Doc Brown's Chemistry  Advanced Level Inorganic Chemistry – Periodic Table Revision Notes

Part 7. s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals – Sections 7.5 to 7.8

 

Group I & Group II Part 7.5 Describes the reaction of Groups 1/2 metals with oxygen and the chemical character of the oxides and peroxides formed. The reaction of the oxides with water and acids is given.

7.6 Describes and explains the reaction with water and the hydroxides formed. The theory behind the group 1/2 reactivity trend is discussed.

 7.7 Where safe to do so! the reaction of Alkali/Alkaline Earth Metals with acids is described.

 7.8 Describes the reaction of Groups 1/2 metals with the halogen element chlorine and the character of the halide salts formed.

 INDEXES BELOW

Pd s block elements d blocks and f blocks of metallic elements  p block elements
Gp1 Gp2 Gp3/13 Gp4/14 Gp5/15 Gp6/16 Gp7/17 Gp0/18
1

1H

2He
2 3Li

lithium

4Be

beryllium

ZSymbol, z = atomic or proton number

highlighting position of Group 1 and Group 2 elements

outer electrons: 1 ns1 and ns2

5B 6C 7N 8O 9F 10Ne
3 11Na

sodium

12Mg

magnesium

13Al 14Si 15P 16S 17Cl 18Ar
4 19K

potassium

20Ca

calcium

21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb

rubidium

38Sr

strontium

39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs

caesium

56Ba

barium

57-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb 83Bi 84Po 85At 86Rn
7 87Fr

francium

88Ra

radium

89-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113Uut 114Fl 115Uup 116Lv 117Uus 118Uuo

7.5 Oxygen reaction & oxides * 7.6 Water reaction & hydroxides * 7.7 Acid reaction & salts * 7.8 chlorine reaction – halides

(c) doc b GCSE/IGCSE Notes Alkali Metals (c) doc b GCSE/IGCSE Periodic Table Notes A level Quiz on basic s–block chemistry

INORGANIC Part 7 s–block Gp 1 Alkali Metals/Gp 2 Alkaline Earth Metals  sub–index: 7.1 Introduction to s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals  * 7.2 Group 1 data and graphs * 7.3 Group 2 data and graphs * 7.4 General trends down groups I & II and formulae *7.5 Oxygen reaction & oxides of s–block metals * 7.6 Water reaction & hydroxides of group 1/2 metals * 7.7 Acid reaction & salts of group1/2 metals * 7.8 chlorine reaction & halide of group I/II metals * 7.9 carbonates & hydrogen carbonates of s–block metals * 7.10 Solubility trends of groups 1/2 OH, NO3,SO4,CO3 compounds * 7.11 Thermal decomposition and stability of group 1 and group 2 carbonates & nitrates * 7.12 Uses of s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals and their compounds

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub–indexes near the top of the pages


7.5. The reaction of s–block metals and oxygen & their oxide (O2–) chemistry

The oxides and hydroxides are white ionic solids.

  • The reaction of Group 1 metals with oxygen (a redox reaction)

  • Group 1 metals: 4M(s) + O2(g) ==> 2M2O(s)    (M = Li, Na, K, Rb, Cs)

    • shows the formation of the 'simple' oxide expected from their position in the periodic table when the element is heated or burned in air.

    • Oxidation state changes: M is 0 to +1, Oxygen is 0 to –2 in the oxide ion O2–.

      • ionically: 4M(s) + O2(g) ==> 2(M+)2O2–(s)

        • the metal is oxidised (0 to +1), electron loss, increase in oxidation state

        • oxygen molecules are reduced (0 to –2), electron gain, decrease in oxidation state

    • The oxides are soluble in water forming the strongly alkaline hydroxide:

      • M2O(s) + H2O(l) ==> 2MOH(aq)

      • ionically: (M+)2O2–(s) + H2O(l) ==> 2M+(aq)  +  2OH(aq) (not a redox change)

        • This is an acid–base reaction, the O2– ion is a strong Bronsted–Lowry base and accepts a proton from water (acting as the Bronsted–Lowry acid).

    • Unfortunately, except for lithium (an anomaly), 'higher' oxides can be formed e.g.

    • 2M(s) + O2(g) ==> M2O2(s)  [redox change, M (0 to +1), O (0 to –1)]

      • shows the formation of the yellow–orange peroxide by Na, K, Rb and Cs

      • each oxygen is in the –1 oxidation state in the peroxide ion O22– 

      • they readily hydrolyse with water forming hydrogen peroxide

        • M2O2(s) + 2H2O(l) ==> 2MOH(aq) + H2O2(aq)  (not a redox change)

    • M(s) + O2(g) ==> MO2(s)  shows the formation of the 'superoxide' by K, Rb and Cs

      • oxidation number changes are M from 0 to +1 as expected, but on average each oxygen changes from 0 to  1/2  in the superoxide ion O2 

      • 2MO2(s) + 2H2O(l) ==> 2MOH(aq) + H2O2(aq) + O2(g)  (redox change)

      • oxidation state changes: M and H no change (+1), four O's change from  –1/2  in superoxide ions to two of –1 in the peroxide molecule and two at zero in the oxygen molecule.

        • This is a case of disproportionation where the oxidation state of an element gives a higher and lower state product from the same 'original species'.

  • The simple oxides readily dissolve in acids and are neutralised to form salts.

    • M2O(s) + 2HCl(aq) ==> 2MCl(aq) + H2O(l)   (M = Li, Na, K, Rb, Cs)

      • to give the soluble chloride salt

      • ionically: (M+)2O2–(s) + 2H+(aq) ==> 2M+(aq) + H2O(l)  (not a redox change)

        • Acid–base reaction, acid donates proton the oxide ion base, applies to all four examples.

        • The chloride Cl, nitrate NO3 and sulphate SO42– are spectator ions.

    • M2O(s) + 2HNO3(aq) ==> 2MNO3(aq) + H2O(l) to give the soluble nitrate salt

    • M2O(s) + H2SO4(aq) ==> M2SO4(aq) + H2O(l) to give the soluble sulphate salt

    • M2O(s) + 2CH3COOH(aq) ==> 2CH3COOM(aq) + H2O(l) to give the soluble ethanoate salt

  • The reaction of Group 2 metals with oxygen (a redox reaction)

  • Group 2 metals:  2M(s) + O2(g) ==> 2MO(s)    (M = Be, Mg, Ca, Sr, Ba)

    • shows the formation of the oxide expected from their position in the periodic table when the element is heated or burned in air. Oxidation state changes: M from 0 to +2, and oxygen from 0 to –2.

    • The oxide, apart from beryllium, is slightly soluble in water forming the alkaline hydroxide, which increases in strength of basic character down the group.

      • MO(s) + H2O(l) ==> M(OH)2(s/aq)   (not a redox change, M = Be, Mg, Ca, Sr, Ba)

        • ionically: M2+O2–(s) + H2O(l) ==> M(OH)2(s/aq)

          • if the hydroxide is soluble: M2+O2–(s) + H2O(l) ==> M2+(aq) + 2OH(aq)

          • Bronsted–Lowry acid–base reaction, the oxide base accepts proton from the water.

          • The mixture of magnesium hydroxide and water is sometimes called milk of magnesia.

          • The formation of calcium hydroxide (slaked lime) when water is added to calcium oxide (quicklime) is very exothermic!

          • The pH of the resulting solution ranges from ~pH 10 to ~pH 13 for Mg(OH)2 to Ba(OH)2

      • All the oxides are basic and readily neutralised by acids (not a redox change).

        • MO(s) + 2HCl(aq) ==> MCl2(aq) + H2O(l)   (M = Be, Mg, Ca, Sr, Ba)

          • to give the soluble chloride salt

          • ionically: M2+O2–(s) + 2H+(aq) ==> M2+(aq) + H2O(l) 

          • This applies to all four acid reactions examples in this section, acid proton donation to the oxide ion base.

          • In each case the chloride Cl, nitrate NO3 and sulphate SO42– are spectator ions.

      • MO(s) + 2HNO3(aq) ==> M(NO3)2(aq) + H2O(l)   (M = Be, Mg, Ca, Sr, Ba)

        • to give the soluble nitrate salt

      • MO(s) + H2SO4(aq) ==> MSO4(aq/s) + H2O(l)    (M = Be, Mg, Ca, Sr, Ba)

        • to form the sulphate salt (soluble => insoluble)

        • but reaction increasingly slower for calcium oxide ==> barium oxide as the sulphate becomes less insoluble.

      • MO(s) + 2CH3COOH(aq) ==> (CH3COO)2M(aq) + H2O(l)  (M = Be, Mg, Ca, Sr, Ba)

        • to give the ethanoate salt

      • Beryllium oxide BeO is amphoteric (another Be Gp 2 anomaly) and dissolves in strong bases like sodium hydroxide.

        • The equation below shows the formation of a hydroxo beryllate complex ion (not a redox change).

        • BeO(s) + 2NaOH(aq) + H2O(l) ==> Na2[Be(OH)4](aq) (beryllate salt)

        • ionically: Be2+O2–(s) + 2OH(aq)  + H2O(l) ==> [Be(OH)4]2–(aq)  


 

7.6. Reaction of s–block metals and water & their hydroxide (OH) chemistry

The oxides and hydroxides are usually white ionic solids.

  • (c) doc bThe reaction of group 1 metals with water (a redox reaction)

  • Group 1 metal hydroxide formation

    • 2M(s) + 2H2O(l) ==> 2M+OH(aq) + H2(g)   (M = Li, Na, K, Rb, Cs)

    • shows the formation of the alkaline metal hydroxide and hydrogen.

      • Oxidation state changes: M from 0 to +1, one H per water remains unchanged in oxidation number and one changes from +1 to 0 in H2.

    • M = Li (slow at first), Na (fast), K (faster – may ignite hydrogen to give a lilac coloured flame* from hot potassium atoms), Rb, Cs, Fr (very explosive) i.e. the reactivity increases down the group.

    • The hydroxides, MOH, are white ionic solids, all very soluble (except  LiOH), strong bases, getting stronger down the group.

      • All Group 1 hydroxides are soluble in water giving strongly alkaline solutions,

      • and their aqueous solutions readily neutralised by acids (not a redox change) e.g.

      • MOH(aq) + HCl(aq) ==> MCl(aq) + H2O(l)   (M = Li, Na, K, Rb, Cs)

        • to give the soluble chloride salt*

        • ionically: OH(aq) + H+(aq) ==> H2O(l)  

        • an acid–base reaction, same for all four examples in this section

      • MOH(aq) + HNO3(aq) ==> MNO3(aq) + H2O(l)   (M = Li, Na, K, Rb, Cs)

        • to give the soluble nitrate salt

      • 2MOH(aq) + H2SO4(aq) ==> M2SO4(aq) + 2H2O(l)   (M = Li, Na, K, Rb, Cs)

        • to give the soluble sulphate salt

      • MOH(aq) + CH3COOH(aq) ==> CH3COOM(aq) + H2O(l)   (M = Li, Na, K, Rb, Cs)

        • to give the soluble ethanoate salt CH3COOM+

      • * The hydroxide solutions are readily titrated with standardised hydrochloric acid (burette) using phenolphthalein indicator, the colour change is from pink to colourless.

  • (c) doc bThe reaction of group 1 metals with water (a redox reaction)

  • Group 2 metal hydroxide formation

    • M(s) + 2H2O(l) ==> M(OH)2(aq/s) + H2(g)    (M = Mg, Ca, Sr, Ba)

    • shows the formation of the hydroxide and hydrogen with cold water.

    • ionically: M(s) + 2H2O(l) ==> M2+(aq) + 2OH(aq) + H2(g)

    • oxidation number changes, M is 0 to +2, for one H per water it changes from +1 to 0 in H2.

    • M = Be (no reaction, anomalous), Mg (very slow reaction), Ca, Sr, Ba (fast to very fast).

      • i.e. the reactivity increases down the group.

      • The reactivity trend for Group 2, and its explanation, are similar to that above for the Group 1 Alkali Metals.

      • The reactivity trend for s–block metals is explained below

      • Magnesium hydroxide and calcium hydroxide (limewater) are sparingly soluble, but the solubility increases down the group, so barium hydroxide is moderately soluble.

      • As previously mentioned, a mixture of magnesium oxide/hydroxide and water is sometimes called milk of magnesia and the saturated aqueous solution of calcium hydroxide is called limewater.

  • If the metal is heated in steam the oxide is formed:

    • e.g. Mg(s) + H2O(g) ==> MgO(s) + H2(g) 

      • NOT an experiment you would do with Alkali Metals! but beryllium gives little reaction.

    • The oxide is formed because the hydroxide is thermally unstable at higher temperatures

      •   M(OH)2(s) ==> MO(s) + H2O(g)    (M = Be, Mg, Ca, Sr, Ba)

  • REACTIVITY TREND THEORY: The Group 1/2 metal gets more reactive down the group because ...

    • When an alkali metal atom reacts, it loses an electron to form a singly positively charged ion.

      • e.g. Na ==> Na+ + e

      • in terms of electrons 2.8.1 ==> 2.8 and so forming a stable ion with a noble gas electron arrangement.

    • As you go down the group from one element down to the next the atomic radius gets bigger due to an extra filled electron shell as you go down from one period to the next one.

    • This means the outer electron is further and further from the nucleus.

    • This also means the outer electron is also shielded by the extra full electron shell of negative charge.

    • Due to this shielding the effective nuclear charge on the external electron is ~ +1 (~ proton number – number of noble gas inner core electrons).

    • Further more, the effective nuclear charge of ~+1 is acting over a larger 'surface area' as the atomic radius increases.

    • Therefore both of these factors combine to make the outer electron less and less strongly held by the positive nucleus as the atomic number increases (down the group).

    • So, the outer electron is more easily lost, and the M+ ion more easily formed, and so the element is more reactive as you go down the group – best seen in the laboratory with their reaction with water.

    • The reactivity argument mainly comes down to increasingly lower ionisation energy down the group (i.e. ease of ion formation) and a similar argument applies to the Group 2 metals, but two electrons are removed to form the cation.

    • The enthalpy change in forming the hydrated cation from the solid metal does not appear to be as important here.

      • At a more advanced and detailed level, this change can be theoretically split into the

      • enthalpies of (i) atomisation, (ii) ionisation, (iii) hydration of gaseous ion ... (BUT not here!).

    • The reactivity trend is also paralleled by the increasingly negative half–cell potential (EθM/M+ and EθM/M2+) down groups, 1 and 2 i.e. increasing potential to acts as a reducing agent – an electron donor.

    • As with water, the reaction of a group 1/2 metal with oxygen or halogens gets more vigorous as you descend the group.

  • All the hydroxides are basic with increasing strength down the group and readily neutralised by acids (not redox reactions). Magnesium hydroxide is sparingly soluble in water but the solubility increases down the group.

    • M(OH)2(aq/s) + 2HCl(aq) ==> MCl2(aq) + 2H2O(l)    (M = Be, Mg, Ca, Sr, Ba)

      • to give the soluble chloride salt*

      • all base (OH) ... acid (H+) reactions

      • ionically if soluble the reaction is: OH(aq) + H+(aq) ==> H2O(l)  

      • ionically if insoluble: M2+(OH)2(s) + 2H+(aq) ==> M2+(aq) + 2H2O(l)

    • M(OH)2(aq/s) + 2HNO3(aq) ==> M(NO3)2(aq) + 2H2O(l)    (M = Be, Mg, Ca, Sr, Ba)  

      • to give the soluble nitrate salt

    • M(OH)2(aq/s) + H2SO4(aq) ==> M2SO4(aq/s) + 2H2O(l)    (M = Be, Mg, Ca, Sr, Ba)  

      • to give the sulphate salt

    • M(OH)2(aq/s) + 2CH3COOH(aq) ==> (CH3COO)2M(aq) + 2H2O(l)    (M = Be, Mg, Ca, Sr, Ba)  

      • to give the ethanoate salt

    • * Saturated calcium hydroxide solution (limewater) can be titrated with standardised hydrochloric acid (burette, low molarity) to determine its solubility. You normally use phenolphthalein indicator and the end–point colour change is from pink to colourless.

  • The Group 2 hydroxides, M(OH)2, get more soluble down the group:

    • If the hydroxide more or less insoluble (e.g. for Be and Mg), they can be made by adding excess sodium/potassium hydroxide solution to a solution of a soluble salt of a Group 2 metal e.g. three 'double decompositions' are shown below ...

      • (i) calcium chloride + sodium hydroxide ==> sodium chloride + calcium hydroxide

        • CaCl2(aq) + 2NaOH(aq) ==> 2NaCl(aq) + Ca(OH)2(s) 

      • (ii) magnesium sulphate + potassium hydroxide ==> potassium sulphate + magnesium hydroxide

        • MgSO4(aq) + 2KOH(aq) ==> K2SO4(aq) + Mg(OH)2(s) 

      • (iii) beryllium nitrate + sodium hydroxide ==> sodium nitrate + beryllium hydroxide

        • Be(NO3)2(aq) + 2NaOH(aq) ==> 2NaNO3(aq) + Be(OH)2(s) 

      • or ionically: M2+(aq) + 2OH(aq) ==> M(OH)2(s) for any Group 2 metal M

      • All the hydroxides are white powders or white gelatinous precipitates.

  • Beryllium hydroxide is amphoteric (an anomaly in the group), because apart from the reactions above, it dissolves in strong alkalis like sodium hydroxide to form a hydroxo–complex ion salts called 'beryllates' e.g.

    • Be(OH)2(s) + 2NaOH(aq) ==> Na2[Be(OH)4](aq)  (not a redox change)

    • ionically: Be(OH)2(s) + 2OH(aq) ==> [Be(OH)4]2–(aq) showing formation of a complex ion

  • For the reaction of Group 1 and 2 hydroxides with carbon dioxide to form the carbonates and hydrogen carbonates, see section 7.9

  • For the thermal decomposition of nitrates see section 7.11

  • See also VOLUMETRIC TITRATION QUESTIONS involving acids and group1 /2 alkaline hydroxides


 

7.7. The reaction of s–block metals with acids

  • Group 1 metals are far too reactive to contemplate adding them to acids in a school laboratory!

  • Group 2 metals, apart from beryllium (another anomaly), readily react with acids, with increasing vigour down the group (explanation in section 7.4). A redox reaction to form the soluble chloride salt.

    • M(s) + 2HCl(aq) ==> MCl2(aq) + H2(g)   (M = Mg, Ca, Sr, Ba)

      • ionically for all four examples: M(s) + 2H+(aq) ==> M2+(aq) + H2(g) 

      • oxidation state changes: one M at (0) and two H's at (+1) ==> one M (+2) and two H's at (0)

        • the metal is oxidised, electron loss, increase in oxidation state

        • hydrogen ions are reduced, electron gain, decrease in oxidation state

    • M(s) + 2HNO3(aq) ==> M(NO3)2(aq) + H2(g)

      • to form the soluble nitrate salt

      • Looks ok in principle, and does this with Mg and very dilute nitric acid, but rarely this simple, the nitrate(V) ion can get reduced to nasty brown nitrogen(IV) oxide gas (nitrogen dioxide, NO2) and other products, NO gas?, NO2 ion?

    • M(s) + H2SO4(aq) ==> MSO4(aq/ s) + H2(g)

      • to form soluble ==> insoluble sulphate salt

      • The reaction from magnesium to barium becomes increasingly slower as the sulphate becomes less soluble, it coats the metal, inhibiting the reaction.

    • M(s) + 2CH3COOH(aq) ==> (CH3COO)2M(aq) + H2(g)  to form soluble ethanoate salt

      • This reaction is much slower than the previous three because ethanoic acid is a weak acid (about 2% ionised, so the fizzing appears a lot less vigorous than the other three acids using solutions of similar molarity).

  • (c) doc bIn aqueous solutions the metal cations formed are hydrated to aqa–complex ions.

    • not quite the simple isolated ions Mn+(aq) which we use in most equations for brevity.

    • e.g. [M(H2O)6]n+(aq) where n=1 for Gp 1 and n=2 for group 2.

      • There may be several layers of water molecules around the ion, so the six is not the whole story, but is typical for the number of 'nearest neighbours', albeit weakly dative covalently bonded water molecules in this case.

      • The six is called the co–ordination number and each water molecule (or anything else attached to the central metal ion) is called a ligand

      • The shape of such an ion is 'octahedral' and its simplified structure is shown above on the right. The middle 'blob' is the metal ion and the six outer 'blobs' are the water molecules.

      • (c) doc b(I will replace with proper diagrams later) 

      • However lithium and beryllium are anomalous (M = Li n = 1, or Be n = 2), because of electronic quantum level restrictions, they can have a maximum co–ordination number of four, so their aqueous cations should be written as [M(H2O)4]n+(aq) which has a tetrahedral shape (shown on the left).

  • As described above, The soluble groups 1/2 salt solutions contain the hydrated cations derived from the metal:

    • tetra–aqua cations [Li(H2O)4]+(aq) and [Be(H2O)4]2+(aq)

    • or the hexa–aqua ions [M(H2O)6]+(aq) M = Na, K etc. for Group 1

    • and [M(H2O)6]2+(aq) where M = Mg, Ca etc. for Group 2

    • The tetraaqua beryllium ion and the hexaaqua magnesium ions generate a slight acidity in their salt solutions due to the significant polarising power of the ions (Be2+ very small and double charged, Mg2+ double charged) e.g.

    • for beryllium: [Be(H2O)4]2+(aq) + H2O(l) (c) doc b [Be(H2O)3(OH)]+(aq) + H3O+(aq) 

    • or magnesium: [Mg(H2O)6]2+(aq) + H2O(l) (c) doc b [Mg(H2O)5(OH)]+(aq) + H3O+(aq) 


 

7.8. The reaction of s–block metals with chlorine & halide (X) salts

The salts are white or colourless crystalline solids

  • Group 1 metals readily react with halogens (a redox reaction)  

    • e.g. heating the metal in chlorine will cause it to burn forming the chloride

    • 2M(s) + Cl2(g) ==> 2MCl(s)  (redox reaction, M = Li, Na, K, Rb, Cs)

      • Oxidation state changes: M from 0 to +1, X = F, Cl, Br & I from 0 to –1

      • The salt products, M+X, are white–colourless crystalline ionic solids that dissolve in water to give neutral solutions of about pH 7. The crystalline solids have high melting and boiling points.

      • The solids do not conduct electricity (no mobile ions or electrons) but will conduct and undergo electrolysis when molten or dissolved in water when ions are free to move to electrodes.

    • The halogen is in the –1 oxidation state in the halide ion X 

    • The halides of groups 1–2 are important raw materials e.g.

      • sodium chloride ==> sodium hydroxide from rock salt by electrolysis of aqueous solution

      • potassium bromide/iodide ==> elemental bromine/iodine from seawater by oxidation

      • calcium chloride ==> calcium metal by electrolysis of molten chloride

  • Group 2 metals (except Be) readily react on heating with halogens (a redox reaction)   

    • e.g. heating in chlorine the chloride is formed

    • M(s) + Cl2(g) ==> MCl2(s)    (M = Mg, Ca, Sr, Ba)

      • Oxidation state changes: M from 0 to +2, X = F, Cl, Br & I from 0 to –1

      • The salt products, M2+(X)2,  are similar in properties to the Group 1 M+X compounds.

      • However, beryllium chloride has a polymeric covalent structure, due to the high polarising influence of beryllium in its +2 oxidation state and the smaller difference in electronegativity between Be–Cl compared to chlorine and the other group 1 and 2 metals.

INORGANIC Part 7 s–block Gp 1 Alkali Metals/Gp 2 Alkaline Earth Metals  sub–index: 7.1 Introduction to s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals  * 7.2 Group 1 data and graphs * 7.3 Group 2 data and graphs * 7.4 General trends down groups I & II and formulae *7.5 Oxygen reaction & oxides of s–block metals * 7.6 Water reaction & hydroxides of group 1/2 metals * 7.7 Acid reaction & salts of group1/2 metals * 7.8 chlorine reaction & halide of group I/II metals * 7.9 carbonates & hydrogen carbonates of s–block metals * 7.10 Solubility trends of groups 1/2 OH, NO3,SO4,CO3 compounds * 7.11 Thermal decomposition and stability of group 1 and group 2 carbonates & nitrates * 7.12 Uses of s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals and their compounds

(c) doc b GCSE/IGCSE Notes Alkali Metals (c) doc b GCSE/IGCSE Periodic Table Notes

A level Quiz on basic s–block chemistry


keywords equations: CaCl2(aq) + 2NaOH(aq) ==> 2NaCl(aq) + Ca(OH)2(s) * MgSO4(aq) + 2KOH(aq) ==> K2SO4(aq) + Mg(OH)2(s) * Ba(NO3)2(aq) + 2NaOH(aq) ==> 2NaNO3(aq) + Ba(OH)2(s) * CaCl2 + 2NaOH ==> 2NaCl + Ca(OH)2 * MgSO4 + 2KOH ==> K2SO4 + Mg(OH)2 * Ba(NO3)2 + 2NaOH ==> 2NaNO3 + Ba(OH)2 *

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