Unit 1: Some Basic Concepts of Chemistry

Comprehensive vocabulary flashcards covering the key terms, historical discoveries, scientific laws, SI units, and quantitative concepts from Unit 1: Some Basic Concepts of Chemistry.

Chemistry

The branch of science that studies the preparation, properties, structure, and reactions of material substances.

Philosopher's Stone (Paras)

A legendary substance sought by alchemists that was believed to convert all baser metals, such as iron and copper, into gold.

Elixir of Life

A mythical potion or substance sought by alchemists that was believed to grant immortality.

Rasayan Shastra

The name given to chemistry in ancient India, which also included practices known as Rastantra, Ras Kriya, or Rasvidya covering metallurgy, medicine, cosmetics, glass, and dyes.

Faience

A kind of glass produced by Harappans that was used in making ornaments.

Nagarjuna

A reputed Indian chemist, alchemist, and metallurgist whose work 'Rasratnakar' deals with the formulation of mercury compounds and extraction methods for gold, silver, tin, and copper.

Chakrapani

An ancient Indian scholar credited with discovering mercury sulphide and inventing soap using mustard oil and alkalies.

Varähmihir’s Brihat Samhita

A sixth-century CE encyclopaedia providing recipes for perfumes, cosmetics, hair dyes, and glutinous materials prepared from plant extracts and resins applied to roofs and walls.

Acharya Kanda

An Indian philosopher born in 600 BCE (originally known as Kashyap) and author of the 'Vaiseshika Sutras' who proposed that matter consists of small, indestructible, spherical units called 'Paramãnu'.

Paramãnu

The term formulated by Acharya Kanda referring to very small, indivisible, eternal, indestructible, and suprasensible building blocks of matter that are in motion in their original state.

Cisplatin and Taxol

Life-saving drugs isolated or synthesised through chemical methods that are effective in cancer therapy.

AZT (Azidothymidine)

A drug synthesised using chemical principles used for helping AIDS patients.

Matter

Anything that has mass and occupies space.

States of Matter

The three physical forms in which matter exists: solids (definite shape and volume), liquids (definite volume but no definite shape), and gases (neither definite volume nor definite shape).

Classification of Matter

The macroscopic division of matter into mixtures (homogeneous and heterogeneous) and pure substances (elements and compounds).

Pure Substance

A form of matter having a fixed composition in which all constituent particles are identical in chemical nature.

Mixture

A material composed of two or more pure substances present in any ratio with a variable composition whose components can be separated by physical methods.

Homogeneous Mixture

A mixture where the components completely mix with each other, resulting in a uniform composition and particle distribution throughout the bulk.

Heterogeneous Mixture

A mixture where the composition is not uniform throughout and different components are sometimes visibly distinguishable.

Element

A pure substance whose constituent particles consist of only one type of atom, existing as individual atoms or molecules.

Compound

A pure substance formed when two or more atoms of different elements combine chemically in a fixed and definite ratio, possessing properties distinct from its constituent elements.

Physical Properties

Characteristics of matter that can be measured or observed without changing the identity or composition of the substance, such as colour, odour, melting point, boiling point, and density.

Chemical Properties

Characteristics of matter whose measurement or observation requires a chemical change to occur, such as combustibility, acidity, or reactivity with acids and bases.

The International System of Units (SI)

A uniform decimal-based measurement system established in 1960 by the 11th CGPM comprising seven base physical quantities.

Metre Convention

A diplomatic treaty signed in Paris in 1875 that created the General Conference on Weights and Measures (CGPM) to establish standard international measurement units.

National Physical Laboratory (NPL), New Delhi

The National Metrology Institute (NMI) in India responsible for realizing base and derived units of measurement and maintaining National Standards of Measurement.

Mass

The amount of matter present in a substance, which remains constant regardless of location.

Weight

The force exerted by gravity on an object, which can vary from one place to another due to changes in gravitational acceleration.

Analytical Balance

A high-precision laboratory instrument used to determine the mass of substances accurately.

Volume

The amount of space occupied by a substance, having the SI unit of $\text{m}^3$ and commonly measured in $\text{dm}^3$, $\text{cm}^3$, or litres ($\text{L}$).

Laboratory Volume Measuring Devices

Apparatus used to measure or prepare specific liquid volumes, including burettes, pipettes, graduated cylinders, and volumetric flasks.

Density

The amount of mass per unit volume of a substance, calculated as $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$ with the SI unit $\text{kg}\,\text{m}^{-3}$.

Fahrenheit to Celsius Formula

The relationship converting degrees Celsius to degrees Fahrenheit: $^{\circ}\text{F} = \frac{9}{5}(^{\circ}\text{C}) + 32$.

Kelvin Scale Formula

The relationship between the SI unit Kelvin and degrees Celsius: $\text{K} = ^{\circ}\text{C} + 273.15$.

Scientific Notation

An exponential notation where any number is written in the form $N \times 10^n$, where $n$ is an integer exponent and $N$ is a digit term between $1.000...$ and $9.999...$.

Significant Figures

Meaningful digits in a measured or calculated quantity that are known with certainty plus one digit that is estimated or uncertain.

Precision

The closeness of various repeated measurements of the same quantity to one another.

Accuracy

The degree of agreement of a particular experimental or measured value with the true value of the result.

Dimensional Analysis

A mathematical technique (also called unit factor method or factor label method) used to convert units from one system to another using unit equivalence ratios.

Law of Conservation of Mass

A chemical law established by Antoine Lavoisier in 1789 stating that matter can neither be created nor destroyed in physical and chemical changes.

Law of Definite Proportions

A chemical law given by Joseph Proust stating that a given compound always contains exactly the same proportion of elements by weight, irrespective of its source.

Law of Multiple Proportions

A chemical law proposed by John Dalton in 1803 stating that if two elements combine to form more than one compound, the masses of one element combining with a fixed mass of the other are in the ratio of small whole numbers.

Gay Lussac’s Law of Gaseous Volumes

A law formulated in 1808 stating that when gases combine or are produced in a chemical reaction, they do so in a simple ratio by volume, provided all gases are at the same temperature and pressure.

Avogadro’s Law

A law proposed in 1811 stating that equal volumes of all gases at the same temperature and pressure should contain equal numbers of molecules.

Dalton's Atomic Theory

A theory published in 1808 in 'A New System of Chemical Philosophy' proposing that matter consists of indivisible atoms, identical for a given element, which combine in fixed ratios to form compounds.

Atomic Mass Unit (amu)

A unit of mass defined as exactly one-twelfth of the mass of one carbon-12 ($^{12}\text{C}$) atom, equal to approximately $1.66056 \times 10^{-24}\,\text{g}$ and now denoted by unified mass ($\text{u}$).

Average Atomic Mass

The weighted mean of the atomic masses of all naturally occurring isotopes of an element computed from their relative abundances.

Molecular Mass

The sum of the atomic masses of all elements present in a molecule, calculated by multiplying each element's atomic mass by the number of its atoms in the formula.

Formula Mass

The sum of atomic masses of the constituent atoms in a formula unit of a substance that does not contain discrete molecules, such as sodium chloride ($\text{NaCl}$).

Mole

The SI base unit (symbol $\text{mol}$) for the amount of substance containing exactly $6.02214076 \times 10^{23}$ elementary entities.

Avogadro Constant

The fixed numerical constant ($N_A = 6.02214076 \times 10^{23}\,\text{mol}^{-1}$) defining the exact number of elementary entities per mole of substance.

Molar Mass

The mass of one mole of a substance expressed in grams, numerically equal to its atomic, molecular, or formula mass in unified mass units ($\text{u}$).

Mass Per Cent of an Element

The mass fraction of an element in a compound multiplied by 100: $\text{Mass}\,\% = \frac{\text{Mass of that element in the compound}}{\text{Molar mass of the compound}} \times 100$.

Empirical Formula

The simplest chemical formula representing the simplest whole-number ratio of the various atoms present in a compound.

Molecular Formula

A chemical formula that shows the exact number of different types of atoms present in a single molecule of a compound.

Stoichiometry

The calculation of the relative masses and volumes of reactants and products involved in a balanced chemical reaction, derived from Greek words 'stoicheion' (element) and 'metron' (measure).

Stoichiometric Coefficients

The numerical coefficients placed before chemical formulas in a balanced chemical equation representing the relative mole or molecule ratios of participants.

Limiting Reagent

The reactant in a chemical reaction that is consumed first, thereby limiting the total quantity of products formed.

Mole Fraction

The concentration unit defined as the ratio of the number of moles of a particular component to the total number of moles of all components in the solution.

Molarity (M)

A temperature-dependent concentration unit defined as the number of moles of solute dissolved in $1\,\text{L}$ of solution: $\text{Molarity} = \frac{\text{Moles of solute}}{\text{Volume of solution in litres}}$.

Molality (m)

A temperature-independent concentration unit defined as the number of moles of solute present in $1\,\text{kg}$ of solvent: $\text{Molality} = \frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}$.

Stock Solution

A concentrated solution of known higher concentration from which solutions of lower desired concentrations are prepared via dilution.