Higher Secondary First Year Chemistry - Volume I Flashcards

This set of vocabulary flashcards covers essential concepts from the first volume of Grade 11 Chemistry, including atomic structure, the periodic table, chemical calculations, states of matter, and thermodynamics.

Unified atomic mass unit (amu or u)

Defined as one twelfth of the mass of a Carbon-12 atom in its ground state, approximately equal to $1.6605 \times 10^{-27}\,kg$.

Relative atomic mass ($A_r$)

The ratio of the average atomic mass to the unified atomic mass unit.

Mole

The SI unit representing the amount of substance containing $6.022 \times 10^{23}$ elementary particles, identical to the number of atoms in $12\,g$ of Carbon-12 isotope.

Avogadro number

The total number of entities present in one mole of any substance, equal to $6.022 \times 10^{23}$.

Molar Volume

The volume occupied by one mole of any substance in the gaseous state at a given temperature and pressure (e.g., $22.71\,dm^3$ at $273\,K$ and $1\,bar$).

Gram equivalent mass

The mass of an element, compound, or ion that combines with or displaces $1.008\,g$ hydrogen, $8\,g$ oxygen, or $35.5\,g$ chlorine.

Empirical formula

The chemical formula written with the simplest ratio of the number of different atoms present in one molecule of a compound.

Molecular formula

The chemical formula written with the actual number of different atoms present in one molecule of a compound.

Stoichiometry

The numerical relationship between chemical quantities (reactants and products) in a balanced chemical equation.

Limiting reagent

The reactant in a chemical reaction that is completely consumed first, thereby limiting the amount of product formed.

Oxidation

According to the classical concept, the addition of oxygen or removal of hydrogen; or according to the electronic concept, the loss of electrons.

Reduction

According to the classical concept, the removal of oxygen or addition of hydrogen; or according to the electronic concept, the gain of electrons.

Oxidation Number

The imaginary charge left on an atom when all other atoms of the compound have been removed in their usual oxidation states.

Disproportionation reaction

A redox reaction in which the same compound undergoes both oxidation and reduction, resulting in an increase and decrease of the oxidation state of the same element.

Stationary orbit

A certain circular path of fixed energy around the nucleus in which an electron can revolve without losing energy.

Heisenberg’s uncertainty principle

States that it is impossible to accurately determine both the position and the momentum of a microscopic particle simultaneously ($\Delta x \cdot \Delta p \ge \frac{h}{4\pi}$).

Orbital

A three-dimensional space around the nucleus in which the probability of finding the electron is maximum.

Principal quantum number ($n$)

Represents the main energy level or shell in which an electron revolves ($n = 1, 2, 3, \dots$).

Azimuthal quantum number ($l$)

Represents the subshell or orbital shape, taking integral values from $0$ to $(n-1)$.

Magnetic quantum number ($m_l$)

Represents the different orientations of orbitals in space, taking values from $-l$ to $+l$ through zero.

Spin quantum number ($m_s$)

Describes the spin of the electron about its own axis, either clockwise ($+\frac{1}{2}$) or anti-clockwise ($-\frac{1}{2}$).

Aufbau principle

States that in the ground state of atoms, orbitals are filled in the order of their increasing energies.

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of values for all four quantum numbers.

Hund’s rule of maximum multiplicity

States that electron pairing in degenerate orbitals does not occur until all available orbitals each contain one electron.

Effective nuclear charge ($Z_{eff}$)

The net positive charge experienced by valence electrons in the outermost shell after accounting for the shielding effect of inner electrons.

Ionization energy

The minimum energy required to remove the most loosely bound electron from the valence shell of an isolated neutral gaseous atom in its ground state.

Electron Affinity

The energy released (or required for noble gases) when an electron is added to the valence shell of an isolated neutral gaseous atom in its ground state.

Electronegativity

The relative tendency of an element in a covalently bonded molecule to attract the shared pair of electrons towards itself.

Diagonal Relationship

The similarity in properties observed between certain elements of the second period and diagonally placed elements of the third period (e.g., $Li$ and $Mg$).

Protium ($^1H_1$)

The predominant isotope of hydrogen, containing one proton and one electron but zero neutrons.

Deuterium ($^1H_2$ or $D$)

Also known as heavy hydrogen, an isotope containing one proton, one electron, and one neutron.

Tritium ($^1H_3$ or $T$)

A radioactive isotope of hydrogen with a half-life of $12.3\,years$, containing one proton, one electron, and two neutrons.

Ortho-hydrogen

A form of molecular hydrogen in which the spins of the two hydrogen nuclei are parallel.

Para-hydrogen

A form of molecular hydrogen in which the spins of the two hydrogen nuclei are anti-parallel.

Temporary Hardness of water

Hardness primarily due to the presence of soluble bicarbonates of magnesium and calcium, which can be removed by boiling.

Zeolites

Hydrated sodium alumino-silicates with a porous structure used as ion-exchange beds to remove water hardness.

Washing soda

Sodium carbonate decahydrate ($Na_{2}CO_{3} \cdot 10H_{2}O$).

Baking soda

Sodium bicarbonate ($NaHCO_{3}$).

Quick lime

Calcium oxide ($CaO$), produced by heating limestone.

Gypsum

A soft mineral composed of calcium sulphate dihydrate ($CaSO_{4} \cdot 2H_{2}O$).

Plaster of Paris

Calcium sulphate hemihydrate ($CaSO_{4} \cdot \frac{1}{2}H_{2}O$), obtained by heating gypsum to $393\,K$.

Boyle’s Law

At a given temperature, the volume occupied by a fixed mass of gas is inversely proportional to its pressure ($PV = k$).

Charles Law

At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature ($V = kT$).

Gay-Lussac’s Law

At constant volume, the pressure of a fixed mass of gas is directly proportional to its absolute temperature ($P = kT$).

Ideal gas equation

The general equation describing the state of any ideal gas: $PV = nRT$.

Dalton’s law of partial pressures

The total pressure of a mixture of non-reacting gases is the sum of the partial pressures of the individual component gases.

Graham’s Law of Diffusion

States that the rate of diffusion or effusion of a gas is inversely proportional to the square root of its molar mass.

Compressibility factor ($Z$)

The ratio of $PV$ to $nRT$ used to measure the deviation of real gases from ideal behavior ($Z = 1$ for ideal gases).

Critical temperature ($T_c$)

The temperature above which a gas cannot be liquefied, regardless of the pressure applied.

Joule-Thomson effect

The phenomenon of lowering the temperature when a compressed gas is allowed to expand adiabatically into a region of low pressure.

Isolated system

A thermodynamic system that can exchange neither matter nor energy with its surroundings.

Intensive property

A property of a system that is independent of the mass or size of the system (e.g., density, temperature).

Extensive property

A property of a system that depends on the mass or size of the system (e.g., volume, internal energy).

State function

A thermodynamic property possessing a specific value for a given state that is independent of the path taken to reach that state (e.g., enthalpy, internal energy).

Zeroth law of thermodynamics

States that if two systems are separately in thermal equilibrium with a third system, they are in thermal equilibrium with each other.

First law of thermodynamics

Also known as the law of conservation of energy, it states that energy can neither be created nor destroyed ($\Delta U = q + w$).

Enthalpy ($H$)

A thermodynamic property defined as the sum of internal energy ($U$) and the product of pressure and volume ($PV$).

Hess’s law of constant heat summation

States that the enthalpy change of a reaction is the same regardless of whether it occurs in one step or multiple steps, provided the initial and final states are same.

Entropy ($S$)

A measure of the molecular disorder or randomness of a system ($dS = \frac{dq_{rev}}{T}$).

Second law of thermodynamics

States that the entropy of the universe increases during a spontaneous process.

Gibbs free energy ($G$)

A state function used to predict spontaneity, defined as $G = H - TS$. A negative $\Delta G$ indicates a spontaneous reaction.

Third law of thermodynamics

States that the entropy of a pure crystalline substance at absolute zero ($0\,K$) is zero.

Lattice energy

The energy required to completely separate one mole of an ionic crystal into its constituent gaseous ions.