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Chapter
Major Concepts
1.1 Branches of Chemistry 1.2 Basic Definitions 1.3 Chemical species 1.4 Avogadro' s Number and Mole 1.5 Chemical Calculations
Students Learning Outcomes
Students will be able to:
- Identify and provide examples of different branches of chemistry.
- Differentiate among branches of chemistry.
- Distinguish between matter and a substance.
- Define ions, molecular ions, formula units and free radicals.
- Define atomic number, atomic mass, atomic mass unit.
- Differentiate among elements, compounds and mixtures.
- Define relative atomic mass based on C-12 scale.
- Differentiate between empirical and molecular formula.
- Distinguish between atoms and ions.
- Differentiate between molecules and molecular ions.
- Distinguish between ion and free radicals.
- Classify the chemical species from given examples.
- Identify the representative particles of elements and compounds.
- Relate gram atomic mass, gram molecular mass and gram formula mass to mole.
- Describe how Avogadro's number is related to a mole of any substance.
- Distinguish among the terms gram atomic mass, gram molecular mass and gram formula mass.
- Change atomic mass, molecular mass and formula mass into gram atomic mass, gram molecular mass and gram formula mass.
Fundamentals of Chemistry
Time allocation
Teaching periods 12
Assessment periods 03
Weightage 10%
Introduction The knowledge that provides understanding of this world and how it works, is science. The branch of science which deals with the composition, structure, properties and reactions of matter is called chemistry. It deals with every aspect of our life. The development of science and technology has provided us a lot of facilities in daily life. Imagine the role and importance of petrochemical products, medicines and drugs, soap, detergents, paper, plastics, paints and pigments, insecticides, pesticides which all are fruit of the efforts of chemists. The development of chemical industry has also generated toxic wastes, contaminated water and polluted air around us. On the other hand, chemistry also provides knowledge and techniques to improve our health and environment and to explore and to conserve the natural resources. In this chapter, we will study about different branches of chemistry, basic definitions and concepts of chemistry.
1.1 BRANCHES OF CHEMISTRY It is a fact that we live in the world of chemicals. We all depend upon different living organisms which require water, oxygen or carbon dioxide for their survival. Today chemistry has a wide scope in all aspects of life and is serving the humanity day and night. Chemistry is divided into following main branches: physical chemistry, organic chemistry, inorganic chemistry, biochemistry, industrial chemistry, nuclear chemistry, environmental chemistry and analytical chemistry.
1.1.1 Physical Chemistry Physical Chemistry is defined as the branch of chemistry that deals with the relationship between the composition and physical properties of matter along with the changes in them. The properties such as structure of atoms or formation of molecules behavior of gases, liquids and solids and the study of the effect of temperature or radiation on matter are studied under this branch.
1.1.2 Organic Chemistry Organic Chemistry is the study of covalent compounds of carbon and hydrogen (hydrocarbons) and their derivatives. Organic compounds occur naturally and are also synthesized in the laboratories. Organic chemists determine the structure and properties of these naturally occurring as well as synthesized compounds. Scope of this branch covers petroleum, petrochemicals and pharmaceutical industries.
1.1.3 Inorganic Chemistry Inorganic chemistry deals with the study of all elements and their compounds except those of compounds of carbon and hydrogen (hydrocarbons) and their derivatives. It has applications in every aspect of the chemical industry such as glass, cement, ceramics and metallurgy (extraction of metals from ores).
Test yourself 1.
1.2 BASIC DEFINITIONS
Matter is simply defined as anything that has mass and occupies space. Our bodies as well as all the things around us are examples of matter. In chemistry, we study all types of matters that can exist in any of three physical states: solid, liquid or gas.
A piece of matter in pure form is termed as a substance. Every substance has a fixed composition and specific properties or characteristics. Whereas, impure matter is called a mixture ; which can be homogeneous or heterogeneous in its composition.
We know that every substance has physical as well as chemical properties. The properties those are associated with the physical state of the substance are called physical properties like colour, smell, taste, hardness, shape of crystal, solubility, melting or boiling points, etc. For example, when ice is heated, it melts to form water. When water is further heated, it boils to give steam. In this entire process only the physical states of water change whereas its chemical composition remains the same.
The chemical properties depend upon the composition of the substance. When a substance undergoes a chemical change, its composition changes and a new substances are formed. For example, decomposition of water is a chemical change as it produces hydrogen and oxygen gases. All materials are either a substance or a mixture. Figure 1. shows simple classification of the matter into different forms.
i. In which branch of chemistry behaviour of gases and liquids is studied? ii. Define biochemistry? iii. Which branch of chemistry deals with preparation of paints and paper? iv. In which branch of chemistry the metabolic processes of carbohydrates and proteins are studied? v. Which branch of chemistry deals with atomic energy and its uses in daily life? vi. Which branch of chemistry deals with the structure and properties of naturally occurring molecules?
Mixture Substance
MATTER
Mixture Substance
MATTER
Homogeneous mixture
Hetrogeneous mixture
Homogeneous mixture
Heterogeneous mixture ElementsElements^ CompoundsCompounds
Fig. 1.1. Classification of matter
‘s crust
1.2.1 Elements, Compounds and Mixtures
1.2.1.1 Elements
In the early ages, only nine elements (carbon, gold, silver, tin, mercury, lead, copper, iron and sulphur) were known. At that time, it was considered that elements were the substances that could not be broken down into simpler units by ordinary chemical processes. Until the end of nineteenth century, sixty-three elements had been discovered. Now 118 elements have been discovered, out of which 92 are naturally occurring elements. Modern definition of element is that it is a substance made up of same type of atoms, having same atomic number and cannot be decomposed into simple substances by ordinary chemical means. It means that each element is made up of unique type of atoms that have very specific properties.
Elements occur in nature in free or combined form. All the naturally occurring elements found in the world have different percentages in the earth's crust, oceans and atmosphere. Table 1.1. shows natural occurrence in percentage by weight of some major elements around us. It shows concentrations of these major elements found in the three main systems of our environment.
Table 1.1 Natural Occurrences by Weight % of Some Major Elements
Elements may be solids, liquids or gases. Majority of the elements exist as solids e.g. sodium, copper, zinc, gold, etc. There are very few elements which occur in liquid state e.g. mercury and bromine. A few elements exist as gases e.g. nitrogen, oxygen, chlorine and hydrogen.
On the basis of their properties, elements are divided into metals, non-metals and metalloids. About 80 percent of the elements are metals.
Table 1.2 Some Elements and Radicals with their Symbols and Common
Valencies
Some elements show more than one valency, i.e. they have variable valency. For example, in ferrous sulphate (FeSO ) the valency of iron is 2. In ferric sulphate 4 (Fe (SO ) ), the valency of iron is 3. Generally, the Latin or Greek name for the element 2 4 3 (e.g., Ferrum) is modified to end in 'ous' for the lower valency (e.g. Ferrous) and to end in 'ic' for the higher valency (e.g. Ferric).
1.2.1.2 Compound
Compound is a substance made up of two or more elements chemically combined together in a fixed ratio by mass. As a result of this combination, elements lose their own properties and produce new substances (compounds) that have entirely different properties. Compounds can't be broken down into its constituent elements by simple physical methods. For example, carbon dioxide is formed when elements of carbon and oxygen combine chemically in a fixed ratio of 12:32 or 3:8 by mass. Similarly, water is a compound formed by a chemical combination between hydrogen and oxygen in a fixed ratio of 1:8 by mass.
Compounds can be classified as ionic or covalent. Ionic compounds do not exist in independent molecular form. They form a three dimensional crystal lattice, in which each ion is surrounded by oppositely charged ions. These oppositely charged ions attract each other very strongly, as a result ionic compounds have high melting and boiling points. These compounds are represented by formula units e.g. NaCl, KBr, CuSO. 4
The covalent compounds mostly exist in molecular form. A molecule is a true representative of the covalent compound and its formula is called molecular formula e.g. H O, HC1, H SO , Ch. 2 2 4 4
Table 1.3 Some Common Compounds with their Formulae
1.2.1.3 Mixture
When two or more elements or compounds mix up physically without any fixed ratio, they form a mixture. On mixing up, the component substances retain their own chemical identities and properties. The mixture can be separated into parent components by physical methods such as distillation, filtration, evaporation, crystallisation or magnetization. Mixtures that have uniform composition throughout are called homogeneous mixtures e.g. air, gasoline, ice cream. Whereas, heterogeneous mixtures are those in which composition is not uniform throughout e.g. soil, rock and wood.
Always use: Standard symbols of elements Chemical formulae of compounds Proper abbreviations of scientific terms Standard values and SI units for constants
Remember
Hence, each element has a specific atomic number termed as its identification number. For example, all hydrogen atoms have 1 proton, their atomic number is Z =l. All atoms in carbon have 6 protons, their atomic number is Z =6. Similarly, in oxygen all atoms have 8 protons having atomic number Z =8 and sulphur having 16 protons shows atomic number Z = 16.
The mass number is the sum of number of protons and neutrons present in the nucleus of an atom. It is represented by symbol ' A '.
It is calculated as A=Z+n where n is the number of neutrons. Each proton and neutron has lamu mass. For example, hydrogen atom has one proton and no neutron in its nucleus, its mass number A =l+0 =1. Carbon atom has 6 protons and 6 neutrons, hence its mass number A =12. Atomic numbers and mass numbers of a few elements are given in Table 1.
Example 1.
How many protons and neutrons are there in an atom having A = 238 and Z = 92.
Solution:
First of all, develop data from the given statement of the example and then solve it with the help of data. Data A= Z= Number of protons? Number of neutrons? Number of protons = Z = 92
Table 1.5 Some Elements along with their Atomic and Mass Numbers
Number of Neutrons = A - Z = 238 – 92 = 146 1.2.3 Relative Atomic Mass and Atomic Mass Unit As we know that the mass of an atom is too small to be determined practically. However, certain instruments enable us to determine the ratio of the atomic masses of various elements to that of carbon-12 atoms. This ratio is known as the relative atomic mass of the element. The relative atomic mass of an element is the average mass of the atoms of that element as compared to 1/12 (one-twelfth) the mass of an atom of carbon-^ th 12 isotope (an element having different mass number but same atomic number). Based on carbon-12 standard, the mass of an atom of carbon is 12 units and l/2 of it comes to be^ th 1 unit. When we compare atomic masses of other elements with atomic mass of carbon- 12 atom, they are expressed as relative atomic masses of those elements. The unit for relative atomic masses is called atomic mass unit , with symbol ' amu '. One atomic mass unit is 1/12^ th the mass of one atom of carbon-12th. When this atomic mass unit is expressed in grams , it is:
For example:
1.2.4 How to write a Chemical Formula
Compounds are represented by chemical formulae as elements are represented by symbols. Chemical formulae of compounds are written keeping the following steps in consideration.
i. Symbols of two elements are written side by side, in the order of positive ion first and negative ion later.
ii. The valency of each ion is written on the right top corner of its symbol, e.g. Na , Ca , CI and O2.^ +^ 2+^ ^
Test yourself 1. It represents the name of the substance e.g. It tells the H O (water).^2 name of the elements as present in the compound.
It indicates the mass of the compound in amus or grams.
It is in fact one molecule or formula unit of the compound.
It also represents one mole of the molecules in the balanced chemical equation.
Significance of chemical formula
i) How many amu 1 g of a substance has? ii) Is atomic mass unit a SI unit of an atomic mass? iii) What is the relationship between atomic number and atomic mass? iv) Define relative atomic mass. v) Why atomic mass of an atom is defined as relative atomic mass?
Molecules are formed by the combination of atoms. These molecules are represented by molecular formulae that show actual number of atoms of each element present in a molecule of that compound. Molecular formula is derived from empirical formula by the following relationship:
Molecular formula = (Empirical formula)n Where n is 1,2,3 and so on. For example, molecular formula of benzene is C H which is derived from the 6 6 empirical formula CH where the value of n is 6.
The molecular formula of a compound may be same or a multiple of the empirical formula. A few compounds having different empirical and molecular formulae are shown in Table 1.6.
Table 1.6 Some Compounds with their Empirical and Molecular Formulae
Some compounds may have same empirical and molecular formula e.g. water (H 0), 2 hydrochloric acid (HC1), etc.
1.2.5 Molecular Mass and Formula Mass
The sum of atomic masses of all the atoms present in one molecule of a molecular substance, is its molecular mass. For example, molecular mass of chlorine (Cl ) is 71.0 2 amu , of water (H O) is 18 2 amu and that of carbon oxide (CO ) is 44 2 amu.
Example 1.
Calculate the molecular mass of Nitric acid, HNO. 3
Solution Atomic mass of H = 1 amu Atomic mass of N = 14 amu Atomic mass of O = 16 amu Molecular formula = HNO 3 Molecular mass = 1 (At. mass of H) + 1 (At. mass of N) + 3 (At. mass of O) = 1 + 14 + 3(16) = 1 + 14 + 48 = 63 amu Some ionic compounds that form three dimensional solid crystals, are represented by their formula units. Formula mass in such cases is the sum of atomic masses of all the atoms present in one formula unit of a substance. For example, formula mass of sodium chloride is 58.5 amu and that of CaCO is 100 3 amu.
Example 1.
Calculate the formula mass of Potassium sulphate K SO 2 4
Solution
Atomic mass of K = 39 amu Atomic mass of S = 32 amu Atomic mass of O = 16 amu Formula unit = K SO 2 4 Formula mass of K SO 2 4 = 2(39) + 1(32) + 4(16) = 78 + 32 + 64 = 174 amu
1.3 CHEMICAL SPECIES
1.3.1 Ions (Cations and Anions), Molecular Ions and Free Radicals
Ion is an atom or group of atoms having a charge on it. The charge may be positive or negative. There are two types of ions i.e. cations and anions. An atom or group of atoms having positive charge on it is called cation. The cations are formed when atoms lose electrons from their outermost shells. For example, Na+, K+ are cations. The following equations show the formation of cations from atoms.
An atom or a group of atoms that has a negative charge on it, is called anion. Anion is formed by the gain or addition of electrons to an atom. For example, Cl^ and O.^2 Following examples show the formation of an anion by addition of electrons to an atom.
Table 1.7 Difference between Atoms and Ions
Test yourself 1.
i. What is the relationship between empirical formula and formula unit? ii. How can you differentiate between molecular formula and empirical formula? iii. Identify the following formulae as formulas or unit molecular formulae: H O , CH , C H O , C H O , BaCO , KBr 2 2 4 6 12 6 12 22 1 3 iv. What is empirical formula of acetic acid (CH COOH)? 3 Find out its molecular mass. v. Calculate the formula masses of: Na S0 , ZnSO and CuCO. 2 4 4 3
Do you know?
Table 1.9 Difference between Ions and Free Radicals
1.3.2 Types of Molecules
A molecule is formed by the chemical combinations of atoms. It is the smallest unit of a substance. It shows all the properties of the substance and can exist independently. There are different types of molecules depending upon the number and types of atoms combining. A few types are discussed here.
A molecule consisting of only one atom is called monoatomic molecule. For example, the inert gases helium, neon and argon all exist independently in atomic form and they are called monoatomic molecules.
If a molecule consists of two atoms, it is called diatomic molecule. For example: hydrogen (H ), oxygen (O ), chlorine (Cl ) and hydrogen chloride (HCl). 2 2 2
If it consists of three atoms, it is called triatomic molecule. For example :H O and 2 CO. If a molecule consists of many atoms, it is called polyatomic. For example: methane 2 (CH ), sulphuric acid (H2SO ) and glucose (C H O ). 4 4 6 12 6
A Molecule containing same type of atoms, is called homoatomic molecule. For example: hydrogen (H ), ozone (O ), sulphur (S ) and phosphorus (P ) are the examples 2 3 8 4 of molecules formed by the same type of atoms. When a molecule consists of different kinds of atoms, it is called heteroatomic molecule. For example: CO , H O and NH. 2 2 3
Most of the universe exists in the form of plasma, the fourth state of matter. Both the cationic and anionic molecular ions are present in it.
sunlight
sunlight
mass of a substance. Avogadro's Number is a collection of 6.02 1023 particles. It is represented by symbol ' N '. Hence, the 6.02 1023 number of atoms, molecules or A formula units is called Avogadro's number that is equivalent to one 'mole' of respective substance. In simple words, 6.02 1023 particles are equal to one mole as twelve eggs are equal to one dozen. To understand the relationship between the Avogadro's number and the mole of a substance let us consider a few examples.
i. 6.02 1023 atoms of carbon are equivalent to one mole of carbon.
ii. 6.02 1023 molecules of H O are equivalent to one mole of^2 water.
iii. 6.02 1023 formula units of NaCl are equivalent to one mole of sodium chloride. Thus, 6.02 1023 atoms of elements or 6.02 1023 molecules of molecular substance or 6.02 1023 formula units of ionic compounds are equivalent to 1 mole. For further explanation about number of atoms in molecular compounds or number of ions in ionic compounds let us discuss two examples: i. One molecule of water is made up of 2 atoms of hydrogen and 1 atom of oxygen, hence 2 6.02 1023 atoms of hydrogen and 6.02 1023 atoms of oxygen constitute one mole of water. ii. One formula unit of sodium chloride consists of one sodium ion and one chloride ion. So there are 6.02 1023 number of Na ions and 6.02 1023 CI ions in one mole of sodium chloride. Thus, the total number of ions in 1 mole of NaCl is 12.04l0^23 or 1.204 10.^24 1.5.2 Mole (Chemist secret unit)
A mole is defined as the amount(mass) of a substance that contains 6.02 l0^23 number of particles (atoms, molecules or formula units). It establishes a link between mass of a substance and number of particles as shown in summary of molar calculations. It is abbreviated as 'mol'. You know that a substance may be an element or compound (molecular or ionic). Mass of a substance is either one of the following: atomic mass, molecular mass or formula mass. These masses are expressed in atomic mass units (amu). But when these masses are expressed in grams, they are called as molar masses. Scientists have agreed that Avogadro's number of particles are present in one molar mass of a substance. Thus, quantitative definition of mole is the atomic mass, molecular mass or formula mass of a substance expressed in grams is called mole.
Amaedo Avogadro (1776-1856) was an Italian scholar. He is famous for molecular t h e o r y c o m m o n l y known as Avogadro's law. In tribute to him, the number of particles (atoms, molecules, ions) in mole of 23 a substance 6.02 10 is k n o w n a s t h e Avogadro's constant.
of the
Example 1. Calculate the gram molecule (number of moles) in 40 g of H PO. 3 4
Solution
Therefore, 40 grams will contain 0.408 gram molecule (mol) of H PO. 3 4
1.6 CHEMICAL CALCULATIONS
In chemical calculations, we calculate number of moles and number of particles of a given mass of a substance or vice versa. These calculations are based upon mole concept. Let us have a few examples of these calculations. Calculating the number of moles and number of particles from known mass of a substance. First calculate the number of moles from given mass by using equation
Then calculate number of particles from the calculated number of moles with the help of following equation:
1.6.1 Mole-Mass Calculations
In these calculations, we calculate the number of moles of a substance from the known mass of the substance with the help of following equation:
When we rearrange the equation to calculate mass of a substance from the number of moles of a substance we get,
Test yourself 1.
i. Which term is used to represent the mass of 1 mole of molecules of a substance? ii. How many atoms are present in one gram atomic mass of a substance? iii. Explain the relationship between mass and mole of a substance. iv. Find out the mass of 3 moles of oxygen atoms. v. How many molecules of water are present in half mole of water?
Example 1.
You have a piece of coal (carbon) weighing 9.0 gram. Calculate the number of moles of coal in the given mass.
Solution The mass is converted to the number of moles by the equation:
So, 9.0 g of coal is equivalent to 0.75 mol. 1.6.2 Mole-Particle Calculations
In these calculations, we can calculate the number of moles of a substance from the given number of particles. (These particles are the atoms, molecules or formula units).
On rearranging above equation we get,
Summary of Molar Calculations:
Example 1.
Calculate the number of moles, number of molecules and number of atoms present in 6 grams of water.
Mass of Substances Number of Particles
mole NA
known mass molar mass
Mole
mole molar mass
number of particles NA
Never calculate the number of particles from mass of the substance or vice versa. Always make calculations through moles. For calculations of the number of atoms in molecular compounds and the number of ions in ionic compounds; first calculate the number of molecules or formula units and then calculate the number of atoms or ions.
Remember
