Plus One Chemistry Notes Chapter 1 Some Basic Concepts of Chemistry is part of Plus One Chemistry Notes. Here we have given Kerala Plus One Chemistry Notes Chapter 1 Some Basic Concepts of Chemistry.
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|Chapter Name||Some Basic Concepts of Chemistry|
|Category||Plus One Kerala|
Kerala Plus One Chemistry Notes Chapter 1 Some Basic Concepts of Chemistry
Chemistry is the science of molecules and their transformations. Chemistry is the branch of science that studies the composition, structure, properties and interaction of matter. Chemistry plays a central role in science and is often interwined with other branches of science.
Nature Of Matter
Matter is anything that occupies space and posses mass. Matter can exist in three physical states- solid, liquid and gas. At the macroscopic or bulk level, matter can be classified as mixtures or pure substances.
A material containing only one substance is called a pure substance. Materials
containing more than one substance are called mixtures. Pure substances are further classified into two types: elements and compounds. Mixtures are also of two types; homogeneous mixtures
A mixture is said to be homogeneous if it has same composition throughout. Some examples of homogeneous mixtures are air, gasoline, kerosene, milk, alloys etc. Heterogeneous mixtures are the mixtures which have different composition in different parts. Some examples of heterogeneous mixtures are cereal and milk, rocks in sand at the beach, muddy water etc.
Properties of Matter and their Measurement
Every substance has its characteristic properties. These are 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. Some examples of physical properties are colour, odour, melting point, boiling point, density etc.
Chemical properties are characteristic reactions of different substances; these include acidity or basicity, combustibility etc.
The International System of units (SI unit)
The SI unit of temperature is kelvin (K). There are three common scales to measure temperature – °C (degree Celsius), °F (degree fahrenheit scale between 32° to 212°) and Kelvin. The relationship between fahrenheit and Celsius scale is
The Kelvin scale is related to Celsius scale by the relation K = °C + 273.15.
Accuracy and Precision
Precision refers to the closeness of various measurements for some quantity, eg., an exact number is absolutely precise but the measured mass of object is not absolutely precise. Accuracy refers to the closeness of a particular value to the true value of result.
The precision in the measurement of mass of an object depends upon the:
- sensitivity of the measuring apparatus.
- observation power (the skill) of the person making the measurement.
Laws of Chemical Combinations
There are 5 laws governing the formation of chemical compounds. The first four laws deals with combination by mass and are applicable to solids, liquids and gases. The fifth law is related to combination by volume and is applicable to gases only.
Law of conservation of mass
The law put forth by Antoine Lavoisier. The law states that matter can neither be created nor I be destroyed. In any chemical reaction the total mass of the reactants is equal to the total mass of the products.
Here in (1), 12g + 32g = 44g and in (2), 2g + 16g= 18g ; Clearly proves the above law.
This law is also called the law of indestructibility of matter.
Law of Definite Proportions
This law was given by, a French chemist, Joseph Proust. He stated that a given compound always contains exactly the same proportion of elements by weight. It is sometimes also referred to as Law of definite composition.
Carbon dioxide can be obtained by different methods such as burning of carbon, heating lime stone, by the action of dil. HCI on marble etc. It has been observed that each sample of C02 contains carbon and oxygen combined in the ratio 3:8 by mass. This means that the composition of a compound is always the same respective of the method by which it is prepared.
Law of Multiple Proportions
This law was proposed by Dalton in 1803. According to this law, if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element, are in the ratio of small whole numbers.
Carbon (C) can combine with oxygen (O) to form more than one compound, namely CO, C02. Here ratio of masses of O that combine with fixed mass of C is 16:32 or 1:2
Law of Reciprocal or Equivalent proportions: Ritcher
This law states that if an element A which combines with an element B in the ratio x:y and combines with another element C in the ratio x:z, then B and C combine they in the ratio y:z or a simple multiple of it.
An element A (hydrogen) combine with element B (carbon) to form CH4 in the ratio 4:12 or 1:3.
An element A (hydrogen) combine with element C (oxygen) to form H2o in the ratio 2:16 or 1:8. .
Then element B (carbon) combine with element C (oxygen) to form Co2 in the ratio 12:32 or 3:8. By experiment, it is found that in the formation of Co2, carbon combines with oxygen in the ratio 12:32 or 3:8.
Gay Lussac’s Law of Gaseous Volume
When gases react to form gaseous products, there exists a simple whole number ratio between the volumes of the reactants and the products at constant temperature and pressure.
Avogadro’s Law or Hypothesis
Avogadro’s law states that equal volumes of gases at the same temperature and pressure should contain equal number of molecules, mathematically Vα N at constant temperature. The volume of gas is directly proportional to number of molecules (N) at constant temperature and pressure. The number of molecules (N) in a given volume of any gas is directly proportional to its number of moles (n). Then V α n at constant temperature and pressure.
22.4 L of every gas at STP (Standard Temperature and Pressure, ie., T = 273 K, P = 1 atm) contains equal number of molecules, which is equal to 6.022 x 1023.
Dalton’s Atomic Theory
The main postulates of the theory are:
- Matter is made up of extremely small, indivisible particles called atoms.
- Atoms of the same element are identical in all respects, ie, size and mass
- Atoms of different elements are different,i.e, they possess different sizes, shapes, masses and chemical properties.
- Atoms of different elements may combine with each other in a simple whole number ratio to form compound atoms or molecules.
- Atoms can neither be created nor destroyed. i.e., atoms are indestructable.
Atomic mass and Molecular mass
The atomic mass or mass of an atom is actually too small because atoms are extremely small.
Actual mass of hydrogen atom is 1.67 x 1024 g, but it not much also used in chemical calculation. The relative mass of an atom or masses of other atoms are expressed in relation to the mass of a standard carbon atom or C-12 is known as atomic mass.
Atomic mass unit is defined as exactly of the mass of a carbon-12 atom. It is represented as amu. [Now a new symbol ‘u’ called unified mass is used]
∴ Mass of 6.02 x 1023 atoms of C12 = 12 g
= 1.660 Χ 10-24 g
Molecular mass is the sum of atomic masses of the elements present in a molecule.
Molecular mass of water = 2 x atomic
mass of hydrogen + 1 x atomic mass of oxygen
= 2 x (1.008 u) + 1 x 16.00 u = 18.02 u
In ionic compounds we use formula mass instead of molecular mass. Formula mass of an ionic compound is the sum of the atomic masses of all atoms in a formula unit of compound.
Mole concept and Molar Masses
‘ Mole’ was introduced as the seventh base quantity for the amount of substance in SI system. One mole of a substance contains many particles and their number is equal to the number of particles in 12 g of the 12C isotope. This number is known as avogardo constant (NA = 6.022 x 1023). The mass of one mole of a substance in grams is called its molar mass. The molar mass is numerically equal to atomic/ molecular/formula mass in u.
Calculate the number of moles of iron in a sample containing 1.0 x 1022 atoms.
Number of Moles
= 0.0166 mol
Mole concept in Gaseous reaction
Molar volume is the mole related to volume of gaseous substance. Since 1 mole of all gases contain 6.02 x 1023 molecules, they occupy the same volume under similar
condition of temperature and pressure. The volume occupied by 1 mol of a gaseous substance is called molar volume. 1 mole occupies 22.414 L or 22414 ml at STP ie., 273 K and 1 atm.
Find the volume of 34 g of NH3 at STP.
Molecular weight of NH3 =14 + 3 x 1 = 17
1 mole of NH3 = 22.4 L
Molar Mass .
The mass of 1 mol of a substance in grams is called its molar mass.
Molar mass of NaCI = 58.5 g [When 58.5 g of NaCI (1 molar) is dissolved in one litre of solution].
Number of moles of Ag in 16.8 g of
Calculate the number of He atoms in
i. 52 u
ii. 52 g
iii. 52 moles of He
Atomic wt. of He is 4 u
i. 4 u is mass of 1 atom
⇒ 52 u is mass of 1/4 x 52= 13 atoms
ii. 4 g of He contains 6.022 x 1023 He atoms
⇒ 52 g of He contains
⇒ 7.8286 x 1024 atoms
iii. 1 mole of He contains 6.022 x 1023 atoms
52 moles of He contains 52x 6.022x 1023
= 3.131 x 1025atoms
Percentage Composition :
Mass percentage of an element
The number of atoms present in one mole of an element is equal to Avogadro number. Which of the following element contains the greatest number of atoms?
a. 4g He
b. 46g Na
c. 0.40g Ca
d. 12g He
For comparing number of atoms, first we calculate the moles as all are monoatomic
Determination of Chemical formulae
Empirical and Molecular formula.
- Empirical formula represent the simplest whole number ratio of various atoms present in a compound, eg., EF of benzene (C6H6) is CH.
- Molecular formula shows the exact number of different types of atoms present in a molecule of a compound, eg., MF of benzene is C6H6.
- Relation between empirical and molecular
formula, MF = (EF) x n;
where n is a simple whole number.
Chemical equation and Balanced chemical equation
Chemical equation is a scientific method of representing a chemical change in terms of symbols and formula of reactants and products involved in it
A chemical equation contains an equal number of atoms of each element in the reactants and products is called a balanced chemical equation. A correct chemical equation must be in accordance with law of conservation of mass.
Balancing a Chemical Equation Steps:
- Determine the correct formulas for all the reactants and products in the reaction.
- Write the formulas for the reactants on the left and the formulas for the products on the right with an arrow in between. If two or more reactants or products are involved, separate their formulas with plus signs.
- Count the number of atoms of each element in the reactants a products.
A polyatomic ion appearing unchanged on both sides of the equation is counted as a single unit.
- Balance the elements one at a time by using coefficients.
- Check each atom or polyatomic ion to be sure that the equation is balanced.
- Finally, make sure that all the coefficients are in the lowest possible ratio.
potassium chlorate gives potassium chloride and oxygen.
- KCIO3 ⇒ KCI + [O] ; unbalanced equation
- KCIO3⇒ KCI + 3[O] ; unbalanced equation
- 2KCIO3⇒ 2KCI + 3O2 ; balanced equation
Stoichiometry and Stoichiometric Calculations
Stoichiometry deals with the calculation of masses and sometimes volumes also. The coefficients of reactants and products in a balanced chemical equation is called the stoichiometric coefficients.
- Write the balanced chemical equation.
- Observe the number of moles of various reactant taking part in the reaction and products formed.
- Calculate the number of moles or amount of substance formed.
We calculate the amount of potassium chlorate required to prepare 2.4 moles of oxygen as follows:
2KCI03(s)→2KCI(s)+ 3O2 (g) is balanced equation.
3 moles of O2 is prepared from 2 moles of 2KCIO3.
2.4 moles of O2 is prepared from 2/4 x 2.4 = 1.6 moles of KCIO3.
Mass of 1.6 moles of KCI03
= No. of moles x Molar mass
= 1.6 mole x [39 + 35.5 + 3 x 16] g/mol
= 1.6 mole x 122.5 g/mol = 196.00 g
Limiting reagent is defined as the reactant which is completely consumed during reaction
Calculate the mass of water formed when 3 g of H2 reacts with 29 g of O2 to form water.
Reactions in Solutions
It is the ratio of number of moles of a particular component to the total number of moles of the solution.
Mole fraction of component A,
It is the most widely used unit and is denoted by M. It is defined as the number of moles of the solute in 1 litre of the solution.
It is defined as the number of moles of solute present in 1 kg of solvent. It is denoted by m.
Note: Molarity of a solution changes with temperature. But molality of a solution does not change with temperature since mass remains uneffected with temperature.
Calculate the molarity of NaOH in the solution prepared by dissolving its 4 gin enough water to form 250 ml of the solution.
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