Some Basic Concepts of Chemistry Class 11 Notes For Chemistry Chapter 1

SOME BASIC CONCEPTS OF CHEMISTRY
Chapter 1 Class 11 Download in pdf

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Notes of class 11 chemistry chapter 1

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Chemistry is the branch of science that deals with the composition, structure and properties of matter. Chemistry is called the science of atoms and molecule.

Branches of Chemistry

1. Organic  Chemistry  - This  branch  deals  with  study  of  carbon  compounds especially hydrocarbons and their derivatives.
2. Inorganic Chemistry-This branch deals with the study of compounds of all other elements except carbon. It  largely  concerns  itself with the  study of minerals found in the Earth’s crust.
3. Physical  Chemistry-The  explanation  of  fundamental  principles  governing various chemical phenomena is the main concern of this branch. It is basically concerned with laws and theories of the different branches of chemistry.
4. Industrial Chemistry-The chemistry involved in industrial processes is studied under this branch.
5. Analytical Chemistry-This branch deals with the qualitative and quantitative analysis of various substances.
6. Biochemistry-This branch deals with the chemical changes going on in the bodies of living organisms; plants and animals.
7. Nuclear Chemistry-Nuclear reactions, such as nuclear fission, nuclear fusion, transmutation processes etc. are studied under this branch.

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PROPERTIES OF MATTER AND THEIR MEASUREMENT

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Every substance has unique or characteristic properties. These properties can be classified into two categories – physical properties and chemical properties.

Physical properties are those properties which can be measured or observed without changing the identity or the composition of the substance. E.g. colour, odour, melting point, boiling point, density etc.The measurement or observation of chemical properties requires a chemical change occur. e.g. Burning of Mg-ribbon in air

Chemical properties are characteristic reactions of different substances; these include acidity or basicity, combustibility etc.Many properties of matter such as length, area, volume, etc., are quantitative in nature.

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The International System of Units (SI) (Class 11 chemistry notes- study rankers)

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The International System of Units (in French Le Systeme International d’Unites– abbreviated as SI) was established by the 11th General Conference on Weights and Measures (CGPM fromConferenceGenerale des Poids at Measures). The SI system has seven base units

Metric System was based on the decimal system.

Mass and Weight- Mass of a substance is the amount of matter present in it while weight is the force exerted by gravity on an object. The mass of a substance is constant whereas its weight may vary from one place to another due to change in gravity. The mass of a substance can be determined very accurately by using an analytical balance

Volume- Volume has the units of (length)³. So volume has units of m³  or cm³  or dm³.A common unit, litre (L) is not an SI unit, is used for measurement of volume of liquids. 1 L = 1000 mL, 1000 cm³ = 1 dm³

Density- Density of a substance is its amount of mass per unit volume.SI unit of density = SI unit of mass/SI unit of volume = kg/m³ or kg m^–3This unit is quite large and a chemist often expresses density in g cm^–3.

Temperature-There are three common scales to measure emperature — °C (degree celsius), °F (degree Fahrenheit) and K (kelvin). Here, K is the SI unit.

K = °C + 273.15

Note- Temperature below 0 °C (i.e. negative values) are possible in Celsius scale but in Kelvin scale, negative temperature is not possible.

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Scientific Notation (Topper notes class 11 chemistry)

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Scientific Notation- In which any number can be represented in the form N × 10ⁿ Where n is an exponent having positive or negative values and N can vary between 1 to 10). e.g. We can write 232.508 as 2.32508 x10² in scientific notation. Similarly, 0.00016 can be written as 1.6 x 10^–4.

Precision refers to the closeness of various measurements for the same quantity.

Accuracy is the agreement of a particular value to the true value of the result

Significant Figures
The  reliability of  a  measurement  is  indicated  by  the  number  of  digits  used  to represent it. To express it more accuratelywe express it with digits that are known with certainty. These are called as Significant figures. They contain all thecertain digits plus one doubtful digit in a number.

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Rules for Determining the Number of Significant Figures

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1. •  All  non-zero  digits  are  significant.  For  example,  6.9  has  two  significant figures, while 2.16 has three significantfigures. The decimal place does not determine the number of significant figures.
2. • A zero becomes significant in case it comes in between non zero numbers. For example, 2.003 has four significantfigures, 4.02 has three significant figures.
3. •  Zeros at the beginning of a number are not significant. For example, 0.002 has one significant figure while 0.0045has two significant figures.
4. •  All zeros placed to the right of a number are significant. For example, 16.0 has three significant figures, while 16.00has four significant figures. Zeros at the end of a number without decimal point are ambiguous.
5. •  In  exponential  notations,  the  numerical  portion  represents  the  number  of significant figures. For example, 0.00045 isexpressed as 4.5 x 10^-4 in terms of scientific notations. The number of significant figures in this number is 2, while inAvogadro’s number (6.023 x 10²³) it is four.
6. •  The decimal point does not count towards the number of significant figures For example, the number 345601 has sixsignificant figures but can be written in different ways, as 345.601 or 0.345601 or 3.45601 all having same number
   ofsignificant figures.

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Retention of Significant Figures - Rounding off Figures

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The rounding off procedure is applied to retain the required number of significant figures.

1. If the digit coming after the desired number of significant figures happens to be more than 5, the precedingsignificant figure is increased by one, 4.317 is rounded off to 4.32.
2. If the digit involved is less than 5, it is neglected and the preceding significant figure remains unchanged, 4.312 isrounded off to 4.31.
3. If the digit happens to be 5, the last mentioned or preceding significant figure is increased by one only in case ithappens to be odd. In case of even figure, the preceding digit remains unchanged. 8.375 is rounded off to 8.38 while8.365 is rounded off to 8.36.

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Dimensional Analysis (Vedantu class 11 chemistry notes)

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Dimensional Analysis During calculations generally there is a need to convert units from one system to other. This is called factor label method or unit factor method or dimensional analysis.   

For example- 5 feet and 2 inches (height of an Indian female) is to converted in SI unit

1 inch = 2.54 x 10-2 m
 
then, 5 feet and 2 inch = 62 inch

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Physical Classification of Matter

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Chemical Classification of Matter

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Elements

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An element is the simplest form of matter that cannot be split into simpler substances or built from simpler substances by any ordinary chemical or physical method. There are 114 elements known to us, out of which 92 are naturally occurring while the rest
have been prepared artificially. Elements are further classified into metals, non-metals and metalloids.

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Compounds

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A compound is a pure substance made up of two or more elements combined in a definite proportion by mass, which could be split by suitable chemical methods. Compounds are  broadly classified  into  inorganic and  organic  compounds. Inorganic compounds are those, which areobtained from non-living sources such as minerals. For example, common salt, marble and limestone. Organiccompounds are those, which occur in living sources such as plants and animals. They all contain carbon. Commonorganic compounds are oils, wax, fats etc.

Characteristics of compound

1. • Compounds always contain a definite proportion of the same elements by mass.
2. • The properties of compounds are totally different from the elements from which they are formed.
3. • Compounds are homogeneous.

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Mixtures

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A mixture is a combination of two or more elements or compounds in any proportion so that the components do not lose their identity. Air is an example of a mixture Mixtures are of two types, homogeneous and heterogeneous.

Heterogeneous mixtures - have the same composition throughout the sample. The components of such mixtures cannot be seen under a powerful microscope. They are also called solutions. Examples of homogeneous mixtures are air, seawater, gasoline, brass etc.
Heterogeneous mixtures - consist of two or more parts (phases), which have different compositions. These mixtures  have visible boundaries of separation between the different constituents and can be seen with the naked eye e.g., sand and salt, chalk powder in water etc.

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LAWS OF CHEMICAL COMBINATIONS  

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Law of Conservation of Mass (Given by Antoine Lavoisier in 1789). It states that matter (mass) can neither be created nor destroyed.

Law of Definite Proportions or Law of Constant Composition: This law was proposed by Louis Proust in 1799, which states that: ’A chemical compound always consists of the same elements combined together in the same ratio, irrespective of the method of preparation or the source from where it is taken’.

Law of Multiple Proportions Proposed by Dalton in 1803, this law states that:’ When two elements combine to form two or more compounds, then the differe nt masses of one element, which combine with a fixed mass of the other, bear a simple ratio to one another’.

Gay Lussac’s Law of Gaseous Volumes (Given by Gay Lussac in 1808.) According to this law when gases combine or are produced in a chemical reaction they do so in a simple ratio by volume provided all gases are at same temperature and pressure.

Example

H2(g) +	+	Cl2(g)	→	2HCl(g)
1V		1V		2V

All reactants and products have simple ratio 1:1:2

Avogadro Law (In 1811, Given by Avogadro)
According to this law equal volumes of gases at the same temperature and pressure should contain equal number of molecules.

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Dalton’s Atomic Theory

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1. • All substances are made up of tiny, indivisible particles called atoms.Atoms  of  the  same  element  are  identical  in  shape,  size,  mass  and  other properties.
2. • Atoms of different elements are different in all respects.
3. • Atom is the smallest unit that takes part in chemical combinations.  
4. •Atoms  combine  with  each  other  in  simple  whole  number  ratios  to  form compound atoms called molecules.
5. •Atoms cannot be created, divided or destroyed during any chemical or physical change.

Atoms and Molecules

The  smallest  particle  of  an  element,  which  may  or  may  not  have  independent existence is called an atom, while the  smallest particle of a substance which is capable of independent existence is called a molecule.
Molecules are classified as homoatomic and heteroatomic. Homoatomic molecules are made up of the atoms of the same element and heteroatomic molecules are made up of the atoms of the different element have different atomicity (number of atoms in a molecule of an element) like monoatomic, diatomic, triatomic and polyatomic.

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Atomic , Unit, Gram Atomic, Molecular and Formula Mass

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Atomic Mass Unit

One atomic mass unit is defined as a mass exactly equal to one twelfth the mass of one carbon -12 atom. And 1 amu = 1.66056×10^–24 g.
Today, ‘amu’ has been replaced by ‘u’ which is known as unified mass.

Atomic Mass

Atomic mass of an element is defined as the average relative mass of an atom of an element as compared to the mass of an atom of carbon -12 taken as 12.

Gram Atomic Mass

The quantity of an element whose mass in grams is numerically equal to its atomic mass. In simple terms, atomic mass of an element expressed in grams is the gram atomic mass or gram atom. For example, the atomic mass of oxygen = 16 amu Therefore gram atomic mass of oxygen = 16 g

Molecular Mass

Molecular mass of a substance is defined as the average relative mass of its molecule as compared to the mass of an atom of C-12 taken as 12. It expresses as to how many times the molecule of a substance is heavier than 1/12th of the mass of an atom of carbon.
For example, a molecule of carbon dioxide is 44 times heavier than 1/12th of the mass of an atom of carbon. Therefore the molecular mass of CO2 is 44 amu.
It is obtained by adding the atomic masses of all the atoms present in one molecule.

Gram Molecular Mass

A quantity of substance whose mass in grams is numerically equal to its molecular mass is called gram molecular mass. In simple terms, molecular mass of a substance expressed in grams is called gram molecular mass.
e.g., the molecular mass of oxygen = 32 amu
Therefore, gram molecular mass of oxygen = 32 g

Formula Mass-

Sum of atomic masses of the elements present in one formula unit of a compound. It is used for the ionic compounds.

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Mole Concept

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Mole Concept - Mole is defined as the amount of a substance, which contains the same number of chemical units (atoms, molecules, ions or electrons) as there are atoms in exactly 12 grams of pure carbon-12.
A mole represents a collection of 6.022 x10²³( Avogadro’s number) chemical units..The mass of one mole of a substance in grams is called its molar mass.

Molar Volume

The volume occupied by one mole of any substance is called its molar volume. It is denoted by Vm. One mole of all gaseous substances at 273 K and 1 atm pressure occupies a volume equal to 22.4 litre or 22,400 mL. The unit of molar volume is litre
per mol or millilitre per mol

PERCENTAGE COMPOSITION—

The mass percentage of each constituent element present in any compound is called its percentage composition
Mass % of the element= (Mass of element in 1 molecule of the compound x100) / Molecular mass of the compound

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Empirical Formula and Molecular Formula

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An empirical formula represents the simplest whole number ratio of various atoms present in a compound. E.g. CH is the empirical formula of benzene.
The molecular formula shows the exact number of different types of atoms present in a molecule of a compound. E.g. C6H6 is the molecular formula of benzene.

Relationship between empirical and molecular formulae
The two formulas are related as Molecular formula = n x empirical formula

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Chemical Equation

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Shorthand representation of a chemical change in terms ofsymbols and formulae of the substances involved in the reaction is called chemical equation.
The substances that react among themselves to bring about the chemical changes are known as reactants, whereas the substances that are produced as a result of the chemical change, are known as products

Limiting Reagent- The reactant which gets consumed first or limits the amount of product formed is known as limiting reagent
Reactions in Solutions-- The concentration of a solution can be expressed in any of the following ways.

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Measuring of Concentration

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1.    Mass Percent is the mass of the solute in grams per 100 grams of the solution.

A 5 % solution of sodium chloride means that 5 g of NaCl is present in 100g of the solution.

2.     Volume percent is the number of units of volume of the solute per 100 units of the volume of solution.

A 5 % (v/v) solution of ethyl alcohol contains 5 cm³ of alcohol in 100 cm³ of the solution

3.     Molarity of the solution is defined as the number of moles of solute dissolved per litre (dm³) of the solution. It isdenoted by the symbol M. Measurements in Molarity can change with the change in temperature because solutionsexpand or contract accordingly.
Molarity of the solution =    Mass of the solute/( Molar mass of the solute X volume of the solution in liter)

Formula :-

Where
W= Mass of the solute
M= Molar mass of the solute
V= volume of the solution in liter)

Molarity equation- To calculate the volume of a definite solution required to prepare solution of other molarity, the following equation is used: M1V1 = M2V2,

where M1= initial molarity, M2= molarity of the new solution, V1= initial volume and V2= volume of the new solution.

4.    Molality- Molality is defined as the number of moles of solute dissolved per 1000 g (1 kg) of solvent. Molality is expressed as ’m’.

5.     Mole Fraction is the ratio of number of moles of one component to the total number of moles (solute and solvents) present in the solution. It is expressed as’x’.

Mole fraction of the solute =  Moles of the solute/( Moles of solute  +  Moles of solvent)

Mole fraction of the solvent =Moles of the solvent/( Moles of solute  +  Moles of solvent)

Always Check

Mole fraction of the solute + Mole fraction of solvent = 1