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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 7.1 to 7.4

Group I & Group II Part 7.1 gives an introduction to the chemistry of Group 1 Alkali Metals and Group 2 Alkaline Earth Metals via their outer electron configuration. 7.2 and 7.3 gives data tables and graphs for the s–block groups and 7.4 discusses the general group trends in 1st ionisation energy (and patterns in successive ionisation energies), atomic radii, ion (cation) radii, electronegativity, melting points, boiling points and the formulae of common compounds. INDEXES BELOW

7.1 Introduction * 7.2 Group 1 data and graphs * 7.3 Group 2 data and graphs * 7.4 General trends down the group and formulae

(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


7.1. Introduction to Group 1 Alkali Metals and Group 2 Alkaline Earth Metals

  • The first two vertical columns of the Periodic Table, i.e. Groups 1 and 2, are called the s–block metals, because they only have 1 or 2 electrons in their outer shell.

  • These outer electrons are of an s–orbital type (s sub–shell or sub–quantum level) and the chemistry of the metals, with their relatively low ionisation energies, is dominated by the loss of these s electrons to form a cation and also accounts for their generally high chemical reactivity ...

    • the outer s1 electron loss by the Group 1 Alkali Metals gives the M+ ion, and,

    • the outer s2 electrons lost by the Group 2 Alkaline Earth Metals forms the M2+ ion,

    • and in each case the cation has a residual very stable noble gas core of electrons.

  • The only chemically stable oxidation states are +1 for Group 1 metals and +2 for Group 2 elements, governed by the relative ease of loss of the outer s electrons and the subsequent very high ionisation energies required to remove a 2nd (for group 1) or 3rd electron (for group 2) from the inner noble gas core of electrons left (i.e. from the next principal quantum level or shell).

  • The relative ease of delocalising the outer 1/2 electrons in the metal lattice makes them good conductors of heat and electricity (bonding model for metals).

  • The low ionisation energies and low electronegativity means that when combined with non–metals, most compounds of the Group 1–2 elements tend to be ionic in nature.

  • Group 1and Group 2 ions are their compounds are important in the natural world of living systems and geology.

    • Calcium carbonate and phosphate minerals are important components of skeletons, teeth as well as bone!

    • For marine organisms the skeletal remains form sedimentary rocks like limestone and chalk which mainly consist of calcium carbonate and magnesium carbonate.

    • In biochemistry, at the heart of the chlorophyll molecules involved with photosynthesis is magnesium ion (Mg2+).

    • Sodium and potassium ions (Na+ and K+) are important components of a balanced electrolyte solutions in living systems and in nerve impulse transmission systems.

  • some other LINKS

Advanced Inorganic Chemistry Page Index and Links


 

7.2. Information and Data Table GROUP 1 ALKALI METALS (c) doc b (from left to right is down the group!)

property\Z symbol, name 3Li Lithium 11Na Sodium 19K Potassium 37Rb Rubidium 55Cs Caesium 87Fr Francium
melting point/oC 181 98 64 39 29 27
boiling point/oC 1347 883 774 688 679 677
density/gcm–3 0.53 0.97 0.86 1.48 1.87 >1.87
1st IE/kJmol–1 513 496 419 403 376 400
2nd IE/kJmol–1 7298 4562 3051 2632 2420 2100
atomic metallic radius/pm 152 186 231 244 262 270
M+ ionic radius/pm 78 98 133 149 165 180
electronegativity 0.98 0.93 0.82 0.82 0.79 0.70
electron configuration 2,1 2,8,1 2,8,8,1 2,8,18,8,1 2,8,18,18,8,1 2,8,18,32,18,8,1
electron configuration [He]2s1 [Ne]3s1 [Ar]4s1 [Kr]5s1 [Xe]6s1 [Rn]7s1
Electrode potential M/M+ –3.04V –2.71V –2.92V –2.92V –2.92V –2.92V
Symbol – flame colour Li – red/crimson Na – yellow K – lilac/purple Rb – red Cs – blue Fr – na
  • Typical metals in some ways e.g. silvery grey lustrous solids*, very good conductors of heat and electricity, relatively high boiling points.

    • * When freshly cut they are quite shiny, but they rapidly tarnish by reaction with oxygen to form an oxide layer, which is why they are stored under oil. Ask your teacher to show you the bottle from which we extract the lumps for doing our demonstrations – not a pretty sight if its a bit old!

    • The 1st ionisation energies are the lowest of any group of elements, but note the jump up to a very high 2nd ionisation energy.

      • Equations for the 1st and 2nd ionization energies.

        • M(g) ==> M+(g) + e–   (very low 1st IE, M = Li, Na, K, Rb, Cs, Fr)

        • M+(g) ==> M2+(g) + e–   (much higher 2nd IE)

        • The very high 2nd ionization energy is due to removing an electron from an electronically very stable noble gas inner core of electrons.

  • Untypical in other ways e.g. relatively soft, low density (Li, Na and K float on water before reacting ...), and very low melting points.

    • The reasons for the low melting points, densities and physical hardness lies in looking at the bonding model for metals.

    • The more electrons that can be delocalised and the closer the atoms (actually ions in a metallic lattice) can approach each other the stronger the bond, thereby raising melting points, densities and tensile strength.

    • However, Group 1 metals can only release one electron per atom to contribute towards the pool of bonding delocalised electrons. Group 2, 3 and transition metals etc. can contribute more electrons for metallic bonding. This why the Group 2 elements have higher melting points, densities and are harder metals.

    • Also, being the first element on a given period, they have the highest relative atomic/ionic radii because you have the minimum nuclear attractive force on the outer electrons.

  • Any metal flame colour is due to electronic transitions in the atom or cation.

    • Electrons are promoted to higher quantum levels via collisions of the high thermal kinetic energy particles in the hot flame. When the promoted electron 'relaxes' or 'falls' back to its more stable electronic level, energy is emitted (this is the basis of an emission spectrum). If the frequency/wavelength/energy of the photons emitted is in the visible region of the electromagnetic spectrum, a 'flame colour' results e.g. as observed in fireworks.

      • The set of quantum levels and associated energies are unique for each atom which means the quantum level differences varies from atom to atom, therefore the frequency/energy of emitted photons is different, hence you see different flame colours in the visible region of light from each Group 1/2 metal – each element has a 'finger print' emission spectrum.

      • The s–block groups 1/2 have the lowest ionisation energies of all the elements in the periodic table, so the high temperature of the flame (>1000oC) means that the kinetic energy of the flame particles is quite sufficient to promote electrons to a higher quantum level when the metal atoms collide with other high KE atoms or molecules.

        • Planck's Equation: ΔE = hν, where

        • ΔE = E2 – E1, the energy difference between e.g. the outer s level E1 and a higher level E2,

        • h = Planck's constant and ν = frequency of light of the emitted photons.

        • For more examples–details see electronic structure and emission/absorption spectra notes.

  • Oxidation state or oxidation number is always +1 in Group 1 Alkali Metal  compounds.

    • Only the single outer s–electron is easily lost, the 2nd, and subsequent ionisation energies are far too high to form chemically stable cations of 2+ etc. i.e. the energy required will not be compensated by ionic bond formation.

    • The stable Group 1 cation has the electron configuration of a noble gas,

      • e.g. the sodium atom, Na, is 2,8,1 or 1s22s22p63s1 or [Ne]3s1

      • so the sodium ion, Na+, is 2,8 or 1s22s22p6 or [Ne]

  • See also section 4. on Group trends and comparison with Group 2 metals.

  • (c) doc bPLEASE NOTE that Francium is highly radioactive and therefore difficult and dangerous to study BUT all its known physical and chemical properties fit in with it being at the foot of Group 1 and other properties could be inferred from the properties and group trends of Li to Cs.

Advanced Inorganic Chemistry Page Index and Links


 

7.3. Information and Data Table GROUP 2 ALKALINE EARTH METALS (c) doc b (from left to right is down the group!)

property\Zsymbol, name 4Be Beryllium 12Mg Magnesium 20Ca Calcium 38Sr Strontium 56Ba Barium 88Ra Radium
melting pt./oC 1278 649 839 769 729 700
boiling pt./oC 2487 1090 1484 1384 1637 1140
density/ gcm–3 1.85 1.74 1.55 2.54 3.51 5.0
1st IE/ kJmol–1 900 738 590 550 503 509?
2nd IE/kJmol–1 1757 1451 1145 1064 965 979
3rd IE/kJmol–1 14848 7733 4910 4210 3600 3300
atomic radius/ pm 111 160 197 215 217 223
M2+ ionic radius/ pm 34 78 106 127 143 152
electronegativity 1.57 1.31 1.00 0.95 0.89 0.89
electron config. 2,2 2,8,2 2,8,8,2 2,8,18,8,2 2,8,18,18,8,2 2,8,18,32,18,8,2
electron config. [He]2s2 [Ne]3s2 [Ar]4s2 [Kr]5s2 [Xe]6s2 [Rn]7s2
Electrode pot'ial M/M2+ –1.97V –2.36V –2.84V –2.89V –2.92V –2.92V
Symbol – flame colour (see chemical tests new window) Be – na Mg – na Ca – brick red Sr – crimson Ba – apple green Ra – na
  • Very typical metals, silvery grey lustrous solids, relatively high melting and boiling points, good conductors of heat and electricity.

    • The first two ionisation energies are relatively low but there is quite a jump to the 3rd ionisation energy.

      • Equations for the 1st, 2nd ionization energies.

      • M(g) ==> M+(g) + e–   (very low 1st, M = Be, Mg, Ca, Sr, Ba, Ra)

      • M+(g) ==> M2+(g) + e–   (2nd higher IE)

      • M2+(g) ==> M3+(g) + e–   (much higher 3rd IE)

      • The very high 3rd ionization energy is due to removing an electron from an electronically very stable noble gas inner core of electrons.

  • Compared to adjacent Group 1 metal on same period:

    • The melting and boiling points are higher, and they are harder, stronger and more dense than the adjacent Group 1 metal on the same period. This is because their are two delocalised electrons per ion in the crystal lattice giving an overall stronger electrical attraction with the more highly charged M2+ ions.

    • Chemically very similar e.g. form mainly ionic compounds but different formulae and less reactive because the 1st ionisation energies are higher (due to extra nuclear charge) and a 2nd ionisation energy input to form the stable M2+ ion.

  • Oxidation state or oxidation number is always +2 in Group 2 Alkaline Earth Metal compounds.

    • The two outer s–electrons are readily lost. The 3rd, and subsequent ionisation energies are far too high to form chemically stable cations of 3+ etc. i.e. the energy required will not be compensated by ionic bond formation.

    • The stable Group 2 cation has electron configuration of noble gas,

      • e.g. the calcium atom, Ca, is 2,8,8,2 or 1s22s22p63s23p64s2 or [Ar]4s2

      • so the calcium ion, Ca2+, is 2,8,8 or 1s22s22p63s23p6 or [Ar]

  • See also section 7.4 below on Group trends and comparison with Group 1 metals.

  • PLEASE NOTE that Radium is highly radioactive and therefore difficult and dangerous to study BUT all its known physical and chemical properties fit in with it being at the foot of Group 2 and other properties could be inferred from the properties of Mg to Ba.

Advanced Inorganic Chemistry Page Index and Links


 

7.4. General Trends down groups 1 & 2 with increasing atomic number and formula patterns

  • The 1st ionisation energy (IE) or 2nd etc. decrease: (important to link to reactivity trend)

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

    • The outer electrons are further and further from the nucleus and are also shielded by the extra full electron shell of negative charge. Therefore the outer electrons are less and less strongly held by the positive nucleus and so less and less energy is needed to remove them.

    • Note (probably not needed for exams, but a Q students might raise):

      • Despite the significant increase in atomic number i.e. positive nuclear charge down the group, this effect is outweighed by 'shielding effects' of inner full electron shells and also the nuclear charge is spread over an increasingly larger surface as the atomic radius increases.

      • Advanced Inorganic Chemistry Page Index and LinksThe effective nuclear charge is NOT what it seems as given by the atomic number and is more related to the number of outer electrons and the size of the atom.

IONISATION ENERGY PATTERNS

  • Successive ionisation energies always increase e.g. ... 3rd > 2nd > 1st ...

    • because the same nuclear charge is attracting fewer electrons and on average closer to the nucleus.

    • The negative electrons are being successively removed from an increasingly more positive ion, so, not surprisingly, more energy is required.

    • BUT note the 2nd IE for Group 1 (e.g. for potassium in the right graph above), and the 3rd IE for Group 2 (e.g. magnesium in the left graph above), show a particularly significant increase in IE compared to the previous ionisation energy or energies.

    • This is due to removing an electron from an electronically highly stable full inner shell and puts an upper limit on the chemically stable oxidation state.

    • These abrupt changes in the successive ionisation energy graphs are clear evidence of principal quantum levels that the electrons occupy.

  • Atomic and ionic radius increases:

    • Advanced Inorganic Chemistry Page Index and LinksBecause from one element to the next, an extra shell of electrons is added, increasing the electron 'bulk' and the outer electrons are increasingly less strongly held (see above).

    • The radii of the adjacent Gp 2 atom is smaller than Gp 1 atom on the same period, because the nuclear charge has increased by one unit (L to R in PT), but is attracting electrons in the same shell.

    • Similarly the radii of a Gp 2 M2+ ion is smaller than the adjacent Gp 1 M+ ion on the same period, because the nuclear charge has increased by one unit (L to R in PT), but is attracting the same number of electrons in the same shells. (see data tables in section 7.2 and section 7.3)

  • REACTIVITY TREND THEORY – relate to atomic radius and ionisation energy

    •  The 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.

    • The outer electron is further and further from the nucleus and is also shielded by the extra full electron shell of negative charge.

    • Therefore the outer electron is less and less strongly held by the positive nucleus.

    • This combination of factors means the outer electron is more easily lost, 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 and a similar argument applies to the Gp 2 metals, but two electrons are removed to form the cation.

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

  • Generally (but not always), the melting and boiling points fall steadily:

    • This is because the ionic radii increase down the group increasing charge separation between the metal cations of the lattice and the free delocalised electrons.

    • This weakens the electrical attractive bonding force and so less thermal KE is needed to weaken the lattice to the 'collapse point' i.e. melting.

    • BUT the situation is not as simple as might be expected, e.g. the metal ions do not always have the same crystal lattice packing arrangement.

  • Advanced Inorganic Chemistry Page Index and LinksThe electronegativity tends to decrease:

    • The electronegativity values are the lowest of the elements, but there is still a group trend.

    • They get lower because the effective nuclear attractive force on the outer electron charge decreases down the group.

    • You can explain it along the lines of the decreasing 1st IE argument (above), by merely changing the last part of the argument from 'easier to lose electron' to 'weaker attraction of electron charge'.

  • Formula patterns:

    • The general formulae are written in the summary tables in two ways

      •  'simple' format M2O or ionic formulae (M+)2O2– where M represents Li to Fr or Be to Ra.

    • Since all compounds can be considered ionic, most formulae needed are readily derived in principle by knowing the formula and charge of 10 ions!

    • All formulae are readily derived from equating the total positive charge of the cation with the total negative charge of the anion, and expressing the formula as the simplest whole number ratio.

    • The Group 7 halide ion, X, can be fluoride F, chloride Cl, bromide Br and iodide I.

    • The ethanoate ion is included as an illustration of carboxylic acid salts (RCOOH acid ==> RCOO in salt).

    • The oxides, hydroxides and carbonates and hydrogencarbonates are usually white ionic solids and the Gp1/2 salts listed in the 2nd table are usually white/colourless crystalline ionic solids.

      • The Group 2 hydrogencarbonates do not exist as stable solids and are only modestly stable in aqueous solution at room temperature. 

    • All relevant equations showing their formation and reactions are in subsequent sections.

s–block cation \ anion oxide hydroxide carbonate

hydrogencarbonate

O2– OH CO32– HCO3
formula derived from Group1 cation  M+ M2O MOH M2CO3 MHCO3
(M+)2O2– M+OH (M+)2CO32– M+HCO3
formula derived from Group 2 cation  M2+ MO M(OH)2 MCO3 M(HCO3)2
M2+O2– M2+(OH)2 M2+CO32– M2+(HCO3)2
  • Advanced Inorganic Chemistry Page Index and LinksThere are also Group 1 hydrogen sulphates (hydrogensulfates/hydrogensulphates) of formula MHSO4, i.e. half neutralised sulphuric acid which are rarely encountered but can be crystallised.

  • The oxidation numbers–states in the compounds listed in the two tables above and below:

    • +1 for metal cation of group 1, +2 for metal cation of group 2

    • oxygen –2, hydrogen +1, carbon +4 in table above and below

      • (except ethanoate, carbon oxidation numbers are awkward in organic compounds, leave em' in AS is my advice!)

    • halogens e.g. Cl –1, nitrogen +5, sulphur +6

    • you need to be able to analyse an anion to understand the relationship between the constituent oxidation states and the charge on the anion, sum of oxidation states = overall charge on ion e.g.

      • carbonate, CO32–, C is +4, 3 O's at –2, sum of ox. states = +4 and –6 = –2 = charge on anion

      • hydrogencarbonate, HCO3, H is +1, C is +4, 3 O's at –2, sum of ox. states = +1 and +4 and –6 = –1 = charge on anion

      • nitrate(V), NO3, N is +5, 3 O's at –2, sum of ox. states = +5 and –6 = –1 = charge on anion

      • sulphate(VI), SO42–, S is +6, 4 O's at –2, sum of ox. states = +6 and –8 = –2 = charge on anion

cation\anion halide nitrate(V) sulphate(VI) ethanoate
X NO3 SO42– CH3COO
formula derived from Group1 cation  M+ MX MNO3 M2SO42– CH3COOM
M+X M+NO3 (M+)2SO42– CH3COOM+
formula derived from Group 2 cation  M2+ MX2 M(NO3)2 MSO4 (CH3COO)2M
M2+(X)2 M2+(NO3)2 M2+SO42– (CH3COO)2M2+
  •   The three strong acids mentioned in reactions in sections 5. to 9. are ...

    • hydrochloric acid, HCl ==> chloride salts

    • nitric acid, HNO3 ==> nitrate salts

    • sulphuric (or sulfuric) acid ==> sulphate salts (or sulfates)

    • and the 4th acid is the weak organic carboxylic acid ethanoic acid*, CH3COOH ==> ethanoates

      • *old name 'acetic acid', and the salts were called 'acetates'


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


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