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Synoptic Chemistry Revision Notes 1) Trends 2) Strength of covalent bonds 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 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 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) 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) 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: 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