Topic 4 : Inorganic Chem & Pt Flashcards
Ion Polarisation
Distortion of the electron could away from completely spherical
Cations (+) have polarising power
Smaller + greater charge mean more polarisation (greater charge density, nc is more effective)
Larger anion: easier polarisation (outer shells held less tightly, more easily distorted)
Ionisation energy and Reactivity trends down Group 2
Decreases down the group
Extra e- and shells cause more shielding : outer are further away from core
Electrostatic attraction decreases, easier to remove an electron so IE decreases while reactivity increases
Group 2 Element: basic reactions
Water -> hydroxides (OH)2 BASE
Oxygen -> oxides 1:1 BASE
Chlorine -> chlorides Cl2
***except Be(OH)2 which is insoluble and BeO, MgO don’t act as bases
Solubility of Group 2 compounds
Depends on the compound anion
Hydroxides increase down
Sulfate decreases (e.g BaSO4 is insoluble)
Thermal Stability of CO3s and NO3s (Group 1 & group 2 compounds)
Increases down a group
Large anions are distorted (greater polarisation, easier thermal decomp)
Larger cations have lower charge density and so cause less distortion-> Ba is more stable than Mg
Comparison G1&G2: Thermally stability
Group 1 are more thermally stable as they are less polarising because of lower charge density
G1 carbonates do not decompose with Bunsen burners (except Li2CO3 -> Li2O and CO2)
G1 nitrates decompose to nitrite (NO2) and oxygen (except LiNO3 -> Li2O + NO2 + O2)
G2 carbonates form oxide and carbon dioxide
G2 nitrates form oxide, nitrogen dioxide and oxygen
TEST: how long it takes for brown gas NO2 to be produced / test for CO2 using lime water / test for O2
Flame colours of Group 1 and 2
Li: red
Na: orange/yellow
K: lilac
Rb: red
Cs: blue
Ca: brick red
Sr: crimson
Ba: green
Flame Test: procedure
Clean nichrome/platinum by dipping it in HCl and burning up it using a Bunsen burner until flame has no colour
Wet wire with acid and dip into solid (powder metal)
Place wire on flame & record colour
Halogens Theory
Reactive non-metals (ns2, np5),
poisonous, corrosive
colour gets darker
Uses: bleach, disinfectant
Group 2: general theory descending
Nuclear charge increases but repulsion between shells & n of shells goes up = ion size up (overall nc less effective)
Ion size up = metallic bonding strength down (less E needed to overcome e-static attraction between e- coils & +ve ions)
Charge density decreases
First Ionisation energy decreases
Halogens: mps
Molecules larger so London forces increase
(Greater n of shells increasingly away from nucleus)
Halogens: electronegativity
Decreases down group as atomic radius increases
(Attraction drops off as distance increases)
Halogens & water
Disproportionation reaction: simultaneous oxidation & reduction of a species (reversible)
E.g Cl2 + H2O -> HCl + HOCL (hydrochlorous acid)
Cl: 0 -1 +1
Halogens & alkalis
Cold, aqueous NaOH: salt, water & metal hypochlorite (OX- where X=halogen and has os = +1)
Hot, concentrated NaOH: salt, water & metal chlorate (XO3- where X has os = +5)
Fluoride/Chloride ions & sulfuric acid:
Weak: act as oxidising agents (are reduced)
very endothermic reaction
X- + H2SO4 -> HX (g) + HSO4-
HX: white misty fumes (hydrochloric acid droplets form)
Bromide ions and sulfuric acid
Reducing agents
2HBr + 2H2SO4 -> Br2 + SO2 + 2H2O
Iodide ions & sulfuric acid
I- is the strongest reducing agent
KI + H2SO4 -> KHSO4 + HI (g) (white misty fumes)
HI + H2SO4 -> I2 + SO2 + 2H2O (Same as Br but then keeps going)
6HI + SO2 -> 3I2 + H2S + 4H2O
Observations: I2 purple fumes, H2S rotten egg smell
Halides: Precipitacions reaction
Two ionic solutions are mixed to give an insoluble compound (used to identify ions)
Metal halide + acidified silver nitrate
Cl, Br, I ions form precipitates of specific colours
To distinguish colours: add dilute ammonia, if it doesn’t dissolve add concentrated ammonia
Silver nitrate + halide precipitate colours
AgCl: white
AgBr: cream
AgI: yellow
Acidified silver nitrate
Acidified to remove ions that could give a false positive result (e.g OH-, CO3-…)
Acidified with HNO3 so that anion from acid does not form a precipitate
