The law implies that mass can neither be created nor destroyed Or We can say that during the whole chemical reaction the total mass of the system remains constant Or Total masses of reactants is equal to total mass of product formed.This law is also known as law of indestructibility of matter.

Experimental verification of law :

The law of conservation of mass is firstly verified by Landolt. Landolt took H shaped tube (now a days called Landolt’s tube) and filled with solutions of sodium chloride(NaCl) and silver nitride(​\( AgNO_3 \)​) separately in two limbs of the tube. This limbed them sealed and H shaped tube weighted accurately. The two solutions then mixed throughly by shaking the tube.Reaction takes place like this :

\[ NaCl(aq)+AgNO_3(aq)\rightarrow{AgCl(s)+NaNO_3(aq)} \]

The tube was again weighted.It as was observed that the mass of tube remained same, before and after the reaction. This verify law of conservation of mass.

LIMITATIONS : In nuclear reaction law of conservation of mass is not obeyed because mass defect is converted into energy according to the following equation.

\( E=\Delta{mc^2} \)

Where, ​\( \Delta{m} \)​ is mass defect, and  c= velocity of light.

(ii) Law of Definite or Constant Proportions

This law was given  by Proust in 1799. This law states “A compound always contains the same elements combined together in fixed proportion by mass and does not depend on its source and method of preparation” i.e., for example carbon-di-oxide can be obtained by using any one one of the following :

(a) By heating calcium carbonate.

(b) By heating sodium bicarbonate.

(c) By burning carbon in oxygen.

(d) By reacting calcium carbonate with hydrochloric acid.

What ever the sample of carbon-di-oxide is taken, it is observed that carbon and oxygen are always combined in the ratio of 12 :32 or 3 :8 .

The conserve of this law, that when same elements combine in the same proportion, the same compound will be formed, is not always true.For example, carbon, hydrogen and oxygen when combine in the ratio of 12 : 3 : 8 may form either ethyl alcohol (​\( C_2H_5OH \)​) or dimethyl ether (​\( CH_3OCH_3 \)​) under different experimental conditions.

LIMITATIONS : Isotopes of an element have different atomic masses. Thus, it is obvious to have same chemical compound with different compositions. For example:

\( CO_2 \)​ (Having ​\( C^{12} \)​ isotope) has C : O : : 12 : 32

\( CO_2 \)​ (Having ​\( C^{14} \)​isotope ) has C : O : : 14 : 32

Similarly, ​\( H_2O^{16} \)​and ​\( H_2O^{18} \)​  have H : O ratio as 1 : 8 and 1 : 9 respectively.

 

(iii) Law of Multiple Proportions

This law was given by Dalton in 1808.This law states “If two elements are combined to form more than one compound, then the different masses of one element which combines with a fixed mass of the other element, bear a simple ratio to one another.” For example :

Hydrogen and oxygen can be combined to form two elements

Water(​\( H_2O \)​)

Hydrogen 2 parts  Oxygen 16 parts

Hydrogen peroxide(​\( H_2O_2 \)​) 

Hydrogen 2 parts  Oxygen 32 parts

The masses of oxygens which combines with same mass of hydrogen in these two compounds bear a simple ratio 1 : 2.

Another example

Nitrogen forms five stable oxides :

\( NO_2 \)​ Nitrogen 28 parts oxygen 16 parts

\( N_2O_2 \)​  Nitrogen 28 parts  Oxygen 32 parts

\( N_2O_3 \)​ Nitrogen 28 parts  Oxygen 48 parts

\( N_2O_4 \)​ Nitrogen 28 parts  Oxygen 64 parts

\( N_2O_5 \)​ Nitrogen 28 parts  Oxygen 80 parts

The masses of oxygen which combines with same mass of nitrogen in the five compounds bear a ratio of 16 : 32 : 48 : 64 : 80 or 1 : 2 : 3 : 4 : 5

Another example

The carbon combines with oxygen to form two compounds, carbon monoxide and carbon dioxides

CO Carbon 12 parts  Oxygen 16 parts

\( CO_2 \)​ Carbon 12 parts  Oxygen 32 parts

The masses of oxygen bears a simple ratio of 1 : 2 or 16 : 32

LIMITATIONS : This law is not obeyed by non-stoichiometric compounds i.e., ​\( Fe_{0.98}O_1 \)​, ​\( TiO_{1.12} \)​, etc

 

(iv) Law of Reciprocal Proportions

This law was given by Richter in 1792. According to this law “When definite mass of an element A combines with two other element B and C to form two compounds . And if B and C also combines to form a compound, then their combining masses are in the same proportion OR bear a simple ratio to the masses of B and C which combines with constant mass of A ” For example :

Let assume Hydrogen (A) , Na (B) and Cl (C) elements for better understanding of law.

Hydrogen combines with sodium(Na) and chlorine(Cl) to form compounds NaH and HCl respectively.

NaH Sodium 23 parts  Hydrogen 1 part

HCl Chlorine 35.5 parts  Hydrogen 1 part 

AND

Sodium(Na) and Chlorine(Cl) combines to form another compound 

NaCl Sodium 23 parts  Chlorine 35.5 parts .These are the same parts which combines with 1 part of hydrogen in NaH and HCl respectively.

Another example

Hydrogen combines with sulphur and oxygen to form compounds

\( H_2S \)​ Hydrogen 2 parts   Sulphur 32 parts

\( H_2O \)​ Hydrogen 2 parts  Oxygen 16 parts

Thus, according to law sulphur should combine with oxygen in the ratio of 32 : 32 or 1 : 1.

 

The law of reciprocal proportions is a special case of a more general law, the law of equivalent masses, which can be stated as “In all chemical reactions, substances always react  in the ratio of their equivalent masses”

 

(v) Law of Gaseous Volumes 

This law is also known as law of combining volumes.This law is given by Gay-lussac in 1808. According to this law “gasses react with each other in the simple ratio of their volumes and if the product is also in gaseous state, the volume of the product also bears a simple ratio with the volumes of gaseous reactants when all volumes are measured under similar conditions of temperature and pressure”. For example :

\[ H_2(1~vol)~+~Cl_2(1~vol)~=~2HCl(1~vol)~~~ratio=1:1:2 \\2H_2(2~vol)~+~O_2(1~vol)~=~2H_2O(2~vol)~~~ratio=2:1:2 \\2CO(2~vol)~+~O_2(1~vol)~=~2CO_2(1~vol)~~~ratio=2:1:2 \\N_2(1~vol)~+~3H_2(3~vol)~=~2NH_3(2~vol)~~~ratio=1:3:2 \]

DALTON’S ATOMIC THEORY

The ultimate particles which were not further divisible are called atoms. The word atom is derived from Greek word “atomos” meaning indivisible. During the years 1803-1808 Dalton gives a theory called Dalton’s Atomic Theory.

The main points are as follows :

(i) Matter is made up of atoms which are extremely small particles, and cannot be sub-divided, or elements consist of minute, indivisible, indestructible particles called atoms.

(ii) Atoms of an element are identical to each other. They have the same mass and size.

(iii) Atoms of different elements differ in properties and have different masses and size.

(iv) Compounds are formed when atoms of different elements combine with each other in simple numerical ratios such as one-to-one , one-to-two,two-to-three, etc.

(v) Atoms cannot be created, destroyed or transformed into atoms of other elements.

(vi) The relative numbers and kinds of atoms are always the same in a given compound.

(vii) Atoms retain their identity in all chemical reactions. However, atom is the smallest particle which can take part in a chemical reaction.

LIMITATION OF DALTON’S ATOMIC THEORY : The theory convincingly explained the various laws of chemical combinations but it has a number of limitations.

(i) It fails to explain the cause of chemical combination.

(ii) It does not explain Gay-lussac’s law of combining gaseous volumes.

(iii) It does not give an idea about isotopes and isobars.

(iv) It fails to explain why atoms of different elements show different properties like mass, size etc.

(v) It does not explain the difference between an atom and a molecule.

The theory has undergone a modification with the modern concepts of structure of atom.Which are as follows:

(i) The atom is not supposed to be indivisible. The atom is not a simple particle but a complex one. Though atom is made up of various sub-atomic particles such as protons, neutron, electron, and also they take part in a chemical reaction.

(ii) Atoms of the element may not necessarily possess the same mass but possess the same atomic number and show similar chemical properties. (Isotopes)

(iii) Atoms of the different elements may possess the same mass but they always have different atomic numbers and differ in chemical properties.

(iv) Atoms of one elements cab be transmuted into atoms of other elements. (Isobars)

(v) In certain organic compounds like proteins, starch, cellulose, etc the ratio in which atoms of different molecules of sugar ​\( C_{12}H_{22}O_{11} \)​, the ratio of C, H,O is not simple and cannot be  called integrals. There are a number of compounds which do not follow the law of constant proportions. Such compounds are called non-stoichiometric compounds.

(vi) Mass and Energy are inter convertible.

(vii) Molecule is the smallest particle of matter which is capable of independent existence.

 

ATOMS, MOLECULES AND FORMULAE

As you know atom is the smallest particle of elements. The atom of hydrogen is smallest and lightest. Atom take part in chemical combination and remain indivisible. All atoms do not occur free in nature. Avogadro introduced the idea of another kind of particles called molecules. A molecule is another smaller particle of element bigger than atom which is stable and independently existence in nature.

A molecule of an element that consist of one atom only is called monoatomic molecule i.e. inert gases. Oxygen is unstable in atomic from but stable in molecular form, they are diatomic in nature also nitrogen, hydrogen , fluorine, chlorine, bromine, iodine are also diatomic in nature. Many more elements exist in nature which are more complex.

[Note : The atoms are components of molecules and molecules are the  components of elements or compound.]

Chemical Formula is a group of symbols of elements which represents one molecule of a substance also shows chemical composition of that substance.

A chemical formula is said to be correct if it shows two information

(i) Show the symbols of component elements to which compound is made up of.

(ii) It shows the combining ratio of atoms of elements of the compound. (number written in subscript)

Example : Nitric acid is a combination of Nitrogen , Hydrogen and Oxygen so the base formula is HNO and the combining ratio is 1 : 1: 3 therefore correct chemical formula is ​\( HNO_3 \)​.

IONS

Ions are different from  atoms and molecules as they are electrically charged example cations (positively charged)and anions (negatively charged).Ions come into existence only by gaining and losing electrons.In a ionic compound the number of positive charge on cation must balance the negative charge on anions. For example :

Calcium Nitrate consist of calcium and nitrate ions. Each calcium ion carries 2 unit positive charge while each nitrate ion carries 1 unit negative charge. Thus, to make electrically neutral compound two nitrate ions combine with one calcium ion and formula will be ​\( Ca(NO_3)_2 \)​or ​\( [Ca^{+2}~+~2NO_{3}^{-}] \)

 

Atomic And Molecular Mass

Atoms are too small to measure their absolute mass. But we can find there relative mass of different atoms if a small unit of mass is taken as standard.

The atomic mass of an elements can be defined as the number which indicates how many times the mass of one atom of the elements is heavier in comparison to the mass of one atom of hydrogen.

A= Atomic mass of an elements

A=​

\[ \dfrac{mass~of~one~atom~of~an~element}{mass~of~one~atom~of~hydrogen} \]

But in 1858, oxygen atom was adopted as a standard on account of the following reasons :

(i) It is much easier to obtain compounds of elements with oxygen than with hydrogen as oxygen is more reactive than hydrogen.

(ii) The atomic masses of most of the elements become approximately whole number but with hydrogen as standard the atomic masses of most of the elements are fractional.

The mass of one atom of oxygen is 16.0. Thus, atomic mass of an element is equal to

\[ =\dfrac{mass~of~one~atom~of~an~element}{\dfrac{1}{16}th~part~of ~mass~of~one~atom~of~oxygen} \]

OR

\[ =\dfrac{mass~of~one~atom~of~an~element}{mass~of~one~atom~of~oxygen}\times{16} \]

Oxygen as a standard, we find atomic mass of hydrogen is 1.008, sodium 22.991 , and sulphur 32.066

In 1961, the international union of chemist selected a new unit for expressing the atomic masses which is stable isotope of carbon (​\( ^{12}C \)​) with mass number of 12 as the standard.

Atomic mass of an element can be defined as the number which indicates how many times the mass of one atom of the element is heavier in comparison to ​\( \dfrac{1}{12}th \)​ part of the mass of one atom of carbon-12  (​\( ^{12}{C} \)​).

A=atomic mass of the element

\[ =\dfrac{mass~of~one~atom~of~the~element}{\dfrac{1}{12}part~of~the~mass~of~one~atom~of~carbon-12} \]

OR

\[ =\dfrac{mass~of~one~atom~of~the~element}{mass~of~one~atom~of~carbon-12}\times12 \]

 

Absolute atomic mass represents actual mass of the atom of the element.However, relative mass represents how the atom is heavier than ​\( \dfrac{1}{12}th \)​ mass of an atom of ​\( ^{12}C \)​ .Relative atomic mass is unit less quantity.

Relative atomic mass = ​\( \dfrac{Absolute~atomic~mass~in~amu}{amu} \)

Example : Relative atomic mass of oxygen is 16 where as absolute atomic mass is 16 amu.

ATOMIC MASS UNIT

The quantity ​\( \dfrac{1}{12} \)​ mass of an atom of carbon-12 (​\( ^{12}C \)​) is known as the atomic mass unit. Abbreviated as amu. The actual mass of carbon-12 is 1.9924​\( \times{10^{-23}} \)​g or 1.9924 ​\( \times{10^{-26}} \)​ kg .

Thus , 1 amu = ​\( \dfrac{1.9924\times{10^{-24}}}{12} \)

\( =1.66\times{10^{-24}} \)​ g or ​\( =1.66\times{10^{-24}} \)​ kg 

A=Atomic mass of an element

A=​\( \dfrac{mass~of~one~atom~of~element}{1~amu} \)

 

The actual mass of an atom of an element = the atomic mass of an element in amu ​\( \times \)​1.66​\( \times{10^{-24}} \)​ g

Example : Actual mass of hydrogen is

=1.008​\( \times \)​1.66​\( \times{10^{-24}} \)​  OR ​\( =1.6736\times{10^{-24}} \)​ g

 

Average Relative Mass (Average Atomic MASS)

Mostly elements found in nature with the mixture of their isotopes (atoms of same element having same atomic number but different atomic masses). For example Chlorine found in nature as a mixture of isotopes of Cl-35 (34.969 amu) and Cl-37 (36.966 amu) in the ratio of 75.53 % and 24.47 % respectively.

Therefore , Average Relative Mass of chlorine is calculated as :

Average relative mass = ​

\[ \sum\dfrac{\%~abundance}{100}\times{Atomic~mass(isotopic~mass)} \]

For chlorine

\[ =\dfrac{75.53}{100}\times{34.969}+\dfrac{24.47}{100}\times{36.966} \]

\[ = 35.457 \]

Also, ratio of ​\( _{17}C^{35} \)​ and ​\( _{17}C^{37} \)​ is 3 : 1. Thus, average relative mass of chlorine may be given as

 ​\( =\dfrac{3\times35+1\times{37}}{3+1} \)

\( =35.5 \)

Gram Atomic Mass or Gram Atom

When we express atomic mass of elements in gram then it is called gram atomic mass or gram atom.
Example : Atomic mass of oxygen is 16 and gram atomic mass is 16 g.
Gram atomic mass or Gram atom of every element consist of same number of atoms. That number is 6.022​\( \times{10^{23}} \)​ and this number is called Avogadro’s number.

Absolute mass of one oxygen atom is

= 16 amu 

=16​\( \times \)​1.66​\( \times{10^{-24}} \)​ g

Therefore, the mass of 6.022​\( \times{10^{23}} \)​ atoms of oxygen will be

\( =16\times{1.66\times{10^{-24}\times{6.022}\times{10^{23}}}} \)

\( =16 \)​ g (gram atomic mass)

Thus, gram atomic mass can be defined as the absolute mass in grams of 6.022​\( \times{10^{23}} \)​ atoms of any element.

 

Number of gram atomic mass of any element can be calculated as 

\[ =\dfrac{mass~of~the~element~(in~gram)}{atomic~mass~of~the~element(in~gram)} \]

Molecular mass

Just like an atom, molecules are also small particle of elements that can exist in nature alone, as they are stable.Just like atomic mass molecular mass is also represented in relative mass with respect to standard substance.

Molecular mass is a number which indicates how many times one molecule of a substance is heavier in comparison to ​\( \dfrac{1}{16}th \)​part of the mass of oxygen atom or ​\( \dfrac{1}{12}th \)​ part of the mass of one atom of carbon-12.

M= molecular mass

M=

\[ \dfrac{mass~of~one~molecule~of~the~substance}{\dfrac{1}{12}th~mass~of~one~atom~of~carbon-12 } \]

OR

M=

\[ \dfrac{mass~of~one~molecule~of~the~substance}{\dfrac{1}{16}th~mass~of~one~atom~of~oxygen~atom } \]

Mass of one molecule is equal to the sum of masses of atoms present in molecule.

Example : 1 molecule of water consist of 2 atoms of hydrogen and 1 atom of oxygen.Thus, molecular mass of water is

\( =(2\times{1.008}+16.0) \)

\( =18.016 \)​ amu

Another example

1 molecule of sulphuric acid ​\( (H_2SO_4) \)​  consist of 2 atoms of hydrogen and 1 atom of sulphur and 4 atoms of oxygen. Thus, molecular mass of sulphuric acid is 

\( =(2\times 1.008)+32.00+(4\times{16.00}) \)

\( =98.016 \)​ amu

Gram molecular mass

When we express molecular mass of elements in gram then it is called gram molecular mass or gram atom.

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