Some of Unit 1 As chemistry for OCR

1. Acids & Salts, Redox, Shells and Structure, Bonding
(Physical properties) and Periodicity

2. Acids and Salts
Molecular formula Name
CH3COOH Ethanoic acid
HCOOH Methanoic acid
HNO3 Nitric acid
HCL Hydrochloric acid
C6H8O7 Citric acid
H2SO4 Sulphuric acid
H3PO4 Phosphoric acid

3. Reactions of acids and Redox
 Acid + carbonate salt +CO 2 H2 O
 Acid + Base salt + H2 O
 Acid + Alkali salt + H2 O
 Acid + Metal salt + H2
OIL RIG:
AN ACID IS A PROTON
DONOR
A BASE IS A PROTON
ACCEPTOR
Oxidation Is Loss,
Reduction Is Gain
(of electrons.)

4. Reactions of acids and Redox
 Acids release H+ ions when they are in aqueous solutions.
 The common bases are metal oxides, metal hydroxides and ammonia.
 An alkali is a soluble base that releases OH– ions in aqueous solutions.
 Metals generally form ions by losing electrons with an increase in
oxidation number to form positive ions.
 Non-metals generally react by gaining electrons with a decrease in
oxidation number to form negative ions.

5. Reactions of acids and Redox
Name Base
Metal Oxides MgO and CuO
Metal Hydroxides NaOH and Mg (OH)2
Ammonia NH 3
Amine CH 3 NH 2
Magnesium
Hydroxide
Mg (OH)2
Calcium Hydroxide Ca(OH)2

6. Key Definitions
Definitions:
 Hydrated
Refers to a crystalline compound containing water molecules
 Anhydrous
Refers to a substance that contains no water molecules
 Water of crystallisation
Refers to water molecules that form and essential part of the
crystalline structure of a compound

7. Shells, atoms, bonds and structure
First ionisation energy:
The amount of energy
required to remove one
electron from each atom
in one mole of gaseous
atoms, to form one mole
of gaseous 1+ ions.
Electrons are arranged in sub shells (s-, p-
and d-sub-shells.) One orbital is a region
that can hold up to two electrons of opposite
spins. For example: Aluminium. The
electrons always fill a subshell by them
selves first, rather like finding a seat on a
bus; you’re not going to sit next to someone
you don’t know.
Successive ionisation energy for Sodium:
Ionisation energy is all about
removing electrons. The equation
is always the same, just with a
change in element obviously. The
definition is key, and is worth 3
marks in exam papers.

8. Shells, atoms, bonds and structure
 Factors effecting ionisation energy’s:
 Atomic radius: Greater atomic radius = smaller nuclear
attraction on outer electrons.
 Nuclear charge: Greater nuclear charge = greater attractive
force on the outer electrons.
 Electron shielding: This is when the inner shells repel the
outer shells of electrons. More inner shells = smaller
nuclear attraction on the outer electrons.

9. Bonding: Ionic, Covalent and
Metallic
 Ionic bonding:
This is the electrostatic attraction
between oppositely charged ions.
 Covalent bonding:
This can be described as a shared pair of electrons.
A dative covalent bond is where a pair of electrons
is provided by one of the atoms.
 Metallic bonding:
This is the attraction of positive
ions to a sea of delocalised electrons. Sea of delocalised electrons

10. Bonding: The forces and attractions
Electronegativity and polarity:
 Non-polar bonds: In a molecule of H2 the bond is identical, as both atoms
have a fair share of electrons (100% perfect covalent bond). This results in
a non-polar molecule. This is also true for Cl2 and other diatomic atoms.
 Polar bonds: In a molecule of HCL, one atom is more electronegative than
the other; in this case, it is the Cl. This means it has a greater attraction
for the bonding pair, which draws the electrons closer to it, creating a
permanent dipole. This results in a polar bond.
 Polar molecules: HCL is a polar molecule with polar bonds because of the
charge difference across the whole molecule. This is also a non-
symmetrical molecule.
 Electronegativity: This increases across the periods and decreases down
the groups. If there is a small difference in the electronegativity, it will
result in a polar covalent bond. If there is a big difference in
electronegativity, one atom will effectively capture the electrons and
result in forming an ionic bond.

11. Bonding: The forces and attractions
Intermolecular forces:
 An intermolecular force: This is the attractive force between neighbouring
molecules.
 Permanent dipoles: A permanent dipole-dipole force is a weak attractive force
between permanent dipoles in neighbouring polar molecules.
 Van der Walls’: These are an attractive force between induced dipoles in
neighbouring molecules. These attractions are caused by electron ‘wobbling ’ within
the molecule. This induces a charge in neighbouring molecules as they attract to
each other. These forces increase with increasing number of electrons, which
increase the attractive forces between the molecules.
 Hydrogen bonding: This temporary attraction between the slightly negative oxygen
and the slightly positive hydrogen in a water molecule. The bond runs from one of
the two lone pairs on the oxygen to the hydrogen. This gives water its special
properties of high surface tension and ice being less dense than water. This is
because the rigid bonds in the ice collapse when it melts, allowing the water
molecules to move closer together.

12. Bonding and physical properties
Structure Type of bonding Properties Malleable or ductile?
Soluble or insoluble?
Simple molecular Within in the structure
there are strong
covalent bonds, but the
molecules are held
together by weak van
der Waal’s forces.
 They have low melting and boiling points.
 Weak forces don’t need a lot of energy to be
broken.
 They don’t conduct electricity as there are
no free ions to move in the structure.
They are soluble in non-polar
solvents as weak van der Waal’s
forces form between the solvent
and simple molecular structure.
This formation of bonds weakens
the lattice.
Giant covalent Within the structure
there are strong
covalent bonds.
 High melting and boiling point.
 High temperatures needed to break the
strong covalent bonds.
 Non-conductors of electricity due to no free
ions.
Giant covalent structures are
soluble in both polar and non-
polar solvents because the
covalent bonds are too strong to
be broken.
Giant ionic Every ion in a giant
ionic lattice is
surrounded by an ion of
the opposite charge.
 Strong electrostatic forces which need
energy to break.
 When a solid, it is a non-conductor of
electricity.
 When molten or dissolved in in water, the
ions can move about and conduct electricity.
Giant ionic lattices can dissolve
in polar solvents such as water.
Giant metallic In a giant metallic
lattice, positive ions
hold a fixed position
while the delocalised
electrons are free to
move throughout the
structure.
 High melting points and boiling points
because the attraction between the positive
ions and negative electrons is strong and
needs high temperatures to break the bonds.
 They also have good electric conductivity
due to the free moving electrons.
Giant metallic lattices are both
malleable and ductile, meaning it
can be stretched out into a wire
(ductile), or hammered into
shape (malleable).

13. Reasons for high or low melting and boiling
points in periods 2 and 3
Period 2 Li Be B C N2 O2 F2 Ne
Period 3 Na Mg Al Si P4 S8 Cl2 Ar
Structure Giant
metallic
Giant
covalent
Simple molecular
structures
Forces Strong forces
between ions
and
delocalised
electrons
Strong
forces
between
atoms
Weak forces between
molecules
Bonding Metallic
bonding
Covalent
bonding
Van der Waal’s forces

14. Periodicity and first ionisation
energy