Physical Chemistry Definitions Flashcards
Atomic / Proton number
Number of protons in the nucleus of an atom.
Mass / Nucleon number
Sum of protons and nucleus in the nucleus of an atom.
Isotopes
Atoms of an element that have the same number of protons but different number of neutrons and hence different nucleon numbers.
S orbitals
Orbital that is spherically symmetrical about the nucleus. The larger the principal quantum number, the orbitals get larger and more diffused.
P orbitals
Three dumbbell shaped orbitals, mutually at right angles to each other. All three orbitals are degenerate, ie have the same energy.
d orbitals
d orbitals have different shapes
First Ionisation Energy
The energy needed to remove one mole of electrons from one mole of the gaseous atoms to form one mole of singly charged gaseous cations.
Relative isotopic mass of an isotope
The mass of one atom of the isotope to 1-12th the mass of one atom of C-12 isotope
Relative atomic mass of an element
The average mass of one atom of the element to 1-12th the mass of one atom of C-12 isotope
Relative molecular mass of a substance
The average mass of one molecule of the substance to 1-12th the mass of one atom of C-12 isotope
Relative formula mass of a substance
The average mass of one formula unit of the substance to 1-12th the mass of one atom of C-12 isotope
One mole
The amount of substance which contains the same number of particles as there are atoms in exactly 12g of C-12
Avogadro’s law
Equal volumes of all gases under the same conditions of temperature and pressure contain the same number of atoms/molecules.
Empirical formula
The simplest formula that shows the relative number of atoms of each element present in a compound.
Molecular formula
The actual number of atoms present in one molecules of the compound
Metallic bonds
Strong electrostatic forces of attraction between the metal ions and the ‘sea’ of delocalised electrons in a giant metallic structure.
Ionic bonds
Strong electrostatic forces of attraction between cations and anions in the crystal lattice of an ionic solid
Covalent bond
Strong electrostatic forces of attraction between a shared pair of electrons and two positively charged nuclei
Lewis Acid
An atom with vacant low lying orbital to accept a lone pair of electron
Lewis base
An atom with a lone pair of electrons, available for donation
Assumptions of Kinetic Theory applied to an Ideal Gas
Volume of gas molecules is negligible compared to the volume of the container.
Forces of attraction between the gas molecules as well as between the gas molecules and the walls of the container are negligible.
All molecular collisions are perfectly elastic. There is no loss of kinetic energy during collision.
Enthalpy Change of Reaction
The amount of heat absorbed or evolved when molar quantities of reactants as shown in the chemical equation react together.
Enthalpy Change of Formation
The amount of heat absorbed or evolved when one mole of a substance is formed from its constituent elements.
Enthalpy Change of Combustion
The amount of heat evolved when one mole of a substance in its standard state is completely burned in excess oxygen.
Enthalpy Change of Neutralisation
The amount of heat evolved when one mole of water is formed from the neutralisation between an acid and a base.
Enthalpy Change of Atomisation
The amount of heat absorbed when one mole of free gaseous atoms is formed from its element.
Bond Energy
The average amount of heat absorbed to break one mole of that particular bond in a particular compound in the gaseous state
Electron Affinity
The amount of heat evolved when 1 mole of electrons is added to one mole of gaseous atoms to form one mole of singly negatively charged ions
Lattice energy
The amount of heat evolved when one mole of the solid ionic compound is formed from its constituent free gaseous ions
Enthalpy change of Hydration
The amount of heat evolved when one mole of free gaseous ions is dissolved in a large amount of water forming a solution at infinite dilution
Enthalpy change of Solution
The amount of heat absorbed or evolved when one mole of a solute is dissolved in an infinite volume of water
Hess’ law
The enthalpy change of a chemical reaction is dependent only on the initial states of the reactants and the final state of the products and is independent of the reaction pathway taken.
Entropy
The degree of disorder, of matter and ways to distribute energy in the system
Rate equation
Relation of the rate of reaction to the concentration of reactants raised to the appropriate power
Order of reaction
The power to which the concentration of the reactant is raised in the experimentally determined rate equation.
Dynamic equilibrium
A state where the rate of forward reaction equals to the rate of backwards reaction for a reversible reaction such that there is no net change in the concentration of reactants and products
Le Chatelier’s Principle
If a change occurs in one of the conditions under which a reversible reaction is in dynamic equilibrium, the position of equilibrium shifts so as to minimise the change.
Saturated solution
A solution containing a maximum amount of solute in a given amount of solvent such that the ions are in equilibrium with the solid
Solubility
Solubility of a substance is the maximum amount of solute that dissolves in 1dm^3 of water to form a saturated solution at a stated temperature
Arrhenius Acid
A substance that releases H+ ions
Arrhenius Base
A substance that releases OH- ions
Bronsted Acid
Proton H+ donor
Bronsted Base
Proton H+ acceptor
Definition of pH
-log [H+]
Buffer solution
A solution where pH remains almost constant when small amounts of acid or alkali are added to it
Suitable indicator
An indicator that has a working pH range which lies within the range of rapid pH change for the titration
Standard electrode potential
The potential difference between a standard hydrogen electrode and the half cell in which reacting species are at molar concentrations of 1 mol dm^-3, 298k and 1bar.
