Group 1 +2

avatar
Created by
samantha macasa

Cards (33)

  • all elements in Groups 1 (alkali metals) = 1 electron in outermost principal quantum shell
    • all elements in Groups 2 (alkali earth metals) = 2 electrons in outermost principal quantum shell
  • all Group 1 + Group 2 metals can form ionic compounds via donating outermost electrons (act as reducing agents)become an ion with either a +1 or +2 charge (so they themselves become oxidised) 
  • Group 1 element reacts = need to lose 1 electron = only 1 electron in outer shell = 1+ ions formed = becomes outermost shell = already full + obtains noble gas configuration

    down group 1, number of shells of electrons increases by 1
    • outermost electron gets further away from nucleus = weaker forces of attraction between outermost electron + nucleus
    • less energy is required to overcome force of attraction as it gets weaker = outer electron is lost easily
  • Group 2 = outermost electron gets further away from nucleus = weaker forces of attraction between the outermost electron + nucleus
    • less energy  required to overcome force of attraction = outer electron is lost more easily
    • can be observed when Group 2 metals react with water
    • magnesium reacts extremely slowly with cold water
    • calcium reacts fairly vigorously with cold water in an exothermic reaction
  • group 1 metal reactions with oxygen
    • alkali metals react with oxygen in the air forming metal oxides = hence why alkali metals tarnish when exposed to air
    • metal oxide produce dull coating = covers surface of metal = metals tarnish more rapidly as down the group
  • Reaction with oxygen (Group 1)
    lithium + oxygen -> lithium oxide
    • 4Li(s) + O2 (g) -> 2Li2O(s)
    sodium + oxygen -> sodium oxide
    • 4Na(s) + O2 (g) -> 2Na2O(s)
    potassium + oxygen -> potassium superoxide
    • K(s) + O2(g) -> KO2(s)
  • Reactions with water (Group 1)
    • lithium + water -> lithium hydroxide + hydrogen
    • 2Li(s) + 2H2O(l) -> 2LiOH(aq) + H2(g)
    • slow reaction, lithium doesn't melt + fizzing seen

    • sodium + water -> sodium hydroxide + hydrogen
    • 2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)
    • heat released causing sodium to melt, hydrogen released catches on fire + sodium slides across surface

    • potassium + water -> lithium hydroxide + hydrogen
    • 2K(s) + 2H2O(l) -> 2KOH(aq) + H2(g)
    • heat released causing hydrogen with lilac flame, potassium melts into ball + slides across surface
  • Reactions with water (Group 2)
    • Mg(s) + H2O(l) -> MgO(s) + H2(g) = reacts with steam
    • Mg(s) + 2H2O(l) -> Mg(OH)2 (aq) + H2(g) = reacts with cold water
    • Ca(s) + 2H2O(l) -> Ca(OH)2 (aq) + H2(g)
    • Sr(s) + 2H2O(l) -> Sr(OH)2 (s) + H2(g)
    • Ba(s) + 2H2O(l) -> Ba(OH)2(s) + H2(g)
  • Reactions with chlorine (Group 2)
    • Group 2 metals + chlorine gas -> metal chloride
    • metal oxide + dilute hydrochloric acid -> metal chloride + water
  • Reactions with sulfuric acid (Group 2)
    • metal oxide + dilute sulfuric acid -> metal sulfate + water
  • Reactions with hydroxides
    • metal hydroxide + dilute hydrochloric acid -> metal chloride + water
    • metal hydroxide + dilute sulfuric acid -> metal sulfate + water
  • Reactions of Group 1 oxides with water
    • Group 2 metal oxides react with water = give colourless alkaline solution (i.e. Na2O (s) + H2O (l) → 2NaOH (aq))
    • solution is alkaline = hydroxide ions released = O2- (s) +  H2O (l) → 2OH- (aq) 
  • Reactions of Group 1 hydroxides with dilute acid
    • Group 1 metal hydroxide acts as an alkali when added to dilute acid  
    • alkali reacts with an acid (neutralisation reaction occurs) = forms salt + water

    NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l) 
    NaOH (aq) + H2SO4 (aq) → Na2SO4 (aq) + H2O (l) 
  • Reactions of Group 2 oxides with water
    • all Group 2 oxides are basic (except BeO's amphoteric acting both an acid + base)
    • Group 2 oxides react water to form alkaline solutions =more alkaline going down group
    • MgO(s) + H2O(l) -> Mg(OH)2(aq) = slightly soluble in water + weakly alkaline solution (pH 10)
    • CaO(s) + H2O(l) -> Ca(OH)2(aq) = releases heat energy causing water to boil + CaO solid expands (pH 11)
    • SrO(s) + H2O(l) -> Sr(OH)2(aq)
    • BaO(s) + H2O(l) -> Ba(OH)2(aq)
  • Reaction of Group 2 oxides with acid
    • Group 2 sulfates form when Group 2 oxide reacts with sulfuric acid
    • insoluble sulfates form at surface of oxide = solid oxide beneath it cannot react with acid
    • prevented via using oxide in powder form + stirring = neutralisation occurs
  • Reaction of Group 2 hydroxides with dilute acid
    • Group 2 metal hydroxides form colourless solutions of metal salts when react with dilute acid
    • sulfates decrease in solubility down the group (barium sulfate is insoluble white precipitate)
  • Reaction of Group 2 hydroxides with dilute hydrochloric acid
    • Mg(OH)2(s) + 2HCl(aq) -> MgCl2(aq) + 2H2O(l)
    • Ca(OH)2(s) + 2HCl(aq) -> CaCl2(aq) + 2H2O(l)
    • Sr(OH)2(s) + 2HCl(aq) -> SrCl2(aq) + 2H2O(l)
    • Ba(OH)2(s) + 2HCl(aq) -> BaCl2(aq) + 2H2O(l)
  • Reaction of Group 2 hydroxides with dilute sulfuric acid
    • Mg(OH)2(s) + H2SO4(aq) -> MgSO4(aq) + 2H2O(l)
    • Ca(OH)2(s) + H2SO4(aq) -> CaSO4(s) + 2H2O(l)
    • Sr(OH)2(s) + H2SO4(aq) -> SrSO4(s) + 2H2O(l)
    • Ba(OH)2(s) + H2SO4(aq) -> BaSO4(s) + 2H2O(l)
  • Group 2 hydroxides solubility trend
    • down the group, solutions formed from reaction of Group 2 oxides with water become more alkaline
    • when oxides dissolved in water = following ionic reaction occurs: O2- (aq) + H2O(l) → 2OH- (aq)
    • higher concentration of OH- ions = more alkaline solution is
    • alkalinity of formed solution = therefore explained by solubility of the Group 2 hydroxides
  • Group 2 sulfates solubility trend
    • solubility of Group 2 sulfates decreases down the group
  • thermal decomposition: breakdown of compound into two or more different substances via heat
  • Thermal decomposition of carbonates (Group 1 + 2)
    • Group 1 = lithium carbonate heated will decompose producing lithium oxide + carbon dioxide = Li2CO3 (s)   Li2O (s)  +  CO2 (g)
    • rest of Group 1 carbonates don't decompose at Bunsen temperatures
    • decomposition temperatures increase as down the Group

    • Group 2 carbonates break down when heated = forms metal oxide + carbon dioxide = XCO3 (s)  XO (s) + CO2 (g) 
    • down the group, more heat needed to break down carbonates
  • Thermal decomposition of nitrates (Group 1)
    • only Group 1 nitrate that decomposes to produce nitrogen dioxide (brown toxic gas) + oxygen is lithium nitrate (LiNO)
    4LiNO3 (s)  2Li2O (s) + 4NO2 (g) + O2 (g)
    • rest of Group 1 don't decompose = producing metal nitrite (NO2-) + oxygen + no nitrogen dioxide
    2XNO3 (s)  2XNO2 (s) + O2 (g)
    • all nitrates from sodium to caesium decompose in same way, only difference the temperature to undergo reaction
  • Theraml decomposition of nitrates (Group 2)
    • Group 2 nitrates decompose when heated to form = metal oxide, oxygen gas + nitrogen dioxide gas 
    • nitrogen dioxide gas is toxic = often carried out in fume cupboard
  • Trend in thermal decomposition of carbonates/nitrates (Group 1 + 2)
    • down Groups 1 + 2 = more heat needed to break down carbonate + nitrate ions
    • thermal stability of Group 1/2 carbonates + nitrates = increases down group 
    • smaller cations will polarise larger anions more = cations attract delocalised electrons in nitrate/carbonate ion
    • higher charge + smaller ion = higher polarising power
    • more polarised = likely to be thermally decompose = bonds in carbonate/nitrate ions become weaker 
  • Testing for gases (carbon dioxide)
    • bubbled through limewater = carbon dioxide gas turn limewater milky (white precipitate of calcium carbonate CaCO3)
    • Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O (l)
  • Testing for gases (oxygen)
    • oxygen gas's present = will relight glowing splint 
    • C (s) + O2 (g) → CO2 (g) 
  • Testing for gases (nitrogen dioxide)
    • NO2 is a toxic brown/orange gas = dissolved in water = give acidic solution
    • NO2 (g) + H2O (l) → 2HNO3 (aq) + HNO2 (aq) 
  • Flame colours characteristics
    • metal ions produce a colour if heated in a flame = ions from different metals produce different colours
    • flame test is used to identify metal ions via colour of the flame

    • the heat causes electrons to move to higher energy levels = electron is unstable at this energy level so falls down 
    • as it falls down down from higher to lower energy level = energy emitted into visible light energy with wavelength of observed light
  • metal ions + colours observed
    1. Li+ = scarlet red
    2. Na+ = yellow
    3. K+ = lilac
    4. Rb+ = red
    5. Cs+ = blue
    6. Mg2+ = no flame colour (emitted outside visible spectrum)
    7. Ca2+ = brick red
    8. Sr2+ = red
    9. Ba2+ = apple green
  • Testing for ammonium ions
    1. add 10 drops of of solution with ammonium ions to clean test tube
    2. add 10 drops of sodium hydroxide via pipette
    3. test tube swirled carefully + place into beaker of water
    4. beaker of water placed above bunsen burner = heat gently causing fumes to be produced
    5. using a pair of tongs place damp red litmus paper near test tube = should change to blue in presence of ammonia gas
    • Equation: NH4Cl(aq) + NaOH(aq) → NH3(g) + H2O(l) + NaCl(aq)
    • Ionic equation: NH4+(aq) + OH-(aq) → NH3(g) + H2O(l)
  • Testing for carbonate ions
    1. add small amount (1cm3) of dilute hydrochloric acid to test tube via pipette
    2. add equal amount of sodium carbonate solutoin to test tube via pippete
    3. add bung with delivery tube to the test tube = transferring gas produced to test tube containing limewater
    4. carbonate ions react with hydrogen ions = producing carbon dioxide gas = causing limewater milky
    • Equation: 2HCl(aq) + Na2CO3(aq) → 2NaCl(aq) + CO2(g) + H2O(l)
    • Ionic equation: 2H+(aq) + CO32-(aq) → CO2(g) + H2O(l)
  • Testing for sulfate ions
    1. acidify sample via dilute hydrochloric acid + add few drops of aqueous barium chloride
    2. sulfate is present = white precipitate of barium sulfate formed:
    • Ba2+ (aq) + SO42- (aq) → BaSO4 (s)