Chemistry 2.1
Isotopes, atoms of the same element with different numbers of neutrons and
different masses/number of protons.
Ions, a charged atom or molecule.
History of atomic structure:
Democritus (BC) – Particles ‘atomos’, sample of matter that is indivisible.
Dalton (early 1800s) – Atomic theory, tiny particles make an element, no divisible, all atom of given element are the
same, atom of one element is different from atom of another element.
J.J Thompson (1897 -1906) – Plum Pudding, discover electron, can be deflected by magnet and electric field, very
small mass. Disproved that atoms cannot be split. +charge and – charge balance.
Rutherford (1909-11) – Gold leaf experiment. Fired a-particles in thin sheet of gold, expected all/most particles to go
straight through – yet found – most went straight through, small % were deflected through large angle, some
deflected straight back.
Calculations made him make a nuclear model -
+ charge is concentrated on centre
- charge orbit atom like planet
Bohr (1913) – Planetary model, explain periodic properties, spectral line in emission spectra, electrons at different
distances from the nucleus.
Moseley (1913) – Found link between x-ray frequency and element atomic number.
Rutherford (1918) – Discovered proton
Broglie and Schrodinger (1923-26) – particle could have nature of wave and particle, and intro to idea of atomic
orbitals.
Chadwick (1932) – Discovered neutrons, radiation of uncharged particles with same mass as proton (approx.)
Relative Atomic Mass: weighed mean mass of an atom compared to 1/12th the mass of an atom of C-12.
Relative Isotopic Mass: mass of an atom of an isotope compared to 1/12th the mass of an atom of C-12.
(Relative Molecular Mass: Mean mass of a molecule, compared to 1/12th of the mass of an atom of C-12)
Mass Spectrometry:
Ionisation, by firing electrons.
Acceleration, of charges ions.
Deflection, high mass means less deflection (vice versa)
Detection, of the mass to charge ratio. (m/z)
For simple molecules, the term relative molecular mass will be used.
For compounds with giant structures, the term relative formula mass will be used
Zn2+ is zinc ions. Mole = Amount of any substance containing as many particles
Ag+ is Silver ion as there are carbon atoms in exactly 12g of carbon-12 isotope.
Ionic Formula Steps – 1 mol of any atom/ molecule = 6.02 x 10^23. Avogadro’s Number
Swap and Drop Method Number of particle = Avogadro’s Number x Number of Moles
Mole Equations:
Mole (Solid) = Mass/ Mr
Mole (Gas) = Volume (dm^3) / 24
Mole (Solution)= Concentration (moldm^-3) x Volume (dm^3)
(Divide cm^3 by 1000 to get dm^3)
Empirical Formula – simplest whole number ratio of atoms
Molecular Formula - the number and type of atoms of each element in a molecule.
(always find mole to be able to calculate)
, Water of Crystallisation: Water molecules that form part of the crystalline structure of a compound.
Hydrated: A crystalline compound that contains water.
Anhydrous: A crystalline compound containing no water.
(always find mole to be able to calculate)
Stoichiometry: the relative quantities of substances in a reaction.
A high atom economy means a process is more
sustainable as there is less waste produced.
Addition reaction = 100% econ
Substitution and elimination = less than 100% econ
Acids release H+ ions in aqueous solution. (PROTON DONORS)
Common acids: HCl, H2SO4, HNO3 and CH3COOH
Strong Acid: An acid that completely dissociates in solution.
Weak Acid: an acid that only partially dissociates in solution.
Alkalis release OH– ions in aqueous solution. (PROTON ACCEPTORS)
Common alkalis: NaOH, KOH and NH3
A base is a something that react with an acid to form water and a salt. An alkali is any base that is soluble in water.
Salt = The H+ ion in an acid has been replaced by a metal ion
Neutralisation Reaction: H+ + OH– = H2O
Acid + Base = Salt + Water (fizz, solid dissolve)
Acid + Metal = Salt + Hydrogen (fizz ,solid dissolve)
Acid + Metal carbonate = Salt + Water + CO2 (fizzing, solid dissolve)
Acid + Metal hydroxide = Salt + Water (fizz, solid dissolve)
Standard Solution (e.g. NaHCO3) :
1) Carefully weigh out the required mass of the solute to two decimal places using a balance.
2) Transfer the solute to a beaker and dissolve it in a small amount of distilled water.
3) Use a funnel to transfer the solution to a volumetric flask. Rinse the beaker with the distilled water and add the
washings to the flask.
4) Add some distilled to the flask, but don't fill it to the graduation line. Then, slowly add more solvent drop by drop until
the bottom of the meniscus reaches the line.
5) Stopper and shake the flask.
Acid Base Titration:
1) Use pipette to add 25cm^3 of acid (known concentration) to a conical flask. Add a few drops of indicator.
2) Pour Alkali (unknown concentration) into the Burette . Record the initial burette volume.
3) Complete a trial titre by slowly adding alkali to acid. The conical flask should be swirling constantly about a white tile.
Stop adding the alkali once end point in reached.
Record the final burette volume.
4) Repeat the titration until 2 concordant results are obtained. Add alkali drop by drop near the end point.
Practice lots of Titration Questions (find moles if any doubts)
,Redox:
Oxidation: loss of electrons/ increase in oxidation number.
Reduction: gain of electrons / decrease in oxidation number.
Redox Reaction: A reaction in which one element is oxidised
and another is reduced.
Chemistry 2.2
Each shell can hold 2n^2 electrons.
Atomic orbitals:
A region around the nucleus that can hold up to two electrons, with opposite spins.
S- orbitals – has 1 orbital can hold 2 electrons: Orbitals – 3 orbitals can hold 3x2 = 6 electrons
Spherical Dumbbell
Subshells Number of subshells
s 1
p 3
d 5
f 7
For orbitals with the same energy, occupation singly before pairing
, Ionic Bonding:
Electrostatic attraction between positive and negative ions.
Dot and cross diagram of CaO.
Solid structures of giant ionic lattices -
Oppositely charged ions strongly attracted in all directions
(electrostatic attraction).
Ionic compounds –
Compounds made up of oppositely charged ions. Large amount of energy is needed to break the strong electrostatic
attraction between the oppositely charged ions.
BP and MP = High
Soluble = Yes, in polar solvents . Ionic lattice is polar, so in water (also a polar solvent), ionic lattice breaks down and
s+ attracts to O-, s- attracts to H+.
Conducting electricity = Yes when liquid, not when solid. Because ions are held in fixed position when solid yet are
free to move when liquid – so are able to carry charge.
Macroscopic Properties: properties of a bulk material rather than the individual atoms/ molecules that
make up the material.
At the macroscopic scale, ionic compounds form lattices, are crystalline solids under normal conditions, and have high melting
points. Most of these solids are soluble in H2O and conduct electricity when dissolved.
Covalent Bonds -
Covalent bond as the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded
atoms.
Types of covalent bonds:
Dative Bond: a type of covalent bond in which both of the electrons in the shared pair come from one atom.
Average bond enthalpy: the average energy
required to break a bond
The larger the value of the average bond
enthalpy, the stronger the covalent bond.
90
Isotopes, atoms of the same element with different numbers of neutrons and
different masses/number of protons.
Ions, a charged atom or molecule.
History of atomic structure:
Democritus (BC) – Particles ‘atomos’, sample of matter that is indivisible.
Dalton (early 1800s) – Atomic theory, tiny particles make an element, no divisible, all atom of given element are the
same, atom of one element is different from atom of another element.
J.J Thompson (1897 -1906) – Plum Pudding, discover electron, can be deflected by magnet and electric field, very
small mass. Disproved that atoms cannot be split. +charge and – charge balance.
Rutherford (1909-11) – Gold leaf experiment. Fired a-particles in thin sheet of gold, expected all/most particles to go
straight through – yet found – most went straight through, small % were deflected through large angle, some
deflected straight back.
Calculations made him make a nuclear model -
+ charge is concentrated on centre
- charge orbit atom like planet
Bohr (1913) – Planetary model, explain periodic properties, spectral line in emission spectra, electrons at different
distances from the nucleus.
Moseley (1913) – Found link between x-ray frequency and element atomic number.
Rutherford (1918) – Discovered proton
Broglie and Schrodinger (1923-26) – particle could have nature of wave and particle, and intro to idea of atomic
orbitals.
Chadwick (1932) – Discovered neutrons, radiation of uncharged particles with same mass as proton (approx.)
Relative Atomic Mass: weighed mean mass of an atom compared to 1/12th the mass of an atom of C-12.
Relative Isotopic Mass: mass of an atom of an isotope compared to 1/12th the mass of an atom of C-12.
(Relative Molecular Mass: Mean mass of a molecule, compared to 1/12th of the mass of an atom of C-12)
Mass Spectrometry:
Ionisation, by firing electrons.
Acceleration, of charges ions.
Deflection, high mass means less deflection (vice versa)
Detection, of the mass to charge ratio. (m/z)
For simple molecules, the term relative molecular mass will be used.
For compounds with giant structures, the term relative formula mass will be used
Zn2+ is zinc ions. Mole = Amount of any substance containing as many particles
Ag+ is Silver ion as there are carbon atoms in exactly 12g of carbon-12 isotope.
Ionic Formula Steps – 1 mol of any atom/ molecule = 6.02 x 10^23. Avogadro’s Number
Swap and Drop Method Number of particle = Avogadro’s Number x Number of Moles
Mole Equations:
Mole (Solid) = Mass/ Mr
Mole (Gas) = Volume (dm^3) / 24
Mole (Solution)= Concentration (moldm^-3) x Volume (dm^3)
(Divide cm^3 by 1000 to get dm^3)
Empirical Formula – simplest whole number ratio of atoms
Molecular Formula - the number and type of atoms of each element in a molecule.
(always find mole to be able to calculate)
, Water of Crystallisation: Water molecules that form part of the crystalline structure of a compound.
Hydrated: A crystalline compound that contains water.
Anhydrous: A crystalline compound containing no water.
(always find mole to be able to calculate)
Stoichiometry: the relative quantities of substances in a reaction.
A high atom economy means a process is more
sustainable as there is less waste produced.
Addition reaction = 100% econ
Substitution and elimination = less than 100% econ
Acids release H+ ions in aqueous solution. (PROTON DONORS)
Common acids: HCl, H2SO4, HNO3 and CH3COOH
Strong Acid: An acid that completely dissociates in solution.
Weak Acid: an acid that only partially dissociates in solution.
Alkalis release OH– ions in aqueous solution. (PROTON ACCEPTORS)
Common alkalis: NaOH, KOH and NH3
A base is a something that react with an acid to form water and a salt. An alkali is any base that is soluble in water.
Salt = The H+ ion in an acid has been replaced by a metal ion
Neutralisation Reaction: H+ + OH– = H2O
Acid + Base = Salt + Water (fizz, solid dissolve)
Acid + Metal = Salt + Hydrogen (fizz ,solid dissolve)
Acid + Metal carbonate = Salt + Water + CO2 (fizzing, solid dissolve)
Acid + Metal hydroxide = Salt + Water (fizz, solid dissolve)
Standard Solution (e.g. NaHCO3) :
1) Carefully weigh out the required mass of the solute to two decimal places using a balance.
2) Transfer the solute to a beaker and dissolve it in a small amount of distilled water.
3) Use a funnel to transfer the solution to a volumetric flask. Rinse the beaker with the distilled water and add the
washings to the flask.
4) Add some distilled to the flask, but don't fill it to the graduation line. Then, slowly add more solvent drop by drop until
the bottom of the meniscus reaches the line.
5) Stopper and shake the flask.
Acid Base Titration:
1) Use pipette to add 25cm^3 of acid (known concentration) to a conical flask. Add a few drops of indicator.
2) Pour Alkali (unknown concentration) into the Burette . Record the initial burette volume.
3) Complete a trial titre by slowly adding alkali to acid. The conical flask should be swirling constantly about a white tile.
Stop adding the alkali once end point in reached.
Record the final burette volume.
4) Repeat the titration until 2 concordant results are obtained. Add alkali drop by drop near the end point.
Practice lots of Titration Questions (find moles if any doubts)
,Redox:
Oxidation: loss of electrons/ increase in oxidation number.
Reduction: gain of electrons / decrease in oxidation number.
Redox Reaction: A reaction in which one element is oxidised
and another is reduced.
Chemistry 2.2
Each shell can hold 2n^2 electrons.
Atomic orbitals:
A region around the nucleus that can hold up to two electrons, with opposite spins.
S- orbitals – has 1 orbital can hold 2 electrons: Orbitals – 3 orbitals can hold 3x2 = 6 electrons
Spherical Dumbbell
Subshells Number of subshells
s 1
p 3
d 5
f 7
For orbitals with the same energy, occupation singly before pairing
, Ionic Bonding:
Electrostatic attraction between positive and negative ions.
Dot and cross diagram of CaO.
Solid structures of giant ionic lattices -
Oppositely charged ions strongly attracted in all directions
(electrostatic attraction).
Ionic compounds –
Compounds made up of oppositely charged ions. Large amount of energy is needed to break the strong electrostatic
attraction between the oppositely charged ions.
BP and MP = High
Soluble = Yes, in polar solvents . Ionic lattice is polar, so in water (also a polar solvent), ionic lattice breaks down and
s+ attracts to O-, s- attracts to H+.
Conducting electricity = Yes when liquid, not when solid. Because ions are held in fixed position when solid yet are
free to move when liquid – so are able to carry charge.
Macroscopic Properties: properties of a bulk material rather than the individual atoms/ molecules that
make up the material.
At the macroscopic scale, ionic compounds form lattices, are crystalline solids under normal conditions, and have high melting
points. Most of these solids are soluble in H2O and conduct electricity when dissolved.
Covalent Bonds -
Covalent bond as the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded
atoms.
Types of covalent bonds:
Dative Bond: a type of covalent bond in which both of the electrons in the shared pair come from one atom.
Average bond enthalpy: the average energy
required to break a bond
The larger the value of the average bond
enthalpy, the stronger the covalent bond.
90
