ISC Physical Sciences — Chemistry Paper 2 Definitions and Concepts

Comprehensive set of 440 practice vocabulary flashcards covering all definitions, functional concepts, equations, laboratory apparatus functions, reaction conditions, and examination principles from the ISC Chemistry Paper 2 curriculum.

Mole

The SI unit for amount of substance, where one mole contains $6{,}02 \times 10^{23}$ specified elementary entities.

Molar Mass

The mass in grams of one mole of that substance, represented by the symbol $M$ with the unit $\text{g}\cdot\text{mol}^{-1}$.

Molar Concentration

The amount of solute per unit volume of solution, expressed with the unit $\text{mol}\cdot\text{dm}^{-3}$.

Solution

A homogenous mixture of solute and solvent with uniform composition.

Solute

The substance that is dissolved in the solution.

Solvent

The substance in which another substance is dissolved, forming a solution.

Avogadro's Constant

The number of elementary entities in one mole of substance, equal to $6{,}02 \times 10^{23}\,\text{mol}^{-1}$.

Molar Volume at STP

One mole of gas occupies $22{,}4\,\text{dm}^3$ at $0\,^\circ\text{C}$ ($273\,\text{K}$) and $1\,\text{atmosphere}$ ($101{,}3\,\text{kPa}$).

Percentage Yield

A measure of the extent of a reaction, determined by comparing the actual amount of product to the theoretical amount of product, expressed as a percentage.

Percentage Purity

The percentage of a sample that is the specified pure substance, calculated as $(\text{mass of pure substance} \div \text{mass of impure sample}) \times 100$.

Limiting Reagent

The reactant that is used up first and limits the amount of product formed.

Formula for Moles from Mass

$n = \frac{m}{M}$, where $n$ is the number of moles, $m$ is the mass of substance in grams, and $M$ is the molar mass in $\text{g}\cdot\text{mol}^{-1}$.

Formula for Number of Particles

$N = n \times N_A$, where $N$ is the number of particles, $n$ is the number of moles, and $N_A$ is Avogadro's constant ($6{,}02 \times 10^{23}\,\text{mol}^{-1}$).

Formula for Moles of Gas at STP

$n = \frac{V}{V_m}$, where $V$ is the gas volume in $\text{dm}^3$ and $V_m$ is the molar volume at STP ($22{,}4\,\text{dm}^3\cdot\text{mol}^{-1}$).

Formula for Molar Concentration ($n$ and $V$)

$c = \frac{n}{V}$, where $c$ is the molar concentration in $\text{mol}\cdot\text{dm}^{-3}$, $n$ is the solute moles, and $V$ is the solution volume in $\text{dm}^3$.

Formula for Molar Concentration ($m$, $M$, and $V$)

$c = \frac{m}{M \times V}$, where $m$ is the mass of solute, $M$ is the molar mass, and $V$ is the volume in $\text{dm}^3$.

Standard Temperature Value at STP

$0\,^\circ\text{C}$ or $273\,\text{K}$.

Standard Pressure Values at STP

$1\,\text{atmosphere}$ or $101{,}3\,\text{kPa}$.

Formula for Percentage Yield

$\text{\% yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$.

Formula for Percentage Purity

$\text{\% purity} = \frac{\text{mass of pure substance}}{\text{mass of impure sample}} \times 100$.

Aqueous Solution Solvent

Water ($H_2O$).

Stock Notation

A system of chemical nomenclature indicating the oxidation state of an element using Roman numerals in parentheses.

Ammonium Ion

A polyatomic cation with the formula $NH_4^+$.

Hydronium Ion

A polyatomic cation with the formula $H_3O^+$.

Chlorate Ion

A polyatomic anion with the formula $ClO_3^-$.

Ethanoate Ion

A polyatomic anion with the formula $CH_3COO^-$.

Hydroxide Ion

A polyatomic anion with the formula $OH^-$.

Nitrate Ion

A polyatomic anion with the formula $NO_3^-$.

Nitrite Ion

A polyatomic anion with the formula $NO_2^-$.

Permanganate Ion

A polyatomic anion with the formula $MnO_4^-$.

Carbonate Ion

A polyatomic anion with the formula $CO_3^{2-}$.

Hydrogen Carbonate Ion

A polyatomic anion with the formula $HCO_3^-$.

Sulfate Ion

A polyatomic anion with the formula $SO_4^{2-}$.

Hydrogen Sulfate Ion

A polyatomic anion with the formula $HSO_4^-$.

Thiosulfate Ion

A polyatomic anion with the formula $S_2O_3^{2-}$.

Sulfite Ion

A polyatomic anion with the formula $SO_3^{2-}$.

Phosphate Ion

A polyatomic anion with the formula $PO_4^{3-}$.

Dichromate Ion

A polyatomic anion with the formula $Cr_2O_7^{2-}$.

State Symbol (s)

Denotes that a substance is in the solid phase.

State Symbol (l) or (\ell)

Denotes that a substance is in the pure liquid phase.

State Symbol (g)

Denotes that a substance is in the gas phase.

State Symbol (aq)

Denotes that a substance is dissolved in water, forming an aqueous solution.

Conversion: \text{cm}^3 to \text{dm}^3

Divide the volume value in $\text{cm}^3$ by $1\,000$ (or multiply by $10^{-3}$) to obtain volume in $\text{dm}^3$.

Law of Conservation of Matter

The principle that mass/atoms are conserved in chemical reactions, allowing calculations of product and reactant masses.

Theoretical Yield

The calculated maximum amount of product that can be formed from the given amounts of reactants according to balanced equation mole ratios.

Actual Yield

The experimentally measured quantity of product obtained from a chemical reaction.

Inert Impurity

Unreactive material present within an impure sample that does not take part in the desired reaction.

Elementary Entities

Specified particles such as atoms, molecules, ions, or formula units counted in a mole.

Relative Atomic Mass

The weighted average mass of an atom of an element compared to one-twelfth the mass of an atom of carbon-12.

Stoichiometric Coefficient

The number written in front of a chemical formula in a balanced equation indicating the relative molar ratio.

Method for Identifying Limiting Reagent

Divide available moles of each reactant by its stoichiometric coefficient; the reactant with the smallest ratio is limiting.

Total Solution Volume vs Solvent Volume

Total solution volume is the final volume of the combined solute and solvent, which must be used in $c = \frac{n}{V}$, not pure solvent volume.

Excess Reagent

A reactant that is not completely consumed when the reaction runs to completion because more is present than required by stoichiometry.

Balanced Chemical Equation

A chemical equation with equal numbers of each type of atom on both reactant and product sides, satisfying the conservation of matter.

Molar Volume Symbol

$V_m$, representing the volume occupied by one mole of gas under specified conditions.

Intramolecular Bond

A bond between atoms within molecules; for the curriculum, only covalent bonds are considered intramolecular.

Electronegativity

A measure of the tendency of an atom to attract a bonding pair of electrons.

Covalent Bond

The sharing of at least one pair of electrons by two atoms.

Non-polar (Pure) Covalent Bond

A covalent bond where electrons are equally shared due to zero difference in electronegativity.

Polar Covalent Bond

A covalent bond where electrons are unequally shared due to a difference in electronegativity.

Bond Polarity Rule

The greater the difference in electronegativity between two bonded atoms, the more polar the covalent bond.

Ionic Bonding

A transfer of electrons to form cations and anions which are electrostatically attracted in an ionic crystal lattice.

Participants in Ionic Bonding

Generally formed between metals (or ammonium cations) and non-metals (or polyatomic anions).

Ionic Crystal Lattice

A rigid 3D repeating regular arrangement of alternating cations and anions held together by strong electrostatic attractions.

Metallic Bonding

The attraction between positive kernels and a sea of delocalised electrons.

Positive Kernel

The nucleus and inner core electrons of a metal atom, excluding its delocalised valence electrons.

Sea of Delocalised Electrons

Valence electrons in metals that are free to move throughout the entire metallic structure.

Intermolecular Force

A force of attraction between molecules or between atoms of noble gases.

Van der Waals Forces

Intermolecular forces including London forces, dipole–dipole forces, and hydrogen bonding as a special case.

London Forces

Intermolecular forces arising from instantaneous dipoles which induce dipoles in neighbouring particles.

Origin of London Forces

Random, temporary fluctuations in electron distribution forming temporary (instantaneous) dipoles that induce dipoles in adjacent particles.

Factors Influencing Strength of London Forces

The size of temporary dipoles, which is determined by the number of electrons (electron cloud size/polarisability) and interacting molecular surface area.

Dipole–Dipole Forces

Intermolecular electrostatic attractions between permanent dipoles of polar molecules.

Hydrogen Bonding

An intermolecular attraction involving a hydrogen atom covalently bonded to N, O or F in one molecule and a lone pair on N, O or F of another molecule.

Atoms That Hydrogen Bond with Hydrogen

Nitrogen ($N$), oxygen ($O$), or fluorine ($F$), which are small, highly electronegative atoms possessing at least one lone pair of electrons.

Reasons Hydrogen Bonds are Stronger than Dipole–Dipole Forces

The high polarity of the covalent bond between H and N, O, or F, and the fact that molecules can approach very closely.

Polar Molecule (Dipole)

A molecule with an overall non-zero dipole moment where electrons are unevenly distributed, resulting in $\delta^+$ and $\delta^-$ ends.

Asymmetrical Molecule Polarity Rule

Asymmetrical molecules are considered polar under the curriculum (except specified non-polar bond cases).

Symmetrical Molecule Polarity Rule

Symmetrical molecules (including hydrocarbons) are non-polar because bond dipoles cancel out.

Giant Covalent Structures (Melting Points)

Substances like diamond, graphite, and silicon dioxide with high melting points due to the large energy required to break many strong covalent bonds.

Giant Ionic Solids (Melting Points)

Determined by the strength of the many electrostatic forces of attraction (ionic bonds) between cations and anions in the crystal lattice.

Simple Molecular Substances (Melting/Boiling Points)

Determined by the strength and the relative number per molecule (for hydrogen bonds) of intermolecular forces.

Boiling Point Trend: Same Group or Homologous Series

Mainly governed by London forces increasing with molar mass/electron cloud size, except when hydrogen bonding is present.

Boiling Point Trend: Similar Mass Across Different Homologous Series

Differences are due to the presence of additional stronger intermolecular forces (dipole–dipole or hydrogen bonding) beside London forces.

Hydrogen Bond Density Effect

Substances forming more hydrogen bonds per molecule exhibit higher melting and boiling points compared to those forming fewer.

Electrical Conductivity of Metals and Graphite

Arises from the presence of mobile delocalised valence electrons that move when a potential difference is applied.

Electrical Conductivity of Ionic Compounds

Conduct electricity only when molten or in aqueous solution because ions become mobile; do not conduct as solids.

Instantaneous (Temporary) Dipole

A temporary unsymmetrical distribution of electron density in an atom or molecule creating a momentary polarity.

Induced Dipole

A dipole created in a neighbouring non-polar particle caused by electrostatic repulsion or attraction from a nearby dipole.

Intermolecular Forces in All Molecules

London forces are present between all atoms and molecules regardless of polarity.

Diamond Structure

A giant covalent network of carbon atoms each tetrahedrally bonded to four other carbon atoms by strong covalent bonds.

Graphite Structure

A giant covalent structure of carbon layers arranged in hexagonal rings with delocalised electrons between sheets.

Silicon Dioxide ($SiO_2$)

A giant covalent substance with high melting point requiring large amounts of energy to break strong Si-O bonds.

Predominant Intermolecular Force

The strongest type of intermolecular force operating between molecules in a substance.

Dipole Moment

A quantitative measure of the separation of positive and negative electrical charges in a system or molecule.

Polarisability

The ease with which the electron cloud of an atom or molecule can be distorted to form a dipole.

Surface Area of Contact (Bonding)

The extent of molecular perimeter available for London force interactions; straight-chain molecules have greater contact than branched isomers.

Noble Gas Intermolecular Forces

Weak London forces arising between individual unbonded noble gas atoms.

Hydrocarbon Molecular Polarity

Hydrocarbons are non-polar symmetrical molecules experiencing only London intermolecular forces.

Non-examined Molecules with Non-polar Bonds

Asymmetrical molecules with non-polar bonds such as $PH_3$, $H_2Te$, and $NCl_3$ are excluded from polarity examination.