NEET Chapter - Some Basic Concepts of Chemistry

Science is described as a constant human effort to organise knowledge in order to describe and comprehend nature. Two examples of changes that occur include the development of curd from milk, as well as the production of sugarcane juice vinegar after a long period of time and the rusting of iron. This is a common occurrence for us. To keep things as simple as possible. Chemistry, biology, geology, and other sciences are among the branches of science. The study of a substance's composition, properties, structure, and reactions is referred to as chemistry. The study of material substances is referred to as chemistry.

Important Topics of Some Basic Concepts of Chemistry

• Properties of Matter and their Measurements

• Uncertainty of Measurement 

•  Percentage composition

• Importance of chemistry 

• Dalton's Atomic Theory

• Stoichiometry and Stoichiometric calculations

• Nature of Matter 

• Atomic and molecular masses

• Mole concept and molar masses

1. Importance of Chemistry

Chemistry is a vital branch of science that is frequently intertwined with other fields. Chemical principles can be used to weather patterns, brain function, computer operation, chemical industry production, fertiliser production, alkalis, acids, salts, dyes, polymers, pharmaceuticals, soaps, detergents, metals, alloys, and new material.

2. Nature of Matter

Matter is defined as everything that has mass and takes up space. Everything around us is made up of matter, including books, pens, pencils, water, air, and all living things.

1. States of Matter

• Anything that occupies space, has mass, and whose presence can be felt by any one or more of our five senses is classified as matter.

• Solid, liquid, and gas are the three physical states in which matter can exist.

• Solid - a substance is said to be solid if it has a fixed volume and shape, such as sugar, iron, gold, or wood.

• Liquid- If a substance has a specific volume but no distinct shape, it is said to be liquid. Water, milk, oil, mercury, alcohol, and other liquids take on the shape of the vessel in which they are placed.

• Gaseous- If a substance has neither a specific volume nor a definite shape, it is said to be gaseous. This is due to the fact that they are completely full.

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2. Classification of Matter

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• Mixtures: A mixture contains two or more substances (in any ratio) that are referred to as its constituents.

• A mixture might be homogeneous or heterogeneous.

• Homogeneous mixture- a homogeneous mixture is one in which all of the components are the same, components are entirely mixed together, and The composition is consistent throughout, i.e. it is made up of only one type of material.  As a result, sugar solution and air are two examples of homogenous mixes are a type of mixture in which all of the ingredients are the same

• Heterogeneous mixtures- A heterogeneous mixture is one in which the constituents are heterogeneous. The composition is not consistent throughout. There are several phases that can be seen. Grains, for example,  Pulses, as well as some soil (typically stone) fragments, are used.

• Pure substances :- A material containing only one substance is called a pure substance.

• A pure substance containing only one type of particle is defined as an element. The elements are further classified based on their physical and chemical properties. Metals are separated into three categories: (1) Metals, (2) Non- Metals, and (3) Metalloids.

• A compound is a pure material that contains other pure substances, two or more parts combined in a fixed configuration.

3. Properties of Matter and their Measurements

• Physical attributes are those that may be measured or observed without affecting the substance's identity or composition. Colour, odour, melting point, boiling temperature, density, and other physical qualities are examples.

• Chemical properties are those in which the substance  must undergo a chemical change. Chemical characteristics include acidity or basicity, combustibility, and other characteristic responses of different compounds.

• Physical quantities refer to all of the quantities we encounter during our scientific investigations. Any physical quantity is evidently measured in two parts: 

• (1) the number, and (2) the unit.

•  A unit is a unit of measurement that is used to measure any physical amount.

• The 11th General Conference on Weights and Measures developed the International System of Units (in French, Le Systeme International d'Unités - shortened as SI) (CGPM from Conference Generale des Poids at Measures).

• The CGPM is an intergovernmental treaty body established by the Metre Convention, a diplomatic convention signed in Paris in 1875.

1. Mass and Weight

• The mass of a material is the amount of matter it contains, whereas the weight of a thing is the force exerted by gravity.

• A substance's mass is constant, but its weight can fluctuate from one location to another due to changes in gravity.

• The kilogram is the SI unit of mass (kg). The newton is the SI derived unit (a unit derived from SI base units) for weight.

2. Volume

• Volume is the amount of three-dimensional space enclosed by a closed boundary, such as the space occupied or contained by a substance (solid, liquid, gas, or plasma). The SI derived unit, the cubic metre, is frequently used to quantify volume numerically.

3. Density

• A material's mass density, sometimes known as density, is defined as its mass per unit volume. The most common symbol for density is (the lowercase Greek letter rho). 

kg m–3 is the SI unit for density.

4. Temperature

• The term "temperature" refers to a physical attribute of matter that quantifies the concepts of hot and cold. Temperature is measured using three different scales: °C (degrees Celsius), °F (degrees Fahrenheit), and K (degrees Kelvin) .

°F = 9/5 (°C) + 32  

K = °C + 273.15 

4. Uncertainty in Measurement

• Significant figures are a combination of significant digits that are known with certainty and one that is estimated or questionable. By printing the certain numerals and the last doubtful digit, the uncertainty is communicated. Thus, if a result is written as 11.2 mL, we can state that the 11 is certain and the 2 is unknown, with the uncertainty being +1 in the last digit. An uncertainty of +1 in the last digit is always assumed unless otherwise stated.

• The number of significant figures is determined according to a set of rules. These are as follows:

1. All digits that are not zero are significant. For example, there are three significant figures in 285 cm and two significant figures in 0.25 mL.

2. The zeros before the first non-zero digit are unimportant. The position of the decimal point is indicated by a zero. As a result, 0.03 has one significant figure, while 0.0052 has two.

3. There are significant zeros between two non-zero digits. As a result, 2.005 is a four-digit number.

4. If the zeros at the end or right of a number are on the right side of the decimal point, they are significant.

• Precision refers to how near different measurements for the same quantity are to each other. 

• Accuracy, on the other hand, is the agreement of a given value with the true value of the result.

• When working with numbers, it's common to need to convert units from one system to another. Factor label method, unit factor method, or dimensional analysis are the methods used to do this.

5. Laws of Chemical Combinations

a. Law of Conservation of Mass

"Mass is conserved in a chemical reaction because the mass of reactants used and the mass of products generated are the same." This is a direct result of the atomic conservation rule. In 1789, Antoine Lavoisier proposed this law.

b. Law of Constant / Definite Proportions

The proportions in which two or more elements combine to form a compound stay constant and are unaffected by the compound's source. Joseph Proust, a French chemist, gave this law.

c. Law of Multiple Proportions

The ratio of masses of one element that combines with a fixed mass of the other element in the two compounds is a simple whole number ratio when two elements mix to generate two or more compounds. Dalton proposed this legislation in 1803.

d. Law of Reciprocal Proportions

When three elements combine in a combination of two to make three compounds, the ratio of the masses of the two elements combining with the fixed mass of the third and the ratio in which they combine are both simple whole number ratios. Richter enacted this law in 1792.

e. Gay Lussac’s Law of Gaseous Volumes

Gay Lussac gave this law in 1808. When gases join or are generated in a chemical process, he discovered that they do so in a simple volume ratio if all gases are at the same temperature and pressure.

f. Avogadro Law

Avogadro hypothesised in 1811 that equal quantities of gases at the same temperature and pressure contain the same number of molecules.

6. Dalton's Atomic Theory

Dalton wrote 'A New System of Chemical Philosophy' in 1808, in which he proposed:

1. Matter is made up of indestructible atoms.

2. The properties of all atoms of a given element are the same, including their mass. The mass of atoms in different elements varies.

3. When atoms of different elements mix in a specific ratio, compounds are produced.

4. Chemical processes entail atom reconfiguration. In a chemical reaction, they are neither generated nor destroyed.

7. Atomic and Molecular Masses

• The mass of one carbon atom is 12 to one-twelfth of the mass of one carbon atom - 12 atoms. Moreover, 1 amu = 1.6605610–24 g.

• ‘Amu’ has been replaced by ‘u’, which is known as a unified mass.

• The total of the atomic masses of the atoms in a molecule is called molecular mass. It's calculated by multiplying each element's atomic mass by the number of atoms in its nucleus and adding the results.

8. Mole Concept and Molar Masses

• The mole (sign mol) is the SI unit of material quantity. There are exactly 

• 6.02214076 x1023 elementary entities in one mole.

• This is the Avogadro number, which is the fixed numerical value of the Avogadro constant, NA, as represented in the unit mol–1. A system's amount of substance, denoted by the symbol n, is a count of the number of defined elementary things.

• The molar mass of a substance is the mass of one mole in grammes. The atomic/molecular/formula mass in u is mathematically equal to the molar mass in grams.

9. Percentage Composition

• The percentage composition of both these elements can be calculated as follows: 

• Mass % of an element  = mass of that element in the compound/ molar mass of the compound ×100

• The molecular formula displays the exact number of distinct types of atoms present in a molecule of a compound, whereas the empirical formula shows the simplest whole number ratio of various atoms present in a compound.

10. Stoichiometry and Stoichiometric Calculations

• Stoichiometry is the study of chemical processes and the calculations that go along with them. Stoichiometric Coefficients are the coefficients that are utilised to balance the reaction.

• If the reactants are not mixed in stoichiometric proportions, the reactant that is present in less than the needed amount affects how much product is created and is known as the Limiting Reagent, while the reactant that is present in excess is known as the Excess Reagent.

• Mole fraction “X “ is defined as the moles of a component / Total moles of solution.

• Mass% is defined as  Mass of solute (in g) present in 100g of solution.

• Molarity (M) is defined as moles of solute / volume of solution (L).

• Molality (m)  is defined as  moles of solute /mass of solvent (kg).

Solved Examples from the Chapter

Example 1: How many electrons are present in 1.6 g of methane ?

Sol. Gram-molecular mass of methane (CH4 ) = 12g + 4g = 16 g

Number of moles in 1.6 g of methane = 1.6/ 16 = 0.1

Number of molecules of methane in 0.1 mole are 

= 0.1 × 6.02 × 1023

= 6.02 × 1022

Key Point to Remember: Need to remember the Mole concept and is defined as the given mass/ molar mass.

Example 2: Classify the following mixtures as homogeneous and heterogeneous.

(i) Air (ii) Smoke (iii) Petrol (iv) Sea water (v) iodized table salt (vi) Aerated water (vii) Mixture of sand and common salt (viii) Gun powder (ix) Milk (x) Muddy water.

Sol. Homogeneous mixture- a homogeneous mixture is one in which all of the components are the same, components are entirely mixed together, and The composition is consistent throughout, i.e. it is made up of only one type of material.  As a result, sugar solution and air are two examples of homogenous mixes, a type of mixture in which all of the ingredients are the same. Examples like-  Air; Petrol; Iodized table salt; Sea water; Aerated water; Milk.

Heterogeneous mixtures areHeterogeneous mixtures- A heterogeneous mixture is one in which the constituents are heterogeneous. The composition is not consistent throughout. There are several phases that can be seen.Examples like- Smoke; Gun powder; Mixture of sand common salt; Muddy water.

Key Point to Remember: Mixtures are categorised as the homogenous and heterogenous mixtures.

Solved Examples from Previous Year Questions

Question 1: Avogardro’s number is the number of molecules present in 

(a) 1 L of molecule 

(b) 1 g of molecule

(c) gram molecular mass 

(d) 1 g-atom of molecules

Solution: The Avogadro constant is defined as the number of constituent particles per mole of a given substance and that is equal to  gram molecular mass of a substance . It is equal to 6.022×1023. Therefore, option (c) is the answer.

Trick: Need to remember the Mole concept in this specially the molar mass concept.

Question 2: 720 g water contains, the number of moles

(a) 2 

(b) 190

(c) 40

(d) 55

Solution: Given, mass of H2O=720 g

We know that, molar mass of water = 18g/mol

Moles of water = Molar mass of water/ Mass of water= 720/18 = 40 moles.

Therefore,  option (c) is the answer.

Trick: Need to remember the Mole concept and is defined as the given mass/ molar mass.

Question 3: A molal solution is one that contains one mole of a solute in

(a) 1000 g of the solvent

(b) one litre of the solvent

(c) one litre of the solution

(d) 22.4 litres of the solution

Solution: Molality is  defined as the number of moles of solute dissolved in 1000 grams of solvent. So, 1 molal solution contains one mole of solute in 1000 grams of solvent.

​Therefore,  option (a) is the answer

Trick: Need to remember the concentration terms and socially the molality case.

Practice Questions

1. Which among the following is the heaviest ?

(a) One mole of oxygen

(b) One molecule of sulphur trioxide

(c) 100 amu of uranium

(d) 44g of carbon dioxide

Answer: (d) 44g of carbon dioxide

Question 2: Choose the wrong statement :

(a) 1 mole means 6.02 × 1023 particles

(b) Molar mass is mass of one molecule

(c) Molar mass is mass of one mole of a substance

(d) Molar mass is molecular mass expressed in grams

Answer: (b) Molar mass is mass of one molecule

Conclusion

Chemists study the properties and structure of substances, as well as how they change through time. All substances contain matter, which can exist in three states: solid, liquid, or gas. The component particles are preserved in various ways and exhibit their own characteristics in these states of matter. All matter is made up of elements, compounds, and mixtures. A single sort of particle, such as atoms or molecules, makes up an element. When atoms from two or more elements combine in a certain ratio, compounds are formed. Many of the substances in our environment are mixtures, and mixtures are common.