∆H = Hproducts – Hreactants
Activation energy of a reaction: the minimum energy required for reaction to occur. It represents
the energy barrier that the reactants must overcome in order to become products.
Bond breaking is an endothermic process; bond making is an exothermic process.
Bond enthalpy is a measure of energy required to break bond. The values quoted are usually
average bond enthalpies, as the strength of a bond between two particular atoms is different in
different molecules. Some bond energies are exact and some bond energies are averages.
∆Hr = ∑B.E (reactants) – ∑B.E (products)
Bond energy is the energy needed to break bonds or released when these bonds form.
q = mc∆T ∆Hr = – mc∆T
Standard conditions = 100 kPa & 298K (25 oC)
Standard enthalpy change of reaction, ∆Hrθ = enthalpy change when amounts of reactants, as
shown in the reaction equation, react together at 25°C (298K) &1 atm (100 kPa) / under standard
conditions) to give products in their standard states.
Standard enthalpy of combustion, ∆Hcθ (always positive) = heat energy evolved when 1 mole of the
substance is completely burnt in oxygen / burnt in excess oxygen under standard conditions
Standard enthalpy of formation, ∆Hfθ = heat change when 1 mole of a substances formed from its
elements in their standard states under standard conditions
- Usually exothermic
Standard enthalpy of change of solution = enthalpy change when 1 mole of a solute is dissolved in
a large amount of solvent under standard conditions
- The solvent must be in excess to make sure all the solute dissolves
- Exothermic or endothermic
- Soluble: negative/small positive value
- Insoluble: large positive values
Enthalpy of change of hydration = enthalpy change when change when 1 mole of a specified
gaseous ion dissolves in sufficient water to form a very dilute solution.
- Always exothermic
- When an ionic solid dissolves in water, bonds are formed between water molecules and the
ions. These bonds are called ion-dipole bonds.
- Water is polar molecule. The δ– oxygen atoms in water molecules are attracted to the positive
ions in the ionic compound. The δ+ hydrogen atoms in the water molecules are attracted to the
negative ions in the ionic compound.
- The energy released in forming ion-dipole bonds is sufficient to compensate for the energy that
must be put in to separate the anions and cations that are bonded together in the crystal
lattice.
- The value of ∆Hθhyd is more exothermic for ions with the same charge but smaller ionic radii
eg. ∆Hθhyd is more exothermic for Li+ than Na+
- The value of ∆Hθhyd is more exothermic for ions with the same radii but a larger charge
eg. ∆Hθhyd is more exothermic for Mg2+ than Li+
, Standard enthalpy of change of neutralisation = enthalpy change when 1 mole of a water is formed
by the reaction of an acid with an alkali under standard conditions.
- Always exothermic
- The heat of neutralisation between strong acids and strong alkali are always the same.
- Complete neutralisation of a strong diprotic acid with alkali produces double amount of heat as
compared to a strong monoprotic acid. This is because these acids produce 2 moles of
hydrogen ions when it dissociates in water.
- The heat given out when strong acid reacts with a strong alkali is higher than the heat given out
when a weak acid reacts with a strong alkali (more negative ∆Hneut value)
Standard enthalpy of change of atomisation = enthalpy change when 1 mole of gaseous atom is
formed from its element under standard conditions.
- Standard enthalpy change of atomisation of diatomic gases = 1⁄2 x bond energy.
- The standard enthalpy change of atomisation of the noble gas (Group VIII) is zero, because all
of them exist as monoatomic gases under standard condition.
- The enthalpy change of atomisation is always endothermic because energy must be supplied to
break the bonds holding the atoms in the element together.
Standard enthalpy of change of dissociation = enthalpy change when 1 mole of a gaseous
substance is completely decomposed into gaseous atoms under standard conditions.
- Standard enthalpy change of dissociation of diatomic gases = bond energy
- The enthalpy change of dissociation is always endothermic because energy must be supplied to
break the bonds holding the atoms in the compound together.
First ionisation energy = the energy needed to remove 1 electron from 1 mole of gaseous atoms of
the element to form 1 mole of gaseous unipositive ions
- Always endothermic
Second ionisation energy = the energy needed to remove 1 electron from 1 mole of gaseous ions
with a single positive charge to form 1 mole of gaseous ions with 2+ charge
- Always endothermic
Electron affinities depend on
- the distance of the electron from the nucleus (size of the atom)
- the size of the positive nuclear charge
- the symmetry of the electronic configuration
With increase in the atomic size, the distance between the nucleus and the incoming
electron also increases → the attraction between the nucleus and the incoming electron
decreases → the electron affinity will have smaller value
For example, potassium has higher atomic size than sodium. The electron affinity of potassium is
lower than that of sodium. Thus, in the group of alkali metals, on moving from top to bottom, with
increase in the atomic size, the electron affinity values decreases. Similar
is the trend observed for halogens.
As the nuclear charge increases, the force of attraction between the nucleus and incoming
electron increases → electron affinity increases
Atoms having half-filled or completely filled subshells in the valence shell have stable electronic
configuration → little tendency to accept an additional electron → very low values of electron
affinity
Activation energy of a reaction: the minimum energy required for reaction to occur. It represents
the energy barrier that the reactants must overcome in order to become products.
Bond breaking is an endothermic process; bond making is an exothermic process.
Bond enthalpy is a measure of energy required to break bond. The values quoted are usually
average bond enthalpies, as the strength of a bond between two particular atoms is different in
different molecules. Some bond energies are exact and some bond energies are averages.
∆Hr = ∑B.E (reactants) – ∑B.E (products)
Bond energy is the energy needed to break bonds or released when these bonds form.
q = mc∆T ∆Hr = – mc∆T
Standard conditions = 100 kPa & 298K (25 oC)
Standard enthalpy change of reaction, ∆Hrθ = enthalpy change when amounts of reactants, as
shown in the reaction equation, react together at 25°C (298K) &1 atm (100 kPa) / under standard
conditions) to give products in their standard states.
Standard enthalpy of combustion, ∆Hcθ (always positive) = heat energy evolved when 1 mole of the
substance is completely burnt in oxygen / burnt in excess oxygen under standard conditions
Standard enthalpy of formation, ∆Hfθ = heat change when 1 mole of a substances formed from its
elements in their standard states under standard conditions
- Usually exothermic
Standard enthalpy of change of solution = enthalpy change when 1 mole of a solute is dissolved in
a large amount of solvent under standard conditions
- The solvent must be in excess to make sure all the solute dissolves
- Exothermic or endothermic
- Soluble: negative/small positive value
- Insoluble: large positive values
Enthalpy of change of hydration = enthalpy change when change when 1 mole of a specified
gaseous ion dissolves in sufficient water to form a very dilute solution.
- Always exothermic
- When an ionic solid dissolves in water, bonds are formed between water molecules and the
ions. These bonds are called ion-dipole bonds.
- Water is polar molecule. The δ– oxygen atoms in water molecules are attracted to the positive
ions in the ionic compound. The δ+ hydrogen atoms in the water molecules are attracted to the
negative ions in the ionic compound.
- The energy released in forming ion-dipole bonds is sufficient to compensate for the energy that
must be put in to separate the anions and cations that are bonded together in the crystal
lattice.
- The value of ∆Hθhyd is more exothermic for ions with the same charge but smaller ionic radii
eg. ∆Hθhyd is more exothermic for Li+ than Na+
- The value of ∆Hθhyd is more exothermic for ions with the same radii but a larger charge
eg. ∆Hθhyd is more exothermic for Mg2+ than Li+
, Standard enthalpy of change of neutralisation = enthalpy change when 1 mole of a water is formed
by the reaction of an acid with an alkali under standard conditions.
- Always exothermic
- The heat of neutralisation between strong acids and strong alkali are always the same.
- Complete neutralisation of a strong diprotic acid with alkali produces double amount of heat as
compared to a strong monoprotic acid. This is because these acids produce 2 moles of
hydrogen ions when it dissociates in water.
- The heat given out when strong acid reacts with a strong alkali is higher than the heat given out
when a weak acid reacts with a strong alkali (more negative ∆Hneut value)
Standard enthalpy of change of atomisation = enthalpy change when 1 mole of gaseous atom is
formed from its element under standard conditions.
- Standard enthalpy change of atomisation of diatomic gases = 1⁄2 x bond energy.
- The standard enthalpy change of atomisation of the noble gas (Group VIII) is zero, because all
of them exist as monoatomic gases under standard condition.
- The enthalpy change of atomisation is always endothermic because energy must be supplied to
break the bonds holding the atoms in the element together.
Standard enthalpy of change of dissociation = enthalpy change when 1 mole of a gaseous
substance is completely decomposed into gaseous atoms under standard conditions.
- Standard enthalpy change of dissociation of diatomic gases = bond energy
- The enthalpy change of dissociation is always endothermic because energy must be supplied to
break the bonds holding the atoms in the compound together.
First ionisation energy = the energy needed to remove 1 electron from 1 mole of gaseous atoms of
the element to form 1 mole of gaseous unipositive ions
- Always endothermic
Second ionisation energy = the energy needed to remove 1 electron from 1 mole of gaseous ions
with a single positive charge to form 1 mole of gaseous ions with 2+ charge
- Always endothermic
Electron affinities depend on
- the distance of the electron from the nucleus (size of the atom)
- the size of the positive nuclear charge
- the symmetry of the electronic configuration
With increase in the atomic size, the distance between the nucleus and the incoming
electron also increases → the attraction between the nucleus and the incoming electron
decreases → the electron affinity will have smaller value
For example, potassium has higher atomic size than sodium. The electron affinity of potassium is
lower than that of sodium. Thus, in the group of alkali metals, on moving from top to bottom, with
increase in the atomic size, the electron affinity values decreases. Similar
is the trend observed for halogens.
As the nuclear charge increases, the force of attraction between the nucleus and incoming
electron increases → electron affinity increases
Atoms having half-filled or completely filled subshells in the valence shell have stable electronic
configuration → little tendency to accept an additional electron → very low values of electron
affinity
