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Synoptic Chemistry
Revision Notes
1)
Trends
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2)
Strength of covalent bonds
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3)
Atomic radius and first ionisation energy depend on the attraction between the
nucleus and the outer shell electrons
Electronegativity depends on the attraction between the nucleus and the shared
pair of electrons in a covalent bond
Strength of metallic bonding depends on the attraction between the nucleus and
the delocalised electrons
Across a period, shielding stays the same and nuclear charge increases. There is
increased attraction between the nucleus and the outer shell/shared
pair/delocalised electrons
Across a period, atomic radius decreases, first IE increases, electronegativity
increases
Down a group, the number of shells increases. This means that shielding
increases and the outer shell/shared pair/delocalised electrons is further from the
nucleus. Attraction between the nucleus and the outer shell/shared
pair/delocalised electrons decreases
Down a group, atomic radius increases, first IE decreases, electronegativity
decreases
Down groups 1 and 2, metallic bonding gets weaker
A covalent bond is a shared pair of electrons
Covalent bonds hold atoms together because both nuclei are attracted to the
shared pair of electrons
The strength of the bond depends on the strength of attraction between the
nuclei and the shared pair
Down a group attraction for the shared pair will decrease (see section 1) so bond
strength will decrease e.g. H-F 562 kJ mol-1, H-Cl 431, H-Br 366, H-I 299
Same trend and explanation for Cl-Cl, Br-Br and I-I (F-F is an exception)
Note the following definitions:
Bond enthalpy – enthalpy needed to break a covalent bond (applies to a bond
that only occurs in 1 molecule e.g. H-H)
Mean bond enthalpy – enthalpy needed to break a covalent bond averaged
over many compounds (applies to a bond that occurs in many compounds e.g.
C-H)
In calculations, mean bond enthalpies give approximate values for enthalpy
changes because they are average values (not specific to the compound being
used)
Bonding and electronegativity
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Electronegativity values can be used to predict whether a compound will have
ionic or covalent bonding
A small difference in electronegativity means that neither atom can pull a shared
pair of electrons completely to itself so the electrons will be shared and the
bonding will be covalent
A large difference in electronegativity means that one atom can pull a pair of
electrons completely to itself so ions will be formed and the bonding will be ionic
4)
Ionic Equations
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Ionic equations leave out ions that are unchanged in a reaction giving a clearer
picture of what is happening in the reaction
Ions can be cancelled if they are (aq) and occur on both sides of the equation
e.g.
Cl2(aq) + 2KBr(aq)  Br2(aq) + 2KCl(aq)
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Cl2 and Br2 are covalent, KBr and KCl are ionic
Aqueous K+ ions can be cancelled leaving:
Cl2(aq) + 2Br-(aq)  Br2(aq) + 2Cl-(aq)
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In ionic equations all acids are represented
Nitric acid and sodium hydroxide:
Sulphuric acid and barium hydroxide:
Hydrochloric acid and magnesium:
Hydrochloric acid and sodium carbonate:
Hydrochloric acid and magnesium oxide:
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In precipitation equations, react the ions needed to form the precipitate
Silver chloride:
Ag+ + Cl-  AgCl
Barium sulphate:
Ba2+ + SO42-  BaSO4
by H+ and all hydroxides by OHH+ + OH-  H2O
H+ + OH-  H2O
2H+ + Mg  Mg2+ + H2
2H+ + CO32-  CO2 + H2O
2H+ + MgO  Mg2+ + H2O