Topic 4A: Elements of Groups 1 and 2

Cards (30)

    1. Understand the reasons for the trend in ionization energy down Group 2
    -The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells.
    -The outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons
  • 2. Understand reasons for the trend in reactivity of the Group 2 elements down the group
    -Reactivity of group 2 metals increases down the group as the atomic radii increases there is more shielding. The nuclear attraction decreases and it is easier to remove (outer) electrons and so cations form mor easily.
  • 3. know the reactions of the elements Mg to Ba in Group 2 with oxygen (Mg)
    Group 2 metals will burn in oxygen. Mg burns with a bright white flame. MgO is a white solid with a high mp due to its ionic bonding.
    2Mg+O2->2MgO Mg is oxidised (oxidation state increases from 0 to +2) and O2 is reduced, forming metal oxides. Mg will also react slowly with oxygen without a flame- Mg ribbon will have a thin layer of MgO on it formed by reaction with O2.
  • 3. . know the reactions of the elements Mg to Ba in Group 2 with oxygen
    The MgO must be cleaned off by emery paper before doing any reactions wih Mg ribbon. If testing for reaction rates with Mg and acid, an un-cleaned Mg and MgO would react but at different rates.
    Mg+2HCl->MgCl2+H2 MgO+2HCl->MgCl2+H2O. Barium is most reactive- stored under oil to keep it from reacting with O2 and H2O vapour in the air, Ba=green flame. No reaction with beryllium. Ca burns with red flame-CaO. Group 2 reacts vigorously with O2 if placed in a glass jar.
  • 3. know the reactions of the elements Mg to Ba in Group 2 with chlorine
    Group 2 reacts with chlorine to form metal chlorides. Mg+Cl2->MgCl2
    white solid. Ca(s)+Cl2(g)->CaCl2(s) Ca is oxidised, oxidation state increases from 0 to +2, Cl2 is reduced, oxidation state decreases from 0 to -1.
  • 3. know the reactions of the elements Mg to Ba in Group 2 with water(Mg)
    Mg will react in steam to produce MgO and H2, burning with a bright white flame. Mg(s)+H2O(g)->MgO(s)+H2(g) MgO=white solid. Mg will also react with warm water, giving Mg+2H2O->Mg(OH)2+H2. This is a much slower reaction than reaction with steam; no flame.
  • 3. know the reactions of the elements Mg to Ba in Group 2 with water
    The other group 2 metals will react with cold water with increasing vigour down the group to form hydroxides. Ca+2H2O(l)->Ca(OH)2(aq)+H2(g). Ca(OH)2, slightly soluble, cloudy precipitate. Sr+2H2O(l)->Sr(OH)2(aq)+H2(g). Ba+H2O(l)->Ba(OH)2(aq)+H2(g), Ba(OH)2 is soluble. Not as vigorous as for Group 1 metals.
  • 3. know the reactions of the elements Mg to Ba in Group 2 with water
    The hydroxides produced make the water alkaline.
    One would observe:
    -fizzing, (more vigorous down the group)
    -metal dissolving (faster down the group)
    -solution heating up (more down the group)
    -and with Ca a white precipitate appearing (less precipitate forms down group)
    Increasing in effervescence going down
  • 4. know the reactions of the oxides of Group 2 elements with water
    Group 2 ionic oxides (solid) react with water to form hydroxides(colourless solutions), (not beryllium oxide). Ionic oxides are basic as the oxide ions accept protons to become hydroxide ions in this reaction (acting as a bronsted lowry base) MgO(s)+H2O(l)->Mg(OH)2(s) pH9. Mg(OH)2 is only slightly soluble in water so fewer free OH- ions are produced and so lower pH. pH increases down the group. CaO(s)+H2O(l)->Ca(OH)2(s) pH12 vigorous reaction, limewater reacts with CO2 to form precipitate-CaCO3. simplifies to O2-+H2O->2OH-
  • 4. know the reactions of the oxides of Group 2 elements with dilute acid
    metal oxide+acid->salt+water. MgO(s)+2HCl(aq)->MgCl2(aq)+H2O(l)
    SrO(s)+2HCl(aq)->SrCl2(aq)+H2O(l) CaO(s)+2HCl(aq)-> CaCl2(aq) + H2O(l) neutralisation therefore exothermic. Metal oxides are white solids, Products=colourless solutions
  • 4. know the reactions of the hydroxides of Group 2 elements with dilute acid
    2HNO3(aq)+Mg(OH)2(aq)->Mg(NO3)2(aq)+2H2O(l)
    2HCl(aq)+Mg(OH)2(aq)->MgCl2(aq)+2H2O(l)
    Ca(OH)2(aq)+2HNO3(aq)->Ca(NO3)2(aq)+2H2O(l)
    lime is used to neutralise acidic soil
  • 5. know the trends in solubility of the hydroxides of Group 2 elements
    Group II hydroxides become more soluble down the group. All Group II hydroxides when not soluble appear as white precipitates. Mg(OH)2 is classed as insoluble in water. Simplest ionic equation for formation of Mg(OH)2(s)= Mg2+(aq)+2OH-(aq)->Mg(OH)2(s). A suspension of Mg(OH)2 in water will appear slightly alkaline (pH9) so some hydroxide ions must therefore have been produced by a very slight dissolving.
  • 5. know the trends in solubility of the hydroxides of Group 2 elements

    Mg(OH)2 is used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation. laxative, antacid. Mg(OH)2+2HCl->MgCl2+H2O. It is safe to use because it so weakly alkaline. It is preferable to using CaCO3 as it will not produce CO2 gas.
  • 5. know the trends in solubility of the hydroxides of Group 2 elements
    Ca(OH)2 is reasonably soluble in water. it is used in agriculture to neutralise acidic soils. An aqueous solution (saturated) of Ca(OH)2 is called limewater and can be used a test for CO2. The limewater turns cloudy as white CaCO3 is produced. Ca(OH)2(aq)+CO2(g)->CaCO3(s)+H2O(l). Ba(OH)2 would easily dissolve in water. The hydroxide ions present would make the solution strongly alkaline. Ba(OH)2(aq)+aq->Ba2+(aq)+2OH-(aq)
  • 5. know the trends in solubility of the hydroxides of Group 2 elements
    Barium meal shows soft tissue of the stomach and intestine. This uses x-rays as they cannot pass through the barium meal. Ba2+ ions in solution (usually BaCl2 or BaNO3) will form a white precipitate of BaSO4 with SO42- ions. Ba2+(aq)+SO42-(aq)->BaSO4(s). Ba2+ ions are poisonous to humans. Barium sulfate is insoluble so safe.
  • 5. know the trends in solubility of the hydroxides of Group 2 elements
    Add dilute HNO3 and Ba(NO3)2 to solution containing SO42- ions
    White precipitate forms
    Ba(NO3)3 (aq)   +   Na2SO4 (aq)  🡺   BaSO4 (s)   +   2NaNO3 (aq)
    Dilute HNO3 or HCl is added to provide H+ ions to prevent BaCO3 forming a white precipitate if CO32- ions are present in the solution (H2SO4 cannot be added)
  • 5. know the trends in solubility of the sulfates of Group 2 elements
    Group II sulfates become less soluble down the group. BaSO4 is the least soluble. If barium meal is reacted with H2SO4 it will only react slowly as the insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack. Ba+H2SO4->BaSO4+H2. The same effect will happen to a lesser extent with metals going up the group as the solubility increases. The same effect does not happen with other acids like HCl or H2NO3 as they form soluble group 2 salts.
  • 6. understand reasons for the trends in thermal stability of the nitrates and the carbonates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved
    Thermal decomposition is defined as the use of heat to break down a reactant into more than one product. Thermal stability- how stable a compound is when heated, beryllium carbonate is so unstable that it does not exist. Group II carbonates decompose on heating to produce group II oxides and CO2 gas. The ease of thermal decomposition decreases down the group.
  • 6. understand reasons for the trends in thermal stability of the nitrates and the carbonates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved

    Group 2 carbonates become more thermally stable going down the group. As the cations get bigger they have less of a polarising effect and distort the carbonate ion less. The C-O bond is weakened less so it less easily breaks down. MgCO3(s)->MgO(s)+CO2(g). White solid, white solid, colourless gas. Group 2 nitrates and carbonates decompose when heated, they don't melt. MCO3 insoluble in water.
  • 6. understand reasons for the trends in thermal stability of the carbonates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved
    Metal ions become larger down group 2 but have the same charge. This means their charge density is reduced. A metal ion with a high charge density has strong polarising power. It can therefore polarise the CO3- ion, making it more likely to split into O2- and CO2 when heated. A metal ions with a low charge density has weaker polarising power, meaning the CO3- ion is less polarised and therefore more thermally stable.
  • 6. understand reasons for the trends in thermal stability of the carbonates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved
    Group 1 carbonates do not decompose with the exception of lithium. As they only have +1 charges they don't have a big enough charge density to polarise the CO3- ion. Li is the exception because its ion is small enough to have a polarising effect. LiCO3(s)->Li2O(s)+CO2(g)
  • 6. understand reasons for the trends in thermal stability of the nitrates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved

    Group 2 nitrates decompose on heating to produce Group 2 oxides, oxygen, nitrogen dioxide. 2M(NO3)2(s)->2MO(s)+4NO2(g)+O2(g). The ease of thermal decomposition decreases down the group. You would observe brown gas evolving (NO2) so do it in a fume cupboard, and the white nitrate solid is seen to melt to a colourless solution and then re-solidify. M(NO3)2 is very soluble in water, white solid.
  • 6. understand reasons for the trends in thermal stability of the nitrates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved
    The explanation for change in thermal stability is the same as for carbonates. Mg(NO3)2 decomposes the easiest because the Mg2+ ion is the smallest and has a greater charge density. It causes more polarisation of the nitrate anion and weakens the N-O bond.
  • 6. understand reasons for the trends in thermal stability of the nitrates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved
    Group 1 nitrates, with the exception of LiNO3, do not decompose in the same way as group 2 nitrates. They decompose to give a nitrate (III) salt and oxygen. 2NaNO3 (sodium nitrate V) ->2NaNO2 (sodium nitrate III)+O2 Lithium nitrate decomposes in the same way as group 2 nitrates. 4LiNO3->2Li2O+4NO2+O2 smallest group 1 cation
  • 7. understand the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions. Students will be expected to know the flame colours for Groups 1 and 2 compounds.
    Flame tests work for some Group 1 and 2 metals.
    Lithium:carmine red, sodium:yellow/orange, potassium:lilac rubidium:red/purple, caesium:blue-violet, magnesium:no flame colour (energy emitted of a wavelength outside of visible spectrum) calcium:brick red, strontium:crimson, barium:(apple)green, copper:blue/green, beryllium:no colour
  • 8. understand experimental procedures to show: ii flame colours in compounds of Group 1 and 2 elements
    Method:Use a nichrome wire (nichrome is an unreactive metal and will not give out any flame colour). Clean the wire by dipping in concentrated HCl and then heating in bunsen flame. If the sample is not powdered then grind it up. Dip wire in solid and put in bunsen flame and observe colour.
  • 7. understand the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions Students will be expected to know the flame colours for Groups 1 and 2 compounds.

    Add a few drops of concentrated HCl to the solid and mix together so the metal compound begins to dissolve-this converts the metal compound to a chloride. Chlorides are more volatile than other salts so give better results.
  • 7. understand the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions Students will be expected to know the flame colours for Groups 1 and 2 compounds.

    Explanation for occurrence of flame: In a flame test the heat causes the electron to move to a higher energy level. Ground->excited state. The electron is unstable at the higher energy level and so drops back down. As it drops back down from the higher to lower energy level, energy is emitted in the form of visible light energy with the wavelength of the observed light.
  • 7. understand the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions Students will be expected to know the flame colours for Groups 1 and 2 compounds.
    Problems with flame tests:
    Many compounds contain small amounts of sodium compounds as impurities, the intense colour of sodium masks the other colours. Colours are subjective as different levels of colour vision in people. "Brick" red depends on the brick. Lilac/lavender/mauve/purple are similar.
    Ammonia gas turns damp red litmus paper blue.
    NH3+HCl(conc)->NH4Cl. NH4+ +OH- ->NH3+H2O
  • 8. understand experimental procedures to show: i patterns in thermal decomposition of Group 1 and 2 nitrates and carbonates
    Heat a known mass of carbonate in a side arm boiling tube and pass the gas produced through lime water. Time for the first permanent cloudiness to appear in the limewater. Repeat for different carbonates using the same moles of carbonate/same volume of limewater/same Bunsen flame and height of tube above flame. (DON'T MEASURE VOLUME OF GAS)