Class 11-Chapter 1(a) : Some Basic Concepts of Chemistry by Subhas C Chakra

 

Some Basic Concepts of Chemistry 

Importance of Chemistry

Chemistry has a direct impact on our life and has a  wide range of applications in different fields. A few of them are listed  below:

(A) In Agriculture and Food:

• For optimal plant growth, Chemistry has introduced various types of chemical fertilizers such as urea, calcium phosphate, sodium nitrate, ammonium phosphate, etc.
•  This has helped us protect our products from harmful insects and bacteria by using effective insecticides, fungicides and insecticides.
• The use of preservatives has helped us to extend the shelf life of foods and beverages such as jam, butter, squash, etc.

(B) In Health and Sanitation:

• It has provided mankind with a variety of life-saving drugs.
• Today, dysentery and pneumonia can be cured thanks to the discovery of life-saving sulfa drugs and penicillin drugs.
• Cisplatin and Taxol have been shown to be highly effective in treating cancer and AZT (azidothymidine) is used for AIDS victims.
•  Disinfectants such as phenol are used to kill microorganisms found in drains, toilets, floors, etc.
•  Low concentration of chlorine, i.e.  0.2 to 0.4 parts per million (ppm), used to sterilize water to make it suitable for drinking.

(C) Saving the Environment:

• Rapid industrialization around the world has caused a lot of pollution.
• Toxic gases such as CFCs and chemicals are constantly released into the atmosphere.
• They are polluting the environment at an alarming rate.
• Scientists are working day and night to develop alternatives that can reduce pollution.
• For example, methane (compressed natural gas), an alternative to gasoline, is very effective in screening for car pollution.

(D) Application in Industry:

• Chemistry has played an important role in the development of many consumer goods produced in industry.
• Commodities such as fertilizers, alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metal alloys and other inorganic and organic chemicals, including new materials, make a major contribution to strengthening the national economy.

• Matter

Matter is anything which has mass and hence occupies space. 

 
For example, A book,  a pen, soil, water, air are composed of matter. 

These these substances  have mass and also  they occupy space.

• Classification of Matter

There are two ways of classifying the matter:
(A) Physical classification
(B) Chemical classification

(A) Physical Classification:

Matter can exist in three physical states:

1. Solids 2. Liquids 3. Gases

1. Solids: The particles of solid are held very close to each other in an orderly fashion and there is not much freedom of movement for the particles.
Characteristics of solids: Solids have definite volume and definite shape.

2. Liquids: In liquids, the particles are so close to each other and still they are able to slide over one another. Their ability to flow gives them a name, Fluid.
Characteristics of liquids: Liquids have definite volume but no definite shape, as they adapt to the shape of the container they are put inside.

3. Gases: In gases, the particles are far apart as compared to those present in solid or liquid states. Their movement is easy, random  and fast.

Characteristics of Gases: Gases have neither definite volume nor definite shape. They completely occupy the container in which they are placed.

(B) Chemical Classification:

Based upon the composition, matter can be divided into two main types:
1. Pure Substances 

2. Mixtures.

1. Pure substances: 

A pure substance may be defined as a single substance (or matter) which cannot be separated by simple physical methods.

Pure substances can be further classified as (i) Elements (ii) Compounds

(i) Elements: 

•  An element consists of only one type of particles. These particles may be atoms or molecules.

• For example, Sodium, Copper, Silver, Hydrogen, Oxygen etc. are some examples of elements.

•  They all contain atoms of the same  type. 

• However, atoms of different elements are different in nature. 

• Some elements such as Sodium  or Copper contain single atoms held together as their constituent particles whereas in some others   two or more atoms combine to give molecules of the element. 

• As in Sulphur(S₈ ) and Phosphorous(P₄ ) there are 8 and 4 atoms respectively to form a Homoatomic molecule. 

• Thus, Hydrogen( H₂ ), Nitrogen( N₂ ) and Oxygen( O₂ ) gases consist of Homoatomic molecules in which two atoms combine to give the respective molecules of the element.

• The Elements have been classified into three types. These are:
• Metals
• Metalloids
• Nonmetals

•  Metals :

Metals  are  substances characterized by high electrical and thermal conductivity as well as by malleability, ductility, rigidity and high reflectivity of light.

• Approximately three-quarters of all known chemical elements are metals. 

• The most abundant varieties in the Earth’s crust are   aluminum,  iron,  calcium,  sodium,  potassium, and magnesium. 

• The vast majority of metals are found in ores (mineral-bearing substances), but a few such as copper, gold, platinum, and silver frequently occur in the free state because they do not readily react with other elements.

•   Semi Metals or Metalloids:

The general chemical properties described as metallic (or base forming),  metalloid (or amphoteric), and nonmetallic (or acid forming) are correlated with the periodic table in a simple way. 

The most metallic elements are to the left and to the bottom of the periodic table and the most nonmetallic elements are to the right and to the top (except the noble gases). 

The metalloids are adjacent to a diagonal line from boron to polonium.

•   Non metals:

• Like metals, Nonmetals may occur in the solid, liquid, or gaseous state.

• However, unlike metals, nonmetals display a wide range of both mechanical and optical properties, ranging from brittleness  to plasticity and from transparency to opaqueness.

From the  chemical point of view, nonmetals may be divided into two classes: 

1) Covalent substances:

Covalent substances contain atoms having small sizes, high electronegativities, low valence vacancy to electron ratios.

These atoms exhibit  a pronounced tendency to form negative ions in chemical reactions and negative oxidation states in their compounds; 

2) Ionic substances:

Ionic substances  contain both small and large atoms. 

• Ions may be formed by adding electrons to (small, electronegative atoms) or by extracting electrons from (large, electropositive) atoms.
• Ions in Ionic substances are bound to each other by strong electrostatic forces of attraction.

(ii) Compounds: 

A Compound  may be defined as a pure substance containing two or more elements combined together in a fixed proportion by weight and can be decomposed into these elements by suitable chemical methods. 

Moreover, the properties of a compound are altogether different from the constituting elements.

The compounds have been classified into two types. These are:

(i) Inorganic Compounds: 

•  These are compounds which are obtained from non-living sources such as rocks and minerals. 
• A few examples are: Common salt, marble, gypsum, washing soda etc.

(ii) Organic Compounds are the compounds which are present in plants and animals. 

All the organic compounds have been found to contain carbon as their essential constituent.

 For example, carbohydrates, proteins, oils, fats etc.

2. Mixtures: 

The combination of two or more elements or compounds which are not chemically combined together and may also be present in any proportion, is called mixture. 

A few examples of mixtures are: milk, sea water, petrol, lime water, paint glass, cement, wood etc.

Types of mixtures: Mixtures are of two types:

(i) Homogeneous mixtures: 

A mixture is said to be homogeneous if it has a uniform composition throughout and there are no visible boundaries of separation between the constituent particles. 

For example: 

A mixture of sugar solution in water has the same sugar water composition throughout and all portions have the same sweetness.

(ii) Heterogeneous mixtures: 

A mixture is said to be heterogeneous if it does not have uniform composition throughout and has visible boundaries of separation between the various constituent particles. 

The different constituents of a heterogeneous mixture can be seen even with naked eye.

For example: 

When iron filings and Sulphur powder are mixed together, the mixture formed is heterogeneous. 

It has greyish-yellow appearance and the two constituents, Iron and Sulphur, can be easily identified with naked eye.

• Differences between Compounds and Mixtures

Compounds

1. In a compound, two or more elements are combined chemically.

2.  In a compound, the elements are present in the fixed ratio by mass. This ratio cannot change.

3. Compounds are always homogeneous i.e., they have the same composition throughout.

4 In a compound, constituent particles cannot be separated by physical methods.

5. In a compound,  the constituent particles lose their identities i.e.,  a compound does not show the characteristics of the constituting elements.

Mixtures

1.  In a mixture, or more elements or compounds are simply mixed and not combined chemically.

2. In a mixture the constituent particles are not present in fixed ratio. The ratio  can vary.

3. Mixtures may be either homogeneous or heterogeneous in nature.

4. Constituent particles of mixtures can be separated by physical methods.

5, In a mixture, the constituent particles do not lose their identities,  i.e., a mixture shows the characteristics of all the constituents .

We have discussed the physical and chemical classification of matter. A flow chart representation of the same is given below.

• Properties of Matter and Their Measurements

Physical Properties:   These are the  properties which can be measured or observed without changing the identity or the composition  of  the  substance.

 Some of the examples of physical properties are colour, odour, melting point, boiling point etc.

 Chemical Properties:

 It requires a chemical change to occur. 

The examples of chemical properties are characteristic reactions of different substances.

 These include acidity, basicity, combustibility etc.

• Units of Measurement

Fundamental Units: 

The quantities mass, length and time are called fundamental quantities and their units are known as fundamental units.

There are seven basic units of measurement for the quantities: length, mass, time, electric current, temperature, ,  luminous intensity and amount of substance.

S.I.  System: This system of measurement is the most common system employed throughout the world.
It has given units of all the seven basic quantities listed above.

• Definitions of Basic SI Units
1. Metre: 

    It is the length of the path travelled by light in vacuum during 

a time interval of 1/299792458 of a second.

2. Kilogram: 

  It is the unit of mass. It is equal to the mass of the international 

prototype of the kilogram.

  
3. Second: 

It is the duration of 9192631, 770 periods of radiation which 

correspond to the transition between the two hyper fine levels 

of the ground state of cesium- 133 atom.

4. Kelvin: 

It is the unit of thermodynamic temperature and is equal to 

1/273.16 of the thermodynamic temperature of the triple 

point of water.

5. Ampere: 

The ampere is that constant current which if maintained in 

two straight parallel conductors of infinite length, of negligible 

circular cross section and placed, 1 metre apart in vacuum, 

would produce between these conductors a force equal to 

2 x 10⁻⁷N per metre of length.

6. Candela: 

It may be defined as the luminous intensity in a given direction, 

from a source which emits monochromatic radiation of frequency 

540 x 10⁻⁻¹² Hz and that has a radiant intensity in that direction 

of 1/ 683 watt per steradian.

7. Mole: It is the amount of substance which contains as many 

elementary entities as there are atoms in 0.012 kilogram of 

carbon -12 (C¹²). Its symbol is ‘mol’.

• Mass and Weight

Mass: 

Mass of a substance is the amount of matter present in it. 

 The mass of a substance is constant.

The mass of a substance can be determined accurately in the 

laboratory by using an analytical balance. 

SI unit of mass is kilogram.

Weight : 

It is the force exerted by gravity on an object. Weight of 

substance may vary from one place to another due to 

change in gravity.

Weight can be measured using a spring balance.

Another form of Spring balance looks like this :

Volume: 

 Volume means the space occupied by matter.
 
It has the units of (length)³.
 
In SI units, volume is expressed in metre³  (m³).

 
However, a popular unit of measuring volume, is litre (L).

 
But it is not an SI unit .

Mathematically,
1L = 1000 mL = 1000 cm³ = 1dm³.

Volume of liquids can be measured by different devices

 
like Burette, Pipette, Cylinder, Measuring flask etc.

 

• All of them have been calibrated.

Temperature: There are three scales in which temperature can be measured. These are known as Celsius scale (°C), Fahrenheit scale (°F) and Kelvin scale (K).

•  Thermometers with Celsius scale are calibrated from 0°C to 100°C.
•  Thermometers with Fahrenheit scale are calibrated from 32°F to 212°F.
•  Kelvin  scale of temperature is S.I. scale and is very common these days. Temperature on this scale is shown by the sign K.

•   The temperature on two scales are related to each other by the relationship :

Idea of the ranges of temperature :

Density: Density of a substance is its amount of mass per unit volume. So, SI unit of density can be obtained as follows:

This unit is quite large and a chemist often expresses density in g cm⁻³ where mass is expressed in gram and volume is expressed in cm⁻³.

• Uncertainty in Measurements
All scientific measurements involve certain degree of error or uncertainty. The errors which arise depend upon two factors.
(i) Skill and accuracy of the worker (ii) Limitations of measuring instruments.

• Scientific Notation
It is an exponential notation in which any number can be represented in the form N x 10n where n is an exponent having positive or negative values and N can vary between 1 to 10. Thus, 232.508 can be written as 2.32508 x 102 in scientific notation.
Now let us see how calculations are carried out with numbers expressed in scientific notation.
(i) Calculation involving multiplication and division

(ii) Calculation involving addition and subtraction: For these two operations, the first numbers are written in such a way that they have the same exponent. After that, the coefficients are added or subtracted as the case may be. For example,

• Significant Figures
Significant figures are meaningful digits which are known with certainty. There are certain rules for determining the number of significant figures. These are stated below:
1. All non-zero digits are significant. For example, in 285 cm, there are three significant figures and in 0.25 mL, there are two significant figures.
2. Zeros preceding to first non-zero digit are not significant. Such zeros indicates the position of decimal point.
For example, 0.03 has one significant figure and 0.0052 has two significant figures.
3. Zeros between two non-zero digits are significant. Thus, 2.005 has four significant figures.
4. Zeros at the end or right of a number are significant provided they are on the right side of the decimal point. For example, 0.200 g has three significant figures.
5. Counting numbers of objects. For example, 2 balls or 20 eggs have infinite significant figures as these are exact numbers and can be represented by writing infinite number of zeros after placing a decimal.
i.e., 2 = 2.000000
or 20 = 20.000000

• Addition and Subtraction of Significant Figures
In addition or subtraction of the numbers having different precisions, the final result should be reported to the same number of decimal places as in the term having the least number of decimal places.

  For example,  If we  carry out the addition of three numbers 3.52, 2.3 and 6.24, having different precisions or different number of decimal places.  

The final result has two decimal places but it has to be expressed with  only up to one decimal place, i.e., the answer would be 12.0.

                                                                                                                                Subtraction of numbers can be done in the same way as the addition. If we substract 15.29  from 28.3470 ,then the final result is 13.0570.

The final result has four decimal places. But it has to be expressed with  only up to two decimal places, i.e., the answer would be 13.05.

• Multiplication and Division of Significant Figures
In the multiplication or division, the final result should be reported up to the same number of significant figures as present in the least precise number.
Multiplication of Numbers: 

2.2120 x 0.011 = 0.024332

According to the rule the final result = 0.024

Division of Numbers: 

4.2211÷3.76 = 1.12263

According to the rule the final result = 1.12

• Dimensional Analysis
Often while calculating, there is a need to convert units from one system to other. The method used to accomplish this is called factor label method or unit factor method or dimensional analysis.

• Laws of Chemical Combinations

The combination of elements to form compounds is governed by the following five basic laws.
(i) Law of Conservation of Mass
(ii) Law of Definite Proportions
(iii) Law of Multiple Proportions
(iv) Law of Gaseous Volume (Gay Lussac’s Law)
(v) Avogadro’s Law

(i) Law of Conservation of Mass
The law was established by a French chemist, A. Lavoisier. The law states:
In all physical and chemical changes, the total mass of the reactants is equal to that of the products.
In other words, matter can neither be created nor destroyed.
The following experiments illustrate the truth of this law. 

 For example, decomposition of mercuric oxide.
(a) When matter undergoes a physical change.

                             

 

It is found that there is no change in weight though a physical change has taken place.

(b) When matter undergoes a chemical change. 

 For example, decomposition of mercuric oxide

 During the above decomposition reaction, matter is neither gained nor lost.

(ii) Law of Definite Proportions :
According to this law:
A pure chemical compound always consists of the same elements combined together in a fixed proportion by weight.
For example, Carbon dioxide may be formed in a number of ways i.e.,

(iii) Law of Multiple Proportions :

According to this law:

If two elements combine to form two or more compounds, the weight of one of the elements which combines with a fixed weight of the other in these compounds, bears simple whole number ratio by weight.
 For example,

    

(iv) Gay Lussac’s Law of Gaseous Volumes
The law states that, 

Under similar conditions of temperature and pressure, whenever gases combine, they do so in volumes which bear simple whole number ratio with each other and also with the gaseous products. 

The law may be illustrated by the following examples.

(a) Combination between Hydrogen and Chlorine:

 

(b) Combination between Nitrogen and Hydrogen: The two gases lead to the formation of ammonia gas under suitable conditions. The chemical equation is

                        

(v) Avogadro’s Law: 

Avogadro's law states that, 

Equal volumes of gases at the same temperature and pressure  contain equal number of molecules.

 For example,
During the reaction of hydrogen and oxygen to produce water, we see that two volumes of hydrogen combine with one volume of oxygen to give two volumes of water without leaving any unreacted oxygen. The ratio of the volumes in this case is found to be a simple whole number ratio.

• Dalton’s Atomic Theory: 

In 1808, Dalton published ‘A New System of Chemical Philosophy’ in which he proposed the following:

• Indivisibility of atoms : Matter consists of indivisible atoms. 
• Identical properties : All the atoms of a given element have identical properties including identical mass.
• Uniqueness of Atoms : Atoms of different elements differ in mass.
• Definite proportions : Compounds are formed when atoms of different elements combine in a fixed ratio.
• Conservation of matter : Chemical reactions involve reorganization of atoms. In a chemical reaction, these are neither created nor destroyed .

• Atomic Mass
The atomic mass of an element is the number of times an atom of that element is heavier than an atom of carbon taken as 12. It may be noted that the atomic masses as stated above are the relative atomic masses and not the actual masses of the atoms.
One atomic mass unit (a m u) is equal to l/12th of the mass of an atom of carbon-12 isotope. It is also known as unified mass.

Average Atomic Mass
Most of the elements exist as isotopes which are different atoms of the same element with different mass numbers and the same atomic number. Therefore, the atomic mass of an element must be its average atomic mass.  

Hence, it may be defined as the average relative mass of an atom of an element as compared to the mass of carbon atoms (C¹²) taken as 12 a.m.u.(or 12 u) .

Molecular Mass
Molecular mass is the sum of atomic masses of the elements present in a molecule. It is obtained by multiplying the atomic mass of each element by number of its atoms and adding them together.
For example,
Molecular mass of methane (CH₄)
= 12.011 u + 4 (1.008 u)
= 16.043 u

 Hence, Molecular  mass  may also be defined as the average relative mass of a  molecule of an element as compared to the mass of carbon atoms (C¹²) taken as 12 a.m.u.(or 12 u) .

Equivalent Mass 

Equivalent weight (also known as gram equivalent) is the mass of   one equivalent, that is the mass of a given substance which will combine with or displace a fixed quantity of another substance.  

The mass of a substance especially in grams that combines with or is chemically equivalent to 1.8 grams of Hydrogen, 8 grams of oxygen or 35.5 grams of Chlorine , the atomic or molecular weight divided by the valency of an element also gives the Equivalent Mass. 

 Hence, Equivalent  mass  of a substance is defined as the mass of a  substance which combines with or displaces 1.008 gram of hydrogen or 8.0 grams of oxygen or 35.5 grams of chlorine.  

The equivalent weight of an element is the mass which combines with or displaces 1.008 gram of hydrogen or 8.0 grams of oxygen or 35.5 grams of chlorine. These values correspond to the atomic weight divided by the usual valence;  for oxygen as example that is 16.0 g / 2 = 8.0 g..

Formula Mass
Ionic compounds such as NaCl, KNO3, Na2C03 etc. do not consist of molecules i.e., single entities but exist “as ions are closely packed together in a three dimensional space as shown in -Fig. 1.5. 

In such cases, the formula is used to calculate the formula mass instead of molecular mass. Thus, formula mass of NaCl = Atomic mass of sodium + atomic mass of chlorine
= 23.0 u + 35.5 u = 58.5 u.

 Hence, Formula  mass  of a substance may also be defined as the mass of a  molecule of the substance which is the sum of the atomic masses of the atoms present in one of its molecules . 

• Mole Concept :

One mole of a substance may be defined as the unit linked to an Avogadro number(N🇦 = 6.022 × 1023 )of particles of matter which may be atoms, ions or molecules of the substance. .

•  One gram atom of any element is found to contain  the same number of atoms and one gram molecule of any substance contains the same number of molecules. 
• This number has been experimentally found to be equal to 6.022137 x 1023 . 
• It is generally called Avogadro’s number or Avogadro’s constant.
• It is usually represented by N🇦 .
• Avogadro’s Number, N🇦 = 6.022 × 1023
• One mole of a substance at NTP weighs equal to the Molecular Mass of the substance.
• One mole of a substance at NTP occupies a volume  equal to the Gram Molecular Volume (22.4 L ) of the substance.
• One mole of a substance at NTP weighs equal to the molar mass of the substance.
• One mole of molecules of a substance at NTP weighs equal to the Molecular mass of the substance.(One gram molecule or GMM)
• One mole of atoms of a substance at NTP weighs equal to the Atomic mass of the substance. (One gram atom or GAM)
• One mole of a substance at NTP contains Avogadro number of particles. 

• Percentage Composition
 One can check the purity of a given sample by analyzing this data. 

 

Let us understand by taking the example of water (H20). Since water contains hydrogen and oxygen, the percentage composition of both these elements can be calculated as follows:

Percentage Composition: The percentage composition of different ingredients of mixtures like solute and solvents in solutions can be calculated as follows.

e 

• Empirical Formula
The formula of the compound which gives the simplest whole number ratio of the atoms of various elements present in one molecule of the compound.
For example, the formula of Hydrogen peroxide is H202. 

In order to express its empirical formula, we have to take out a common factor 2. The simplest whole number ratio of the atoms is 1:1 and the empirical formula is HO. 

Similarly, the formula of Glucose is C6H1206. 

In order to get the simplest whole number of the atoms,
Common factor = 6
The ratio is = 1 : 2 : 1 The empirical formula of glucose = CH20

• Molecular Formula
The formula of a compound which gives the actual ratio of the atoms of various elements present in one molecule of the compound.
For example, 

Molecular formula of hydrogen peroxide = H202 and of  Glucose = C6H1206
Molecular formula = n x Empirical formula
Where n is the common factor and also called multiplying factor. The value of n may be 1, 2, 3, 4, 5, 6 etc.
If case n= 1,  then 

Molecular formula of a compound = Empirical formula of the compound.

• Stoichiometry and Stoichiometric Calculations
The word ‘stoichiometry’ is derived from two Greek words—Stoicheon (meaning element) and metron (meaning measure). Stoichiometry, thus deals with the calculation of masses (sometimes volume also) of the reactants and the products involved in a chemical reaction. 

Let us consider the Formation of Ammonia from Nitrogen and Hydrogen. A balanced equation for this reaction is as given below:

Limiting Reactant/Reagent
Sometimes, in alchemical equation, the reactants present are not the amount as required according to the balanced equation. The amount of products formed then depends upon the reactant which has reacted completely. This reactant which reacts completely in the reaction is called the limiting reactant or limiting reagent. The reactant which is not consumed completely in the reaction is called excess reactant.

Reactions in Solutions
When the reactions are carried out in solutions, the amount of substance present in its given volume can be expressed in any of the following ways:
1. Mass percent or weight percent (w/w%)
2. Mole fraction
3. Molarity
4. Molality

1. Mass percent: It is obtained by using the following relation:

Example 1

Example 2

2. Mole fraction: It is the ratio of number of moles of a particular component to the total number of moles of the solution.

 

For a solution containing n₁ moles of the solute dissolved in n₂  moles of the solvent,

3. Molarity: It is defined as the number of moles of solute in 1 litre of the solution.

  4. Molality: It is defined as the number of moles of solute present in 1 kg of   solvent. It is denoted by m.

Summary :

• All substances contain matter which can exist in three states — solid, liquid or gas.
• Matter can also be classified into elements, compounds and mixtures.
• Element: An element contains particles of only one type which may be atoms or molecules.
• Compounds: Compounds are formed when atoms of two or more elements combine in a fixed ratio to each other.
• Mixtures: Mixtures are the substances present around us which contain two or more pure substances in different ratios, combined physically or chemically.
• Scientific notation: The measurement of quantities in chemistry are spread over a wide range of 10-31to 1023. Hence, a convenient system of expressing the number in scientific notation is used.
• Scientific figures: The uncertainty is taken care of by specifying the number of significant figures in which the observations are reported.
• Dimensional analysis: It helps to express the measured quantities in different systems of units.

• Laws of Chemical Combinations are:
(i) Law of Conservation of Mass
(ii) Law of Definite Proportions
(iii) Law of Multiple Proportions
(iv) Gay Lussac’s Law of Gaseous Volumes
(v) Avogadro’s Law.
• Atomic mass: The atomic mass of an element is expressed relative to 12C isotope of carbon which has an exact value of 12u.
• Average atomic mass: Obtained by taking into account the natural abundance of different isotopes of that element.
• Molecular mass: The molecular mass of a molecule is obtained by taking sum of atomic masses of different atoms present in a molecule.
• Avogadro number: The number of atoms, molecules or any other particles present in a given system are expressed in terms of Avogadro constant.
= 6.022 x 1023
• Balanced chemical equation: A balanced equation has the same number of atoms of each element on both sides of the equation.
• Stoichiometry: The quantitative study of the reactants required or the products formed is called stoichiometry. Using stoichiometric calculations, the amounts of one or more reactants required to produce a particular amount of product can be determined and vice-versa.