1. STPM Chemistry Form 6 Definition List:
1. The Hund Rule ± orbital with the same energy level (degenerate orbitals), electron will
occupy different orbital singly/one electron first with the parallel spin, before
pairing.(Reject: same spin, spin with same direction) STPM 2008, 2007
2. Pauli Exclusive Principle ± Each orbital can hold a maximum of two electrons with
opposite spin. (Reject: different spin) STPM 2007
3. Aufbau principle ± Electrons occupy orbitals with the lowest energy level first. (Reject:
occupy lower energy first) STPM 2006, 2007
4. Vapour pressure ± the pressure exerted by a vapour that is in equilibrium with its liquid
at a fixed temperature in a closed system. It is due to the collision of the vapour particles
on the walls of the container. STPM 2004
STOICHIOMETRY
1. Molarity (M) [concentration of a fluid solution] ± defined as the moles of a solute per
volume of solution.
2. Density [concentration of a fluid solution] ± defined as the mass of solution per volume
of solution.
3. Dilution ± addition of solvent to a solution resulting in an increase in the volume of the
solution and a decrease in the concentration of the solute in solution.
4. Precipitation reaction (also refer to as double-displacement reaction) ± a reaction that
involve two aqueous salts being added together to form ions and a solid salt precipitate.
ACID-BASE EQUILIBRIA
1. Arrhenius acid ± acid yields H3O+ when added to water
2. Arrhenius base ± base yields OH- when added to water
3. Bronsted-Lowry acid (proton donor) ± a substance that donate a proton, H+ to a base.
4. Bronsted-Lowry base (proton acceptor) ± a substance that accept a proton, H+ from an
acid.
5. Conjugate acid ± a species (molecule or ion) formed when a proton is added to a base.
6. Conjugate base ± a species (molecule or ion) formed when a proton is removed from an
acid.
7. Lewis acid ± an electron-pair acceptor.
8. Lewis base ± an electron-pair donor.
9. Electrolyte (also electrolytic conductor) ± a chemical compound that will conduct
electricity in the molten state or in aqueous solution.
10. Strong acid ± an acid that is almost completely dissociated in aqueous solution.
(Stronger the acid, the weaker its conjugate base)
11. Weak acid ± an acid that is only partially dissociated in aqueous solution. (Weaker the
acid, the stronger its conjugate base)
2. STPM Chemistry Form 6 Definition List ± Part 2:
Acid-Base Equilibria
Term Definition Example
Arrhenius acid Yields H3O+ when added to aq: [H3O+] > [OH-]
H2O
Arrhenius base Yields OH- when added to aq: [OH-] > [H3O+]
H2O
Bronsted-Lowry acid Proton donor HX in protic solvent
Bronsted-Lowry base Proton acceptor KOH in protic solvent
Lewis acid Electron pair acceptor BF3 in aprotic solvent
Lewis base Electron pair donor NH3 in aprotic solvent
Four important concepts (just the berry essence):
1. Acid dissociation ± (expressed quantitatively) acid dissociation constant, Ka, is nothing
more than the equilibrium constant for the dissociation reaction of an acid in water.
Relative strength of an acid (increases), its Ka (increases) and its pKa (decreases).
(The Ka and pKa of an acid depend on the strength of an acid, but not its concentration.)
2. Base hydrolysis ± base hydrolysis constant, Kb, nothing more than the equilibrium
constant for the hydrolysis reaction of a base in water.
Relative strength of a base (increases), its Kb (increases) and its pKb (decreases).
(The Kb and pKb of an acid depend on the strength of an acid, but not its concentration.)
3. Overall Relationship : Acid strength , Ka , pKa , conjugate base strength , Kb , pKb
4. Strength of a reagent (Ka / pKa and Kb / pKb) ± the completeness of a reaction in water.
(dissociation = ionisation or electrolytic nature)
The stronger the acid, the more electrolytic it is, because it conducts electricity better due
to the greater number of ions in solution.
The stronger the base, the more readily it undergoes hydrolysis when mixed with water
Acid Name pKa
Cl3CCOOH Trichloroacetic acid 0.64
Cl2HCCOOH Dichloroacetic acid 1.27
H2SO3 Sulfurous acid 1.82
HClO2 Chloroacetic acid 1.90
ClH2CCOOH Chloroacetic acid 2.82
HF Hydrofluoric acid 3.15
HNO2 Nitrous acid 3.41
HCOOH Formic acid 3.74
H3CCOOH Acetic acid 4.74
![STPM Chemistry Form 6 Definition List:
1. The Hund Rule ± orbital with the same energy level (degenerate orbitals), electron will
occupy different orbital singly/one electron first with the parallel spin, before
pairing.(Reject: same spin, spin with same direction) STPM 2008, 2007
2. Pauli Exclusive Principle ± Each orbital can hold a maximum of two electrons with
opposite spin. (Reject: different spin) STPM 2007
3. Aufbau principle ± Electrons occupy orbitals with the lowest energy level first. (Reject:
occupy lower energy first) STPM 2006, 2007
4. Vapour pressure ± the pressure exerted by a vapour that is in equilibrium with its liquid
at a fixed temperature in a closed system. It is due to the collision of the vapour particles
on the walls of the container. STPM 2004
STOICHIOMETRY
1. Molarity (M) [concentration of a fluid solution] ± defined as the moles of a solute per
volume of solution.
2. Density [concentration of a fluid solution] ± defined as the mass of solution per volume
of solution.
3. Dilution ± addition of solvent to a solution resulting in an increase in the volume of the
solution and a decrease in the concentration of the solute in solution.
4. Precipitation reaction (also refer to as double-displacement reaction) ± a reaction that
involve two aqueous salts being added together to form ions and a solid salt precipitate.
ACID-BASE EQUILIBRIA
1. Arrhenius acid ± acid yields H3O+ when added to water
2. Arrhenius base ± base yields OH- when added to water
3. Bronsted-Lowry acid (proton donor) ± a substance that donate a proton, H+ to a base.
4. Bronsted-Lowry base (proton acceptor) ± a substance that accept a proton, H+ from an
acid.
5. Conjugate acid ± a species (molecule or ion) formed when a proton is added to a base.
6. Conjugate base ± a species (molecule or ion) formed when a proton is removed from an
acid.
7. Lewis acid ± an electron-pair acceptor.
8. Lewis base ± an electron-pair donor.
9. Electrolyte (also electrolytic conductor) ± a chemical compound that will conduct
electricity in the molten state or in aqueous solution.
10. Strong acid ± an acid that is almost completely dissociated in aqueous solution.
(Stronger the acid, the weaker its conjugate base)
11. Weak acid ± an acid that is only partially dissociated in aqueous solution. (Weaker the
acid, the stronger its conjugate base)](https://image.slidesharecdn.com/10067578-111108013553-phpapp02/85/57996218-STPM-Chemistry-Form-6-1-320.jpg)
![STPM Chemistry Form 6 Definition List ± Part 2:
Acid-Base Equilibria
Term Definition Example
Arrhenius acid Yields H3O+ when added to aq: [H3O+] > [OH-]
H2O
Arrhenius base Yields OH- when added to aq: [OH-] > [H3O+]
H2O
Bronsted-Lowry acid Proton donor HX in protic solvent
Bronsted-Lowry base Proton acceptor KOH in protic solvent
Lewis acid Electron pair acceptor BF3 in aprotic solvent
Lewis base Electron pair donor NH3 in aprotic solvent
Four important concepts (just the berry essence):
1. Acid dissociation ± (expressed quantitatively) acid dissociation constant, Ka, is nothing
more than the equilibrium constant for the dissociation reaction of an acid in water.
Relative strength of an acid (increases), its Ka (increases) and its pKa (decreases).
(The Ka and pKa of an acid depend on the strength of an acid, but not its concentration.)
2. Base hydrolysis ± base hydrolysis constant, Kb, nothing more than the equilibrium
constant for the hydrolysis reaction of a base in water.
Relative strength of a base (increases), its Kb (increases) and its pKb (decreases).
(The Kb and pKb of an acid depend on the strength of an acid, but not its concentration.)
3. Overall Relationship : Acid strength , Ka , pKa , conjugate base strength , Kb , pKb
4. Strength of a reagent (Ka / pKa and Kb / pKb) ± the completeness of a reaction in water.
(dissociation = ionisation or electrolytic nature)
The stronger the acid, the more electrolytic it is, because it conducts electricity better due
to the greater number of ions in solution.
The stronger the base, the more readily it undergoes hydrolysis when mixed with water
Acid Name pKa
Cl3CCOOH Trichloroacetic acid 0.64
Cl2HCCOOH Dichloroacetic acid 1.27
H2SO3 Sulfurous acid 1.82
HClO2 Chloroacetic acid 1.90
ClH2CCOOH Chloroacetic acid 2.82
HF Hydrofluoric acid 3.15
HNO2 Nitrous acid 3.41
HCOOH Formic acid 3.74
H3CCOOH Acetic acid 4.74](data:image/gif;base64,R0lGODlhAQABAIAAAAAAAP///yH5BAEAAAAALAAAAAABAAEAAAIBRAA7)