Doc Brown's Chemistry  Periodic Table Revision Notes

Part 7. s-block Groups 1/2 Alkali Metals/Alkaline Earth Metals - Sections 7.1 to 7.4

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

Revision notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters Chemistry CIE Chemistry revising courses for pre-university students (equal to US grade 11 and grade 12 and Honours/honors level courses) GCSE/IGCSE notes Alkali Metals * GCSE Periodic Table notes * A level Quiz on basics of s-block chemistry * EMAIL query?comment

INORGANIC Part 7 s-block Gp 1 Alkali Metals/Gp 2 Alkaline Earth Metals  sub-index: 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 * 7.5 Oxygen reaction & oxides * 7.6 Water reaction & hydroxides * 7.7 Acid reaction & salts * 7.8 chlorine reaction - halides * 7.9 carbonates & hydrogen carbonates * 7.10 Solubility trends of OH, NO₃,SO₄,CO₃'s * 7.11 Thermal decomposition and stability of carbonates & nitrates * 7.12 Uses of group 1-2 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, 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 relatively their low ionisation energies, is dominated by the loss of these s electrons to form a cation.

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

  • the s² electron loss by the Group 2 Alkaline Earth Metals forms the M²⁺ ion.

• The principal (only!) oxidation states of +1 for Group 1 and +2 for Group 2 elements are governed by the relative ease of loss of the outer s electrons and the subsequent very high ionisation energies required to remove a 2nd or 3rd electron respectively.

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

• Consequently most of the compounds of Group 1-2 elements tend to be ionic in nature.

• For introduction-revision GCSE notes on Alkali Metals and GCSE Quiz on Alkali Metals

• and an ASA2 A level Quiz on the basics of s-block metal chemistry.

7.2. Information and Data Table GROUP 1 ALKALI METALS

• Some of the data is tabulated and plotted on a separate web-page in a new window.

• 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!

• Untypical in other ways e.g. relatively soft, low density (Li-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.

    • Note: Because the set of quantum level energies are unique for each atom, it means the quantum level difference varies from atom to atom, therefore the frequency of emitted photons is different, hence you see different flame colours in the visible region of light from each Group 1/2 metal.

      • 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.

• 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.

    • Detailed notes on oxidation state and redox reaction theory.

  • 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.

  • and for help notes on data, see the inorganic data page. (opens in new window)

• PLEASE 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.

7.3. Information and Data Table GROUP 2 ALKALINE EARTH METALS

• Very typical metals, silvery grey lustrous solids, relatively high melting and boiling points, good conductors of heat and electricity.

• 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.

  • and for help notes on data, see the inorganic data page (opens in new window)

• 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.

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. The 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:

  • 

  • Because 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 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)

• 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.

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

  • Also consider the increasingly negative half-cell potentials (EM/M+ and EM/M2+) down the groups, i.e. increasing potential to acts as a reducing agent.

• The 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 'simple' format M₂O or ionic formulae (M⁺)₂O^2- 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 derive from these 10 ions by 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) that some GCSE courses introduce.

  • 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.

• There are also Group 1 hydrogen sulphates 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, CO₃^2-, 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), SO₄^2-, S is +6, 4 O's at -2, sum of ox. states = +6 and -8 = -2 = charge on anion

•   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'

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