Survey of Period 2: Li across to Ne (8 elements, Z = 3 to 10)

4.3 Period 2 trends in bonding, structure, oxidation state, formulae and reactions

M⁺ X^- ionic bond, M^δ+-X^δ+ polar bond and M-X a relatively non-polar bond (no partial charges shown)

Sub-index for this page

(1) DATA TABLE for period 2 survey of element, oxides, chlorides and hydrides

(2) The structure and physical properties of period 2 elements

(3) Electron configuration and oxidation states of period 2 elements

(4) Reaction of element with oxygen and the structure of the oxide and its reaction with water, acid or alkali

(5) Reaction of element with chlorine and the structure of the chloride and its reaction with water

(6) Reaction of period 2 element with water

(7) The period 2 hydrides MH_x

(8) Radii of isoelectronic ions - both cations and anions overlapping periods 2 and 3

(1) DATA TABLE for period 2 survey of element, oxides, chlorides and hydrides of period 2

(2) The structure and physical properties of the elements of period 2

• The trend is metal lattice ==> giant covalent structure ==> small covalent molecules.

• Lithium and beryllium are silvery solids, with a metal lattice structure, and are good conductors of heat/electricity due to the delocalised free electrons moving between the immobile metal ions.

  • The melting/boiling points increase from Li ==> Be due to double the potential number of delocalised bond electrons that may contribute to bonding.

• Boron B_n and carbon C_n have a non-metallic giant covalent structure (where n is a very large number indeed) and are poor conductors of heat/electricity (though C_n(graphite) shows some metallic character as a 'moderate' conductor). The strong 3D bonding gives boron and carbon (diamond) very high melting/boiling points and great hardness.

  • Actually, the diamond/graphite allotropes of carbon sublime at temperatures over 3500^oC. Graphite with its huge 2D planes of atoms also has a very high melting/boiling point, though it is physically weak because the 2D layers of carbon atoms are only held by weak intermolecular forces and slip over each other under stress

• Nitrogen N₂, oxygen O₂ and fluorine F₂ have a simple-small covalent molecule structure and neon consists of single atoms. The molecules are only held together by the weakest of the intermolecular forces, namely the instantaneous dipole - induced dipole forces, and consequently have very low melting/boiling points.

  • From left to right the elements become less metallic and more non-metallic.

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(3) Electron configuration and oxidation states of the elements of period 2

• Electron configurations of  2,1 or 1s²2s¹ to 2,8 or 1s²2s²2p⁶  

  • Filling the s orbital (max 2 e-'s) gives the metallic s-block elements of Groups 1-2,

  • filling the p orbitals gives the predominantly non-metallic p block elements of Groups 3-7 & 0 (13-18) for Period 2.

• Oxidation states in compounds (numerically = valency) are: Li (+1 only), Be (+2 only), B (+3 only), C (usually +4^*), N (-3, +1 to +5), O (usually -2, but can be -1 and +2), F (-1 only), Ne has no stable compounds due to the full outer quantum level (shell) being full, conferring extra electronic stability on the atom.

  • From Li to N the maximum oxidation state is equal to the 'old' group number and the 'highest' oxide formulae can be predicted up to N and the chloride formula up to C (period 2 elements cannot exceed a co-ordination number of 4 due to valance orbital restrictions, in periods 3-4 the maximum is raised to 6.

    • So in the 'highest' oxides you can go from +1 to +5 for Groups 1 to 5, then drops to +2 for fluorine

    • Li₂O, BeO, B₂O₃, CO₂, N₂O₅ using all available 1-5 outer 2s and 2p valence electrons, lastly F₂O

    • and the chloride formula is derived from oxidation state +1 to +4 for Groups 1 to 4/14 (1-4 outer valency electrons), then declines +3, +2, -1

    • LiCl, BeCl₂, BCl₃, CCl₄ then NCl₃, Cl₂O (OCl₂), ClF (FCl)

  • ^* C can be (+2) in CO and from -1 to + 4 in organic compounds - a complex area in terms of oxidation state - but this knowledge is not needed for UK ASA2 level, if interested see Redox Reactions Part 3.

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(4) Reaction of period 2 element with oxygen and the structure of the oxide and its reaction with water, acid or alkali

• Reaction period 2 element with oxygen and the structure of the oxide

  • The metal Li burns in air/oxygen to form a giant ionic lattice oxide Li₂O or (Li⁺)₂O^2-

  • Beryllium forms a giant lattice oxide (BeO has an intermediate ionic-covalent structure).

  • Boron forms a giant covalent lattice of B₂O₃.

  • Carbon forms simple covalent molecular CO₂ gas (or CO).

  • Nitrogen forms a variety of simple molecular covalent gaseous oxides up to N₂O₅.

  • Fluorine can form F₂O gas.

  • In terms of the maximum oxidation state the observed formulae comply with +1 to + 5 for Groups 1-5, for fluorine it is -1

  • Li₂O, BeO, B₂O₃, CO₂, N₂O₅, (O₂), F₂O

  • The overall bonding pattern, from left to right is giant ionic lattice => giant covalent lattice => small covalent molecules.

  • The change in bonding character from ionic to covalent in the oxide, follows the decreasing difference in electronegativity between that of the period 2 element and oxygen.

  • -

• Reaction of the period 2 oxides with water, acids and alkalis

• The chemical character of the period 2 oxides - reaction with water, acids and alkalis

  • Lithium oxide Li₂O is basic and forming the alkali lithium hydroxide in water and this gives salts with acids.

    • Li₂O(s) + H₂O(l) ==> 2LiOH(aq)

    • LiOH_(aq) + HCl_(aq) ==> LiCl_(aq) + H₂O_(l)

  • Beryllium oxide BeO has no reaction but is amphoteric and forms salts with acids and alkalis.

    • BeO_(s) + 2HCl_(aq) ==> BeCl_2(aq) + H₂O_(l)

    • BeO_(s) + 2NaOH_(aq) + H₂O_(l) ==> Na₂[Be(OH)₄]_(aq) 

  • Boron oxide B₂O₃ is weakly acidic and reacts with strong alkalis to form borate salts.

  • Carbon dioxide CO₂ is weakly acidic forming salts with alkalis e.g.

    • CO_2(aq) + NaOH_(aq) ==> NaHCO₃(aq) + H₂O(l)

  • Nitrogen(I) oxide N₂O (dinitrogen oxide) and nitrogen(II) oxide NO (nitrogen monoxide) are neutral oxides,

    • but nitrogen(III) oxide N₂O₃, nitrogen(IV) oxide NO₂ (nitrogen dioxide) and nitrogen(V) oxide N₂O₅ are moderately to strongly acidic oxides.

    • Generally speaking, in a series of oxides of the same element, the higher the oxidation state of X in a 'X_xO_y' series, the more acidic is the oxide.

      • Nitrogen(V) oxide reacts with water to form nitric(V) acid,

      • N₂O₅(l) + H₂O(l) ==> 2HNO₃(aq)

      • and nitric(V) acid is neutralised by alkali to form salts e.g.

      • 2HNO₃(aq) + NaOH(aq) ==> 2NaNO₃(aq) + H₂O(l)

  • Oxygen(II) fluoride (oxygen difluoride) F₂O is also acidic in principal but I don't know if it reacts with water?

  • So, the general patterns across the period from left to right are

  • ionic lattice bonding ==> giant covalent structure ==> small covalent molecules

    • The change in bonding character from ionic to covalent in the oxide follows the decreasing difference in electronegativity between that of the element and oxygen.

  • In terms of chemical character ...

  • basic oxide => amphoteric oxide => weakly acidic oxide => strongly acidic oxide

  • and metal basic oxides ==> metal amphoteric oxides ==> non-metal oxides

  • This is chemically characteristic of metallic compound ==> non-metallic compound character.

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(5) Reaction of period 2 element with chlorine and the structure of the chloride and its reaction with water

• Reaction with chlorine and the structure of the chloride

  • The first three element will combine directly on heating in chlorine to give the expected chloride formula.

  • Lithium gives a giant ionic lattice of LiCl or Li⁺Cl^-,

  • Beryllium a polymeric covalent lattice of BeCl₂.

  • Boron forms small covalent molecules of BCl₃.

  • The others are also small covalent molecules, made indirectly, CCl₄, NCl₃, Cl₂O

  • In terms of the chloride the (oxidation states) in the chlorides are

  • (+1) LiCl, (+2) BeCl₂, (+3) BCl₃, (+4) CCl₄, (+3) NCl₃, (-2) Cl₂O, (+1) ClF, no neon chloride

  • So the number of atoms of chlorine combined with the Period 2 element (the valency) follows the pattern

    • 1  2  3  4  3  2  1  0

  • The overall pattern, from left to right across period 2 is ...

  • giant ionic lattice => polymeric covalent lattice ==> small covalent molecules.

  • The change in bonding character from ionic to covalent in the chloride, follows the decreasing difference in electronegativity between that of the element and chlorine, as in the case of oxides.

  • This is chemically characteristic of metallic ==> non-metallic element character.

• Reaction of the period 2 chlorides with water

• The reaction of chlorides with water

  • Ionic LiCl dissolves to form a nearly neutral solution of hydrated lithium and ions.

  • The covalent beryllium chloride BeCl₂ hydrolyses to give beryllium hydroxide and hydrochloric acid.

  • Boron trichloride BCl₃ similarly hydrolyse to form acid solutions.

  • Tetrachloromethane CCl₄ (carbon tetrachloride) is quite stable in water.

  • Nitrogen trichloride NCl₃ slowly hydrolyses to give an acid solution.

  • Chlorine(I) oxide Cl₂O hydrolyses to give an acidic solution.

  • FCl or ClF gives acidic solution ???

  • The general trend is for ionic chloride salts to give nearly neutral solutions => covalent chlorides that hydrolyse to give acidic solutions.

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(6) Reaction of period 2 element with water

• The reaction of the element with water

  • The reactive metal Li gives the hydroxide and hydrogen.

  • Beryllium and boron have no reaction with water.

  • Carbon reacts at high temperature.

  • Nitrogen and oxygen have no reaction with water.

  • Fluorine reacts violently forming oxygen and hydrogen fluoride.

    • Fluorine is such a powerful oxidising agent that it can oxidise water itself!

  • The 'limited' pattern for period 2 (or any other period), is to have reactive metals on the left forming an alkaline solution of a hydroxide and hydrogen and non-metals on the right forming acid solutions IF they react with water (only fluorine).

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(7) The period 2 hydrides MH_x

• A hydride is formed from the combination of an element with hydrogen.

• For hydrides the difference in electronegativity works both ways!

• From left to right across the period you change from an ionic lithium hydride crystal lattice e.g.

  • Li⁺H^- to small non-polar molecule covalent hydrides (methane CH₄)

  • and then a polar weakly basic covalent hydride molecule (ammonia NH₃)

  • and finally a quite acidic polar covalent molecule (hydrogen fluoride HF).

• The formulae follow a simple period pattern of rising and falling valency combinations.

  • LiH  BeH₂  B₂H₆ (= to BH₃)  CH₄  NH₃ H₂O  HF no neon hydride

  • valency combing power of the element gives a 1  2  3  4  3  2  1  0 pattern

• On reaction with water, the ionic metal hydrides at the start of the period give an alkaline solution

  • e.g. LiH(s) + H₂O(l) ==> LiOH(aq) + H₂(g)

    • MgH₂(s) + 2H₂O(l) ==> Mg(OH)₂(aq/s) + 2H₂(g)

• In the middle are neutral hydrides like methane which in contact with water do not change the pH.

• Then you get weak base ==> moderately strongly acidic hydrides when they dissolve in water e.g.

  • weak base: NH_3(aq) + H₂O_(l) <=> NH₄⁺_(aq) + OH^-_(aq)

  • quite strong acid: HF_(aq) + H₂O_(l) ==> H₃O⁺_(aq) + F^-_(aq)

• So things are a bit complicated with hydrides on period 2 and not the greatest of patterns!

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(8) Radii of isoelectronic ions - both cations and anions overlapping periods 2 and 3

• Isoelectronic means species having the same total number of electrons, hence the same electron configuration.

• The table below considers the isoelectronic anions associated with Periods 2 (and the table continues with Period 3 cations).

• isoelectronic system	Group 4/14	Group 5/15	Group 6/16	Group 7/17	(Group 0/18)	Group 1	Group 2	Group 3/13
  Period	Period 2					Period 3
  [Ne] 10e 1s22s22p6	C4-	N3-	O2-	F-	(Ne)	Na+	Mg2+	Al3+
  total nuclear charge	+6	+7	+8	+9	(+10)	+11	+12	+13
  radius in picometre (pm)	260	171	140	136	(38-112*)	95	65	50
  name of ion	carbide	nitride	oxide	fluoride	(neon)	sodium	magnesium	aluminium

• Excluding the noble gases themselves, there is a clear pattern of decreasing ionic radius with increase in nuclear charge (+ atomic/proton number) for the two isoelectronic series tabulated above.

• From left to right the proton/electron ratio is steadily increasing so that the electrons are experiencing an increasingly greater attractive force of the nucleus for the same number of electrons, hence the steady decrease in radii for an isoelectronic series.

• * all sorts of values are quoted for noble gas radii e.g. atomic, covalent and ionic, but most don't fit in the pattern above which is quite clear for all the cations and anions listed.

• More isoelectronic radii are at the end of the Period 3 survey pages or the end of the Period 4 survey pages.

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