Some Basic concepts of chemistry-Notes

Matter is anything that possesses mass and occupies space. The nature of matter refers to its fundamental character—how matter is built from extremely small particles, how these particles behave, and how their interactions give rise to the observable properties of substances. Although matter appears continuous to our senses, careful study reveals that it is discrete, made up of tiny entities (atoms and molecules) that are in constant motion.

This particulate view of matter explains why substances can mix, why gases expand, why solids maintain shape, and why chemical changes occur. The modern understanding of matter evolved from early atomic ideas, notably formalized by scientists such as John Dalton, whose work emphasized that all material substances are composed of indivisible particles called atoms.

A state of matter refers to a distinct physical form in which matter exists under given conditions of temperature and pressure. Although all matter is made up of particles, the way these particles are arranged, how strongly they attract one another, and how freely they move decides the observable form of a substance.

In everyday life, matter is commonly found in three physical states: solid, liquid, and gas. These states differ mainly in shape, volume, compressibility, and particle motion.

Interconversion of States of Matter

Matter can change from one state to another by altering temperature or pressure. These changes are physical and reversible. \[\boxed{\bbox[blue,5pt]{ \text{Solid} \;\xRightleftharpoons[\text{cool}]{\text{heat}}\; \text{Liquid} \;\xRightleftharpoons[\text{cool}]{\text{heat}}\; \text{Gas}}} \]

Every substance shows certain observable and measurable features. These features are called the properties of matter. A property is any characteristic that helps us describe, identify, or compare substances. Some properties can be observed directly, while others require careful measurement using instruments.

In chemistry, the study of matter begins by recognizing these properties and learning how to measure them accurately. This forms the foundation for understanding chemical behavior and performing calculations.

Classification of Properties

Properties of matter are broadly divided into physical properties and chemical properties.

1. Physical Properties
   

   Physical properties are those characteristics that can be observed or measured without changing the chemical identity of the substance.

   Common physical properties include:

   • Mass
   • Volume
   • Density
   • Colour
   • Melting point
   • Boiling point
   • Solubility
   • Electrical conductivity

   For example, measuring the mass of copper or observing the colour of sulphur does not alter their composition.
   These properties are especially useful for identification and comparison of substances.

2. Chemical Properties
   

   Chemical properties describe how a substance behaves when it undergoes a chemical change.

   Examples include:
   • Combustibility
   • Rusting
   • Reactivity with acids or bases
   • Ability to undergo oxidation or reduction

   These properties become evident only when new substances are formed.

Measurement of Properties of Matter

To study matter scientifically, observations must be expressed in numerical form. This requires measurement, which involves comparing an unknown quantity with a known standard.

• A numerical value
• A unit

For example: 250 g, 2.5 L, 300 K

Physical Quantities Commonly Used in Chemistry

Some important quantities measured in chemistry are:

• Mass
• Length
• Time
• Temperature
• Amount of substance

To maintain uniformity worldwide, scientists follow the International System of Units, which provides standard units for all physical quantities.

Fundamental Units Used in Chemistry

Definitions of SI Base Units

Prefixes used in the SI System

Mass

Mass is the amount of matter present in a body. It represents how much material an object contains and does not depend on where the object is placed.

Mass is a scalar quantity—it has magnitude only. Whether an object is on Earth, the Moon, or in space, its mass remains the same because the quantity of matter does not change.

In laboratories, mass is usually measured with a beam balance or electronic balance, and its SI unit is the kilogram (kg).

Key idea:

Mass reflects quantity of matter and remains constant everywhere.

Weight

Weight is the force with which a body is attracted towards the centre of the Earth (or any celestial body). It depends on both the mass of the object and the local value of acceleration due to gravity.

Mathematically:

where \(g\) is the acceleration due to gravity.

Weight is a vector quantity because it has both magnitude and direction (always towards the centre of gravity). It is measured using a spring balance, and its SI unit is the newton (N).

Since gravity varies from place to place, weight changes with location. For example, a person weighs less on the Moon than on Earth, even though their mass remains unchanged.

Important Differences Between Mass and Weight

Important Aspects to Remember

• Mass is an intrinsic property of matter; weight is an external effect due to gravity.
• Chemical calculations always use mass, not weight.
• In space, astronauts become weightless, but their mass remains the same.
• Weight can change with altitude or planetary location, while mass is invariant.
• Understanding this distinction is essential for measurements, stoichiometry, and laboratory work.

Volume refers to the space occupied by matter. Liquids are commonly measured using graduated cylinders or burettes, while volumes of gases are measured using gas syringes or calculated from container dimensions.

For liquids:

Density is a physical property of matter that expresses how much mass is contained in a given volume of a substance. It helps us understand how closely matter is packed within a material.

In simple terms, density tells us how heavy a substance is for its size. Two substances may have the same volume but different masses; the one with greater mass has higher density.

Mathematically, Density is defined as the ratio of mass to volume:

Units of Density

Since mass is measured in kilograms (kg) and volume in cubic metres (m³), the SI unit of density is: \[\mathrm{kg\,m^{-3}}\]

Temperature is a physical quantity that indicates the degree of hotness or coldness of a body. In scientific terms, temperature is a measure of the average kinetic energy of the particles present in matter.

When a substance is heated, its particles move faster, and the temperature rises. When it is cooled, particle motion slows down, and the temperature falls. Thus, temperature reflects how energetic the particles of a system are.

Microscopic Interpretation

At the particle level:

• Higher temperature \(\Rightarrow\) faster particle motion \(\Rightarrow\) greater kinetic energy
• Lower temperature \(\Rightarrow\) slower particle motion \(\Rightarrow\) lesser kinetic energy

Temperature does not depend on the amount of substance present. A cup of hot tea and a bucket of hot water may have the same temperature even though their masses are very different.

Measurement of Temperature

Temperature is measured using a thermometer. In chemical laboratories, temperature is expressed on the absolute scale, whose SI unit is the kelvin (K).

Other commonly used scales are:

• Celsius (°C)
• Fahrenheit (°F)

The relationship between Celsius and Kelvin scales is: \[\boxed{\bbox[blue,5pt]{T(K)=t(^\circ C)+273}}\]

Absolute Zero

Absolute zero (0 K) is the lowest possible temperature at which molecular motion is minimum. At this temperature, particles possess the least kinetic energy achievable.

Though absolute zero cannot be attained practically, it provides a theoretical reference point for thermodynamic measurements.

Temperature vs Heat

It is important not to confuse temperature with heat.

• Temperature indicates the intensity of thermal energy.
• Heat refers to the total thermal energy transferred from one body to another due to temperature difference.

A small object at high temperature may contain less heat than a large object at a lower temperature.

In science, no measurement is perfectly exact. Every observed value carries a small doubt about its correctness. This unavoidable doubt is called uncertainty in measurement.

Uncertainty tells us the possible range within which the true value of a quantity is expected to lie. It arises because measuring instruments have limited precision and because human observation cannot be flawless.

For example, when the mass of a substance is recorded as 25.4 g, the last digit is uncertain—it is an estimate based on the scale of the balance.

Reasons for uncertainty in measurement

Uncertainty mainly originates from three sources:

1. Instrumental Limitations:
   

   Every instrument has a least count (smallest readable division). Values smaller than this cannot be measured exactly.

   Example:
   If a balance reads up to 0.01 g, then measurements are uncertain in the second decimal place.

2. Human Observation:
   Errors may occur while reading scales, especially due to:
   • Parallax (eye not aligned with scale)
   • Improper estimation between markings
   • Reaction time in timing experiments
3. Environmental Factors:
   Temperature, air currents, humidity, and vibrations can affect measurements, particularly in sensitive experiments.

In scientific measurements, every recorded value contains certain reliable digits and one doubtful digit. These meaningful digits are called significant figures.

Simply stated, significant figures are the digits in a measured quantity that convey information about its precision. The last digit is always uncertain because it is estimated from the instrument scale.

For example, in a mass recorded as 2.45 g, all three digits are significant, but the final “5” carries uncertainty.

Significant figures therefore tell us how accurately a quantity has been measured.

Rules for Counting Significant Figures

1. All non-zero digits are significant
   

   Example:
   25.6 → three significant figures

2. Zeros between non-zero digits are significant
   

   Example:
   1002 → four significant figures

3. Leading zeros are not significant
   

   Example:
   0.0045 → two significant figures
   (zeros only locate the decimal point)

4. Trailing zeros after a decimal point are significant
   

   Example:
   2.300 → four significant figures

5. Trailing zeros in whole numbers without decimal point are ambiguous
   Example:
   1500 may have two, three, or four significant figures unless written in scientific notation.
6. Scientific Notation and Significant Figures
   

   To avoid confusion, large or small numbers are written in scientific notation: \[N=a\times 10^n\] where \(a\) contains all significant figures.
   
   Example:
   4500 written as \(\mathrm{4.50\times 10^3}\) clearly shows three significant figures.

Precision

precision refers to the closeness of various measurements for the same quantity.

Accuracy

Accuracy is the agreement of a particular value to the true value of the result.

The law of conservation of mass states that mass is neither created nor destroyed during a chemical reaction. In any chemical change, the total mass of substances present before the reaction is exactly equal to the total mass of substances formed after the reaction, provided the reaction occurs in a closed system.

This law establishes that chemical reactions involve only the rearrangement of particles, not their creation or annihilation.

The principle was firmly established through careful experimentation by Antoine Lavoisier, who emphasized accurate measurement as the foundation of chemistry.

The law of definite proportions states that a given chemical compound always contains the same elements combined in a fixed ratio by mass, irrespective of its source or method of preparation.

In other words, once a compound is formed, its composition becomes permanent. Whether the substance is prepared in a laboratory or obtained from nature, the relative masses of its constituent elements remain unchanged.

This law was established through systematic experiments by Joseph Proust, who demonstrated that compounds possess constant composition.

Significance of the Law

The law of definite proportions is fundamental to chemical science because:

• It confirms that compounds have fixed composition.
• It supports the atomic nature of matter.
• It forms the basis of chemical formulae.
• It enables quantitative calculations in reactions.
• It helps distinguish compounds from mixtures.

Unlike mixtures, where composition can vary, compounds obey this law strictly.

The law of multiple proportions states that when two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other bear a simple whole-number ratio to one another.

This law highlights that chemical combination is not arbitrary. Instead, elements unite in definite and orderly ways, producing compounds whose compositions follow simple numerical relationships.

The principle was proposed by John Dalton and provided strong support to the atomic view of matter.

Avogadro proposed that equal volumes of all gases at the same temperature and pressure should contain equal number of molecules.

Conceptual Basis of the Law

Gas particles are far apart and experience very weak intermolecular forces. Under the same temperature and pressure:

• Each gas molecule occupies nearly the same average space.
• Volume depends mainly on the number of molecules, not on their size or mass.

Therefore, doubling the number of gas molecules doubles the volume, and halving the number halves the volume.

Key Points to Remember

• Equal gas volumes at same temperature and pressure contain equal numbers of molecules.
• Volume is directly proportional to number of moles.
• Nature of gas does not affect particle count per unit volume.
• Law links microscopic molecules with macroscopic volume
• Essential for gas stoichiometry and mole calculations.

Dalton’s atomic theory states that all matter is composed of extremely small particles known as atoms, and that chemical reactions involve the rearrangement of these atoms in definite numerical ratios.

The theory explains why elements combine in fixed proportions and why mass is conserved during chemical change.

Main Postulates of Dalton’s Atomic Theory

Dalton proposed the following fundamental ideas:

1. All matter is made of atoms.
   Atoms are tiny, indivisible particles that form the basic building blocks of substances.
2. Atoms of the same element are identical.
   They have the same mass and properties, while atoms of different elements differ in mass and nature.
3. Atoms cannot be created, destroyed, or subdivided in chemical reactions.
   Chemical changes only rearrange atoms.
4. Compounds are formed when atoms of different elements combine in simple whole-number ratios.
5. In a given compound, the relative number and kinds of atoms remain fixed.

These statements provided a clear explanation for observed experimental laws.

Atomic mass is the average mass of an atom of an element expressed relative to a standard reference. Instead of measuring the actual mass of a single atom (which is extremely small), atomic masses are compared on a relative scale.

On this scale, the mass of one atom of carbon-12 is taken as exactly 12 units, and the masses of all other atoms are measured with respect to this standard.

Thus, atomic mass tells us how heavy an atom is compared to one-twelfth the mass of a carbon-12 atom.

Atomic Mass Unit (u)

The unit used to express atomic mass is the atomic mass unit (u).

Average Atomic Mass

Most elements exist naturally as a mixture of isotopes (atoms of the same element having different masses). Therefore, the atomic mass shown in tables is not the mass of a single atom but a weighted average of all naturally occurring isotopes.

Suppose an element has two isotopes:

• Isotope A with mass \(m_1\) and abundance \(x%\)
• Isotope B with mass \(m_2\) and abundance \(y%\)

This explains why atomic masses are often fractional values.

Key Points to Remember

• Atomic mass is based on carbon-12 standard.
• Unit used is atomic mass unit (u).
• Tabulated atomic mass is an average value.
• Fractional atomic masses arise due to isotopes.
• Atomic mass connects microscopic atoms with macroscopic measurements.

Molecular mass is the sum of the atomic masses of all the atoms present in one molecule of a substance. It tells us how heavy a molecule is on the atomic mass scale.

Since individual molecules are extremely small, their masses are not measured directly. Instead, molecular mass is calculated using the atomic masses of constituent elements.

In simple terms, molecular mass represents the relative mass of a molecule compared to one-twelfth the mass of a carbon-12 atom.

Not all substances exist as individual molecules. Many compounds, especially ionic solids, are made up of vast networks of oppositely charged ions arranged in a crystal lattice. Such substances do not have discrete molecules; instead, they are represented by the simplest whole-number ratio of ions, called a formula unit.

The formula mass of a substance is defined as the sum of the atomic masses of all atoms present in one formula unit of that compound.

In simple words, formula mass tells us how heavy one formula unit of an ionic compound is on the atomic mass scale.

For molecular substances (like \(\mathrm{H_2O}\) or \(\mathrm{CO_2}\)), we calculate molecular mass.

But for ionic compounds (like \(\mathrm{NaCl}\) or \(\mathrm{MgO}\)), molecules do not exist independently. Therefore, instead of molecular mass, we use formula mass.

Atoms and molecules are far too small to be counted individually. To deal with such enormous numbers in a practical way, chemistry uses a counting unit called the mole.

A mole is defined as the amount of substance that contains exactly \(6.022\times 10^{23}\) elementary entities (atoms, molecules, ions, or formula units). This fixed number is known as Avogadro’s constant, introduced following the ideas of Amedeo Avogadro.

Thus \[\boxed{\bbox[blue,5pt]{\mathrm{1\, mol=6.022\times 10^{23} \,particles}}}\]

Molar Mass

The molar mass of a substance is defined as the mass of one mole of that substance. It is expressed in grams per mole (g mol⁻¹).

Numerically, molar mass is equal to:

• atomic mass (for elements),
• molecular mass (for molecular compounds),

Relationship Between Mass, Moles, and Particles

• Mole–Mass Relation
  \[\scriptsize\boxed{\bbox[blue,5pt]{\text{Number of Moles}=\dfrac{\text{Given Mass}}{\text{Molar Mass}}}}\]
• Mole–Particle Relation
  \[\scriptsize\boxed{\bbox[blue,5pt]{\begin{aligned}&\text{Number of particles}=\\&\text{Number of moles}\times 6.022\times 10^{23}\end{aligned}}}\]

Percentage composition tells us how much of each element is present in a compound, expressed as a percentage by mass.

Every compound is formed when elements combine in fixed proportions by mass. If we know the molecular or formula mass of a compound and the atomic masses of its constituent elements, we can determine what fraction of the total mass comes from each element.

Because composition is constant for a given compound, these percentages remain the same regardless of how much of the substance is taken.

Formula for Percentage Composition

Empirical Formula

The empirical formula of a compound shows the simplest whole-number ratio of atoms of different elements present in that compound.

It does not give the actual number of atoms in a molecule—only their relative proportions.

For example:

Molecular formula of glucose is \(\mathrm{C_6H_{12}O_6}\)
Simplest ratio = \(\mathrm{CH_2O}\)
So, \(\mathrm{CH_2O}\) is the empirical formula of glucose.
In simple terms, the empirical formula tells us how elements are combined proportionally.

Molecular Formula

The molecular formula gives the actual number of atoms of each element in one molecule of a compound.
It is always a whole-number multiple of the empirical formula and represents the true composition of a molecule.
Example:
Empirical formula: \(\mathrm{CH_2O}\)
Molecular formula: \(\mathrm{C_6H_{12}O_6}\)
Here, every subscript in the molecular formula is six times that in the empirical formula.

Determination of Molecular Formula

To find molecular formula, we must know the molar mass of the compound.

Key Points to Remember

• Empirical formula = simplest ratio of atoms
• Molecular formula = actual number of atoms
• Molecular formula is a whole-number multiple of empirical formula
• Determination requires percentage composition and molar mass
• Central to quantitative chemistry

Stoichiometry is the branch of chemistry that deals with the quantitative relationship between reactants and products in a chemical reaction. It tells us how much of each substance is required or formed.

The word stoichiometry comes from Greek roots meaning measure of elements. In practical terms, stoichiometry connects chemical equations with actual laboratory quantities such as mass, volume, and number of moles.

Every stoichiometric calculation is based on a balanced chemical equation, which represents the fixed ratios in which substances react.

Conceptual Basis

Chemical reactions occur because atoms rearrange themselves in definite numbers. Since atoms combine in whole-number ratios, reactions also follow exact numerical relationships.

For example, consider a simple reaction: \[\mathrm{2H_2+O_2\rightarrow 2H_2O}\]

This equation tells us that:

• 2 molecules (or moles) of hydrogen react with
  
• 1 molecule (or mole) of oxygen to produce
  
• 2 molecules (or moles) of water.
  

These coefficients form the stoichiometric ratios of the reaction.
Thus, stoichiometry rests on three fundamental ideas:

• conservation of mass,
• fixed composition of compounds,
• mole concept.

Types of Stoichiometric Calculations

1. Mass–Mass Calculations:
   Used to find the mass of a product formed from a given mass of reactant (or vice versa).
2. Mole–Mole Calculations
   Directly use coefficients of the balanced equation to relate amounts of substances.
3. Volume Relationships (for gases)
   At the same temperature and pressure, gaseous reactants and products combine in simple volume ratios corresponding to mole ratios.

In most chemical reactions, reactants are not mixed in exactly the required stoichiometric ratio. As the reaction proceeds, one reactant gets completely used up first. This reactant is called the limiting reagent (or limiting reactant).

The limiting reagent decides the maximum amount of product that can be formed. Once it is exhausted, the reaction stops—even if other reactants are still present.

Any substance left behind after the reaction is known as the excess reagent.

In simple words:
The limiting reagent controls the yield of the reaction.

Most chemical reactions in the laboratory are carried out in solutions rather than between dry solids. In solution chemistry, reactants are present as dissolved particles (ions or molecules), and their amounts are expressed using concentration terms. To predict how much reactant will react or how much product will form, we must describe solutions quantitatively.

Four commonly used ways to express concentration are:

• Mass percent
• Mole fraction
• Molarity
• Molality

Each represents composition differently, but all serve the same purpose: linking the quantity of dissolved substance with the extent of reaction.

Mass Percent

Mass percent expresses the mass of solute present in 100 parts by mass of solution. \[\boxed{\bbox[blue,5pt]{\text{Mass Percent}=\dfrac{\text{Mass of Solute}}{\text{Mass of Solution}}\times 100}}\]

Role in Reactions

Mass percent is useful when solutions are prepared or supplied by weight. From this value, the actual mass—and hence moles—of reactant participating in a reaction can be calculated.

Mole Fraction (\(X\))

Mole fraction is the ratio of moles of one component to the total moles of all components in the solution. \[\boxed{\bbox[blue,5pt]{X_A=\dfrac{n_A}{n_A+n_B}}}\] where \(n_A\) and \(n_B\) are moles of solute and solvent.

Mole fraction has no unit.

Role in Reactions

Mole fraction directly represents the particle proportion in solution. It is especially useful in:

• gas–solution equilibria,
• vapour pressure studies,
• thermodynamic calculations.

In reaction chemistry, it helps describe composition at the molecular level.

Molarity (M)

Molarity is the number of moles of solute dissolved in one litre of solution. \[\boxed{\bbox[blue,5pt]{\text{Molarity}= \dfrac{\text{Moles of Solute}}{\text{Volume of Solution in litres}}}}\] unit: \(\mathrm{mol\,L^{-1}\; or\; M}\)

Role in Reactions

Molarity is the most widely used concentration unit in solution reactions because volumes of liquids are easy to measure.

In reactions such as titrations, molarity allows direct calculation of reacting moles: \[\text{Moles}=M\times V\] Thus, balanced equations combined with molarity give quick access to unknown concentrations or required volumes.

Molality (m)

Molality is the number of moles of solute present in one kilogram of solvent. \[\boxed{\bbox[blue,5pt]{\text{Molality}=\dfrac{\text{Moles of Solute}}{\text{Mass of Solvent (kg)}}}}\] Unit: mol kg⁻¹

Role in Reactions

Molality depends on mass, not volume. Since mass does not change with temperature, molality remains constant under temperature variation.

It is preferred in studies involving:

• temperature-dependent properties,
• colligative properties,
• precise thermodynamic measurements.