Basic Concepts of Chemistry

• Chemistry
• Matter
• Atoms and Molecules
• Physical Quantities and their Measurement Units
• Dimensional Analysis
• Scientific Notation
• Precision and Accuracy
• Laws of Chemical Combinations
• Dalton's Atomic Theory
• Mole Concept
• Atomic Mass
• Molecular Mass
• Equivalent Mass
• Stoichiometry
• Percent Yield
• Empirical and Molecular Formulae

Chemistry

It is the branch of science which deals with the composition, structure, and properties of matter.

Antoine Laurent Lavoisier is called the father of chemistry.

Branches of Chemistry

Inorganic chemistry is concerned with the study of elements (other than carbon) and their compounds.

Organic chemistry is the branch of chemistry which is concerned with organic compounds or substances produced by living organisms.

Physical chemistry is concerned with the explanation of fundamental principles.

Analytical chemistry is the branch of chemistry which is concerned with qualitative and quantitative analysis of chemical substances.

In addition to these, biochemistry, war chemistry, nuclear chemistry, forensic chemistry, earth chemistry, etc., are other branches of chemistry.

Matter

Anything which occupies some space and has some mass is called matter. It is made up of small particles which have space between them. The matter particles attract each other and are in a state of continuous motion.

Classification of Matter

Pure Substances

They have characteristics different from the mixtures. They have fixed composition, whereas mixtures may contain the components in any ratio and their composition is variable.

Elements

It is the simplest form of pure substance, which can neither be decomposed nor be built from simpler substances by ordinary physical and chemical methods. It contains only one kind of atom. The number of elements known till date is 118.

An element can be a metal, a non-metal, or a metalloid.

Hydrogen is the most abundant element in the universe.

Oxygen (46.6%), a non-metal, is the most abundant element in the earth's crust.

Aluminium (Al) is the most abundant metal in the earth's crust.

Compounds

It is also the form of matter which can be formed by combining two or more elements in a definite ratio by mass. It can be decomposed into its constituent elements by suitable chemical methods, e.g., water (H₂O) is made of hydrogen and oxygen in the ratio 1:8 by mass.

Types of Compounds

• Inorganic Compounds: Previously, it was believed that these compounds are derived from non-living sources, like rocks and minerals. But they are, in fact, the compounds of all the elements except hydrides of carbon (hydrocarbons) and their derivatives.
• Organic Compounds: According to earlier scientists, these compounds are derived from living sources like plants and animals or remain buried under the earth (e.g., petroleum). According to modern concepts, these are the hydrides of carbon and their derivatives.

Mixtures

These are made up of two or more pure substances. They can possess variable composition and can be separated into their components by some physical methods.

Mixtures may be homogeneous (when composition is uniform throughout) or heterogeneous (when composition is not uniform throughout).

Mixture Separation Methods

Common methods for the separation of mixtures are:

• Filtration: The process of separating solids suspended in liquids by pouring the mixture into a filter funnel. The solid particles are held on the filter while the liquid passes through.
• Distillation: The process of heating a liquid to form vapours and then cooling the vapours to get back the liquid. This method is used to separate volatile substances.
• Sublimation: The process of conversion of a solid directly into vapours on heating. Substances showing this property are called sublimates (e.g., iodine, naphthalene, camphor). This method is used to separate a sublimate from non-sublimate substances.
• Crystallisation: It is a process of seperating solids having different solubilities in a perticular solvent.
• Magnetic Seperation: This process is based upon the fact that a magnet attracts the magnetic components of a mixture of magnetic and non magnetic substances. The non magnetic substances remains unaffected. Thus it can be used to seperate magnetic components from non-magnetic components.
• Atmolysis: This method is based upon the rates of diffusion of gases and used for their seperation from a gaseous mixture.

Atoms and Molecules

Atom: It is the smallest particle of an element which can take part in a chemical reaction. It may or may not be capable of independent existance.

Molecule: It is the particle of matter that has independent existence. It may be a homoatomic, e.g. H2, Cl2 etc, or hetroatomic, e.g. HCL, NH3, CH4, etc

Physical Quantities and their Measurement Units

Physical Quantities:is a physical property of a material that can be quantified by measurement and their measurement does not involve any chemical reactions. such as mass, length, and time have standard units of measurement.

To express the measurement of any physical qunatity to things are considered: (i) It's Unit, (ii) The Numerical Value.

Unit

It is defined as "some fixed standard against which the comparison of a physical quantity can be done during measurement."

Units are of two types:

(i) Basic units

(ii) Derived units

(ⅰ) The basic or fundamental units are:

length (m), mass (kg) time (s), electric current (A), thermodynamic temperature (K amount of substance (mol) and luminous intensity (Cd).

(ii)Derived units

These are basically derived from the fundamental unit e.g. unit of density is derived from units of mass and volume

Different systems used for describing measurements of various physical quantities are:

CGS system:It is based on centimetre, gram and second as the units of length, mass and time respectively.

FPS system: A British system which used foot (ft), pound (lb) and second (s) as the fundamental units of length, mass and time respectively.

MKS system: It is the system which uses metre (m), kilogram (kg) and second (s) respectively for length, mass and time; ampere (A) was added later on for electric current.

SI system: In (1960) International system of units or SI units contains following seven basic and two supplementary units:

Basic Physical Quantities and Their Corresponding SI Units

Supplementary units: It includes plane angle in radian and solid angle in steradian.

Prefixes

The SI units of some physical quantities are either too small or too large. To change the order of magnitude, these are expressed by using prefixes before the name of base units. The various prefixes are listed as:

Some Physical Quantities

(i) Mass: It is the amount of matter present in a substance. I butains constant for a substance at all the places. Its unit is kg laboratories usually gram is used.

(ii)Weight: It is the from place to plac force exertedge in gravity. Its unit is Newton 1 by gravity on an object. It varies from place to place due to change in gravity. Its unit is Newton(N).

(iii)Temperature:There are three common scale to measure temperature °C (degree celsius), °F (degree fahrenheit) and K (kelvin). K is the SI unit. The temperature on two scales (°C and °F) are related to each other by the following relationship:

The Fahrenheit & Celsius Scale relates as fallows ^ F= 9/5 (^ C)+32

The kelvin scale & celsius scale relates as follows: K=^ C + 273.15

(iv)Volume: The space matter (usually by liquid or a gas) is called its volume. Its unit is m ^ 3

(v)Density: It is defined as the amount or mass per unit volume and has units kg * m ^ - 3 or gcm-3.

Dimensional Analysis

Often while calculating, there is a need to convert units from one system to other. The method use to accomplish this is called factor label mehtod of unit factor method or Dimensional analysis.

In this,

Information sought = Information Given X Conversion factor

Important Conversion Factors

Scientific Notation

In such a notation all measurements(how so ever large or small) are expressed as a number betweeen 1.000 and 9.999 multiplied or divided by 10.

In simple words we can say that the Scientific notation is a way to express very large or very small numbers in a concise format.

Precision and Accuracy

Precision refers to the closeness of the set of values obtained from identical measurements of a quantity. Precision is simply a measure of reproducibility of an experiment.

Precision = individual value - arithmetic mean value

Accuracy is a measure of the difference between the experimental value or the mean value of a set of measurements and the true value.

Accuracy = mean value - true value

In physical measurements, accurate results are generally precise but precise results need not be accurate.

Laws of Chemical Combinations

These laws describe how elements combine to form compounds in definite proportions.

Law of conservation of mass (Lavoisier, 1789)

This law states that during any physical or chemical change, the total mass of the products is equal to the total mass of reactants. It does not hold good for nuclear reactions.

Law of definite proportions (Proust, 1799)

According to this law, a chemical compound obtained by different sources always contains same percentage of each constituent element.

Law of multiple proportions (Dalton, 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, e.g. in NH3 and N2H4, fixed mass of nitrogen requires hydrogen in the ratio 3: 2.

Law of reciprocal proportions (Richter, 1792)

According to this law, when two elements (say A and B) combine separately with the same weight of a third element (say C), the ratio in which they do so is the same or simple multiple of the ratio in which they (A and B) combine with each other. Law of definite proportions, law of multiple proportions and law of reciprocal proportions do not hold good when same compound is obtained by using different isotopes of the same element, e.g. H₂O and D2O.

Gay Lussac's law of gaseous volumes (In 1808)

It states that under similar conditions of temperature and pressure, whenever gases react together, the volumes of the reacting gases as well as products (if gases) bear a simple whole number ratio.

Avogadro's hypothesis

It states that equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules.

Dalton's Atomic Theory (1803)

This theory was based on laws of chemical combinations. It's basic postulates are:

1. All substances are made up of tiny, indivisible particles, called atoms.

2. In each element, the atoms are all alike and have the same mass. The atoms of different elements differ in mass.

3. Atoms can neither be created nor destroyed during any physica or chemical change.

4. Compounds or molecules result from combination of atoms in some simple numerical ratio.

Limitations

(i) It failed to explain how atoms combine to form molecules.

(ii) It does not explain the difference in masses, sizes and valencie of the atoms of different elements.

Mole Concept

Term mole was suggested by Ostwald (Latin word mole heap) A mole is defined as the amount of substance which contains same number of elementary particles (atoms, molecules or ions) as the number of atoms present in 12 g of carbon (C-12). 1mol = 6.023 * 10 ^ 23 atoms = one gram-atom gram atomic mass [ mol = 6.023 * 10 ^ 23 molecules gram molecular masa In gaseous state at STP GT = 273K p=1. atm ) Gram molecular mass 1 mol = 22.4 L = 6.022 * 10 ^ 23 molecules

Standard number 6.023 * 10 ^ 23 is called Avogadro number in honour of Avogadro (he did not give this number) and is denoted by N_{A} The volume occupied by one mole molecules of a gaseous substance is called molar volume or gram molecular volume.

Number of molecules = number of moles * N_{A}/ g-molar mass

Number of molecules in 1g compound N_{A} Number of molecules in 1c * m ^ 3 (1 mL) of an ideal gas at STP is called Loschmidt number (2.69 * 10 ^ 19)

[one amu or u (unified mass) is equal to exactly the 1/12 * th of the mass of ^ 12 C atom, i.e. 1 amu or u= 1/12 * mass of one carbon (C ^ 12) atom 1 amu = 1 NA 1 Avogram = 1 Aston =1 Dalton = 1.66x 10-24 g

One mole of electrons weighs 0.55 mg (5.5x 10 g).

Atomic Mass

Atomic mass is the average atomic mass of an atom. It indicates that how many times an atom of that elementis heavier as compared with 1/12th part of mass of one atom of carbon -12.

The word average has been used in the above definition and is significant because elements occur in nature as mixture of seve isotopes. So, atomic mass can be computed as

• Average atomic mass = RA (1) x at. mass (1) + RA (2) x at. mas/RA(1) + RA(2)

Here, RA is relative abundance of different isotopes.

• In case of volatile chlorides, the atomic weight is calculated as

At. wt. = Eq. wt. x valency

And

valency = 2 x vapour density of chloride/eq. wt. of metal + 35.5

• According to Dulong and Petit's rule, Atomic weight x specific heat = 6.4

Gram Atomic Mass (GAM)

Atomic mass of an element expressed in gram is called its gram at mass or gram-atom or mole-atom.

Molecular Mass

Molecular mass is the sum of the atomic masses of all atoms in a molecule.

Equivalent Mass

Equivalent mass is the mass of a substance that reacts with a fixed amount of another substance.

Stoichiometry

Stoichiometry is the branch of chemistry that deals with the quantitative relationships of reactants and products in a chemical reaction. It is based on the law of conservation of mass, which states that matter is neither created nor destroyed in a chemical reaction.

Importance of Stoichiometry

Predicting Product Yields – Helps determine the amount of product formed from given reactants.

Determining Reactant Quantities – Identifies the exact amount of reactants needed for a reaction.

Optimizing Industrial Reactions – Essential for maximizing efficiency in chemical industries.

Balancing Chemical Equations – Ensures mass and charge conservation.

Types of Stoichiometry

Reaction Stoichiometry

Deals with the quantitative relationship between reactants and products in a balanced chemical equation.

Example: 2H2+O2→2H2O 2H2​+O2​→2H2​O This equation shows that 2 moles of hydrogen react with 1 mole of oxygen to form 2 moles of water.

Composition Stoichiometry

Focuses on the ratio of elements within a compound.

Example: In water (H₂O), the ratio of hydrogen to oxygen atoms is 2:1.

Mole Stoichiometry

Determines the relationship between moles of reactants and products.

Example: C+O2→CO2 C+O2​→CO2​ In this reaction 1 mole of carbon reacts with 1 mole of oxygen to form 1 mole of CO₂.

Key Concepts in Stoichiometry

Mole Concept

1 mole of a substance contains Avogadro's number of particles (6.022 × 10²³). The molar mass (grams per mole) is used to relate mass and number of moles.

Balancing Chemical Equations

Ensures the number of atoms of each element is the same on both sides of the equation.

Example: Fe+O2→Fe2O3 Fe+O2​→Fe2​O3​

Balanced form: 4Fe+3O2→2Fe2O3 4Fe+3O2​→2Fe2​O3​

Limiting and Excess Reactants

Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

Excess Reactant: The reactant left over after the reaction is complete.

Example: If 5 moles of H₂ react with 2 moles of O₂, oxygen is the limiting reactant.

Theoretical, Actual, and Percentage Yield

Theoretical Yield: Maximum product expected from a reaction.

Actual Yield: Product actually obtained.

Percentage Yield: Efficiency of the reaction.

Percentage Yield=(Actual YieldTheoretical Yield)×100 Percentage Yield=(Theoretical YieldActual Yield​)×100

Percent Yield

Percent yield is a measure of the efficiency of a chemical reaction. It compares the actual yield of a product obtained from a reaction to the theoretical yield, which is the maximum amount of product that could be produced based on stoichiometric calculations.

Mathematically, percent yield is given by:

Percent Yield=(Actual Yield/Theoretical Yield)×100%

Where:

Actual yield is the measured amount of product obtained from the experiment. Theoretical yield is the calculated amount of product expected based on the balanced chemical equation.

Significance of Percent Yield

Evaluates Reaction Efficiency: A high percent yield indicates an efficient reaction with minimal losses. Identifies Practical Limitations: Percent yield helps chemists understand practical inefficiencies such as side reactions, incomplete reactions, or product loss during purification.

Essential for Industrial Chemistry: In pharmaceutical, food, and chemical industries, maximizing percent yield reduces costs and improves production efficiency.

Empirical and Molecular Formulae

The empirical formula of a compound represents the simplest whole-number ratio of atoms of each element present in the compound. It does not show the actual number of atoms but gives the basic atomic ratio.

Key Features of Empirical Formula

Simplest Ratio – Expresses elements in the smallest whole numbers.

Does Not Indicate Molecular Structure – Only provides composition, not how atoms are arranged.

Different Compounds Can Have the Same Empirical Formula – Example: Glucose (C₆H₁₂O₆) and Acetic Acid (C₂H₄O₂) both have the empirical formula CH₂O.

How to Determine the Empirical Formula

To find the empirical formula, follow these steps:

Obtain the Mass Percentages of Elements (from experiment or problem statement).

Convert Percentages to Masses (assuming 100 g of the compound).

Convert Masses to Moles (by dividing by atomic mass).

Find the Simplest Whole-Number Ratio (by dividing by the smallest mole value).

Multiply if Necessary (to eliminate fractions).

Example Calculation

Problem:

A compound contains 40% carbon (C), 6.7% hydrogen (H), and 53.3% oxygen (O). Determine its empirical formula.

Solution:

Assume 100 g of the compound, so the mass of each element:

C = 40 g, H = 6.7 g, O = 53.3 g

Convert mass to moles (using atomic masses: C = 12, H = 1, O = 16)

Moles of C = 40 ÷ 12 = 3.33

Moles of H = 6.7 ÷ 1 = 6.7

Moles of O = 53.3 ÷ 16 = 3.33

Divide by the smallest value (3.33)

C : 3.33 ÷ 3.33 = 1

H : 6.7 ÷ 3.33 = 2

O : 3.33 ÷ 3.33 = 1

Write the empirical formula

Empirical Formula = CH₂O

Difference Between Empirical and Molecular Formula

Property Empirical Formula Molecular Formula Definition Simplest ratio of atoms Actual number of atoms Example CH₂O (for glucose) C₆H₁₂O₆ (glucose)

Relationship

Molecular Formula = (Empirical Formula) × n Can be a multiple of the empirical formula

Example:

Benzene (C₆H₆) has an empirical formula of CH.

Glucose (C₆H₁₂O₆) has an empirical formula of CH₂O.

Applications of Empirical Formula

Used in Chemical Analysis – Helps determine unknown compounds.

Important in Stoichiometry – Essential for reaction calculations.

Determining Molecular Formula – When combined with molar mass data.

Conclusion

The empirical formula is a fundamental concept in chemistry that simplifies the representation of compounds. It helps chemists analyze compositions, compare substances, and determine molecular formulas.

Empirical formula shows the simplest ratio of atoms, while molecular formula gives the exact number of atoms in a molecule.