3.1.2 Group 2 Elements Flashcards
how do alkali metals act as “reducing agents” in redox
- reduce the species they interact with, as in redox:
- metal atom is oxidised
- loses 2 e- in outer s-subshell
- form 2+ ion, with same configuration as noble gas
- e.g. Ca ===> Ca2+ + 2e- ([Ar]4s2 ===> [Ar])
how do alkali metals react with oxygen
- undergo redox, forming metal oxides (GF: MO, where M2+ and O2-)
give example of magnesium reacting with oxygen
- burns with O2 to give bright light and form white oxide
- 2Mg(s) + O2(g) ===> 2MgO(s)
- Mg oxidised (0 to +2)
- O2 reduced (0 to -2)
how do alkali metals react with water
- form an alkaline hydroxide (GF: M(OH)2) and hydrogen gas
- originally a slow reaction, like with Mg, but increases down group as does reactivity
give example of strontium reacting with water
Sr(s) + 2H20(l) ===> Sr(OH)2(aq) + H2(g)
-Sr oxidised (from 0 to +2)
-H reduced from +1 to 0)
how do alkali metals react with dilute acids
metal + dilute acid ===> salt + hydrogen
- reactivity increases down the group
- can have MONObasic acids, e.g. HCl, which donate 1 proton per molecule, or DIbasic (H2SO4) and TRIbasic (H3PO4)
give example of magnesium reacting with dilute hydrochloric acid
Mg(s) + 2HCl(aq) ===> MgCl2(aq) + H2(g)
- Mg oxidised ( 0 to +2 )
- H reduced ( +1 to 0 )
what is the pattern of reactivity down group 2 elements
reactivity INCREASES
what is the main energy input when alkali metals react and how does it occur
- 2 ionisation energies
- lose 2 e- and form 2+ gaseous ions
- other energy changes do take place, but IEs are the main one
why does ionisation energy DECREASE down group 2 elements
- increase in atomic radius
- increase in shielding
- (increase in nuclear charge too, but this is outweighed by other factors)
- decrease in ES attraction between nucleus and outer shell electrons
- so more reactive
- and more stronger reducing agents
what are group 2 compounds
oxides, hydroxides, carbonates
how do oxides react with water
- release hydroxide ions (OH-) and form alkaline solutions
- go on to form metal hydroxides
give example of CaO reacting with water
CaO(s) + H2O(l) ===> Ca2+ + 2OH-
- ions add to make Ca(OH)2 (s)
why do most reactions of oxides with water result in solid precipitate formed
- hydroxides aren’t very soluble
- one solution is saturated, form solid precipitate form
- which is typical form - (s) and (aq)
what happens to the alkalinity of hydroxides as you go down the group
INCREASES
why does alkalinity increase down group 2
- solubility increases down group
- so solution contains more OH- ions
- increasing alkalinity
- pH range from 10 (Mg) to 13 (Ba)
explain an experiment to show the varying alkalinity of the group 2 hydroxides
- add spatula of each oxide to water
- shake
- should get a saturated solution and remaining white precipitate which has not dissolved, and insufficient amount of H2O at this scale to dissolve all hydroxide
- measure pH of each hydroxide
- value should increase
what are the 2 uses of alkaline compounds
1) AGRICULTURE - neutralise acidic soil
2) MEDICINE - neutralise stomach acid
explain use of alkaline compounds in agriculture
- Ca(OH)2, calcium carbonate, is added to fields as lime, which increases the pH of acidic soil, neutralising it, and forming H2O
- Ca(OH)2(aq) + 2H+(aq) ===> Ca2+(aq) + 2H20(l)
explain the use of alkaline compounds in medicine
- used as antacids to treat stomach indigestion
- neutralise excess stomach acid, which is mainly HCl
- use MgCO3 (magnesium carbonate), CaCO3 (calcium carbonate)
- or use “milk of magnesia”, Mg(OH)2 (magnesium hydroxide), which is a suspension of white Mg(OH)2 in water, as only slightly soluble in water
CaCO3(s) + 2HCl ===> CaCl2 + H20 + CO2
Mg(OH)2(s) + 2HCl ===> MgCl2 + 2H20
