Doc Brown's Chemistry Revision

Extra Notes on chemical bonding for advanced A level chemistry students

 Extra A level Notes on Ionic Bonding and Ionic Compounds

All the advanced A level 'basics' with lots of examples of dot and cross diagrams of ionic bonding, Lewis diagrams, properties of ionic compounds etc. is on a separate page.

IONIC BONDING

However, advanced electron notation using s, p and d notation was NOT included.


Ionic bonding and advanced electron notation

I'm not going to repeat the diagrams, there's no point, so I'm just outlining the ionic bond formation in s, p and d terms.

Most involve the 'octet rule' to form a noble gas electron configuration.

Examples

You need to be able to write the electron configuration of ions in terms of s, p and d orbital notation.

e.g. sodium ion Na+ is 1s22s2p6, [Ne], and the chloride ion Cl- is 1s22s2p63s23p6, [Ar]

More on electron configuration of ions and oxidation states


Which elements form ionic compounds? and how to work out and write an ionic formula?

Pd metals Part of the modern Periodic Table

related to selected elements that form ionic compounds

Pd = period, Gp = group

metals => non–metals
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0
1

1H  Note that H does not readily fit into any group

2He
2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn
Gp 1 Alkali Metals  Gp 2 Alkaline Earth Metals  Gp 7/17 Halogens  Gp 0/18 Noble Gases

Chemical bonding comments about the selected elements highlighted in white

e.g. When the electropositive metals on the left combine with the non–metals on the right, quite often ionic bond is formed e.g. the formation of an ionic compound like sodium chloride NaCl

Note: Throughout this page to form stable ions with a noble gas electron arrangement by electron transfer ...

(a) Group 1 metals lose their 1 outer electron to form a singly charged positive ion: M ==> M+ + e

(b) Group 2 metals lose their 2 outer electrons to form a doubly charged positive ion: M ==> M2+ + 2e

(c) Group 6 non–metals gain 2 electrons to form a doubly charged negative ion: X + 2e ==> X2–

(d) Group 7 halogen non–metals gain 1 electron to form a singly charged negative ion: X + e ==> X

In a correct ionic formula: total positive ion charge = total negative ion charge

therefore we can predict the following formula where  M = a group 1/2 metal  and  X = a group 6/7 non–metal:

(a) Group 1  +  (c) Group 6  ===>  M2X  or  (M+)2X2–   (eg group 1 oxides or sulfides)

(a) Group 1  + (d) Group 7  ===>  MX  or  M+X   (eg group 1 halides)

(b) Group 2  +  (c) Group 6  ===>  MX  or  M2+X2–   (eg group 2 oxides or sulfides)

(b) Group 2  +  (d) Group 7  ===>  MX2  or  M2+(X)2   (eg group 2 halides)


Atoms of groups 4/14 and 5/15 may also form ions in combination with very electropositive metals

examples – showing the formation of a noble gas structure

Group 4/14: carbon forms ionic carbides: C (2.4) + 4e ==> C4– (carbide ion, 2.8)

examples: sodium carbide Na4C, magnesium carbide Mg2C

Group 5/15: nitrogen forms ionic nitrides: N (2.5) + 3 ==> N3– (nitride ion, 2.8)

Group 5/15: phosphorus forms ionic phosphides: P (2.8.5) + 3– ==> P3– (nitride ion, 2.8)

examples: potassium nitride K3N, magnesium nitride Mg2N3, sodium phosphide Na3P

formulae derived from Na+, K+, Mg2+ (see next section on working out ionic formulae)

Note that the

3d block – transition metals also form many simple ionic compounds with the more electronegative non–metals.

 

Examples of how to work out an ionic formula

numerical ion charges = the valency of A and B to deduce the formula AxBy

i.e. the valence or ionic charge = the combining power of the ion

'molecular' or ionic style of formula and compound name are shown

In the electrically balanced formula for a potentially stable compound, the total positive ionic charge must equal the total negative ionic charge.

number of positive ion 'A' x positive charge of ion 'A' = number of negative ion 'B' x negative charge of ion 'B' (you ignore charge sign)

Example: A moderately difficult example to work out!

Aluminium oxide consists of aluminium ions Al3+ and oxide ions O2– 

number of Al3+ x charge on Al3+ = number of  O2– x charge on O2– 

the simplest whole number ratios are 2 of Al3+ x 3 = 3 of  O2– x 2 (total 6+ balances total 6–)

so the simplest whole number formula for aluminium oxide is Al2O3 which is its empirical formula

More examples of how to work out ionic formulae

1 of K+ balances 1 of Br because 1 x 1 = 1 x 1 gives KBr or K+Br    potassium bromide

2 of Na+ balances 1 of O2– because 2 x 1 = 1 x 2 gives Na2O or (Na+)2O2–    sodium oxide

1 of Mg2+ balances 2 of Cl because 1 x 2 = 2 x 1 gives MgCl2 or Mg2+(Cl)2    magnesium chloride

1 of Fe3+ balances 3 of F because 1 x 3 = 3 x 1 gives FeF3 or   Fe3+(F)3    iron(III) fluoride

1 of Ca2+ balances 2 of NO3 because 1 x 2 = 2 x 1 gives Ca(NO3)2 or Ca2+(NO3)2     calcium nitrate

2 of Fe3+ balances 3 of SO42– because 2 x 3 = 3 x 2 gives Fe2(SO4)3 or (Fe3+)2(SO42–)3    iron(III) sulfate


LINK to Table of lots of common formulae of ionic compounds

(includes common oxides, hydroxides, carbonates, hydrogencarbonates, halides, sulfates, nitrates)

At Advanced A level you would be expected to work out any ionic formulae from given ions and you should know many of them anyway, or work out the charge on the ion from the position of the element in the periodic table and using the octet electron rule eg for groups 1, 2, 3/13, 4/14, 5/15, 6/16, 7/17, 0/18.


Naming ionic inorganic compounds

When combined with other elements in simple compounds the name of the non-metallic element changes slightly from ...??? to ...ide.

Sulfur forms a sulfide (ion S2-), oxygen forms an oxide (ion O2-), fluorine forms a fluoride (ion F-), chlorine forms a chloride (ion Cl-), bromine a bromide (ion Br-) and iodine an iodide (ion I-).

The other element at the start of the compound name e.g. hydrogen or a metal like sodium, potassium, magnesium, calcium, etc. remains unchanged because there is only one oxidation state.

So typical compound names are, sodium sulfide, hydrogen sulfide, magnesium oxide, potassium fluoride, hydrogen chloride, sodium chloride, calcium bromide, magnesium iodide etc.

However, with different oxidation states the complications will arise e.g.

(i) Where an element can form two different compounds with different formulae with the same  element there needs to be a way of expressing it in the name as well as in the formula e.g.

iron(II) chloride, FeCl2 and iron(III) chloride, FeCl3

copper(I) oxide, Cu2O and copper(II) oxide, CuO

Hear chlorine has a combining power of 1 (valence 1) and oxygen 2 in both compounds.

However, iron can have a valence of 2 or 3 and copper 1 or 2 and these also correspond numerically to the charge on the metal ions in such compounds e.g. Fe2+ and Fe3+, Cu+ and Cu2+.

Therefore the 'Roman numerals' number in (brackets) gives the valence of the element in that particular compound. At a higher academic level this is known as the oxidation state.

(ii) When the non-metal is combined with oxygen to form a negative ion (anion) ion which combines with a positive ion (cation) from hydrogen or a metal, then the end of the 2nd part of the name ends in ...ate or ...ite e.g.

NO3 in a compound formula is nitrate e.g. KNO3, potassium nitrate.

SO3 in a formula is sulphite, e.g. Na2SO3, sodium sulphite,

SO4 is sulfate, e.g. MgSO4, magnesium sulfate,

PO4 is phosphate, e.g. Na2HPO4, disodium hydrogen phosphate

  • Other examples

    For metallic or non–metallic elements the name of the element is used if NOT in an anion

    • Some of the old names are still in common use, but try to use the correct systematic name e.g.

    • copper(I) oxide Cu2O and copper(II) oxide CuO

      • (once called cuprous oxide and cupric oxide)

    • iron(II) chloride FeCl2 and iron(III) chloride FeCl3

      • (once called ferrous chloride and ferric chloride)

    • iron(II) oxide FeO, iron(III) oxide Fe2O3 and diiron(II) iron(III) oxide, Fe3O4

      • (once called ferrous oxide and ferric oxide and tri–iron tetroxide)

      • Historic note: ...ous was the lower oxidation state, ...ic the higher.

    • vanadium(II) sulfate for VSO4 or V2+SO42– and vanadium(III) sulfate V2(SO4)3

    • sulfur(IV) oxide SO2 (sulfur dioxide) and sulfur(VI) oxide, SO3 (sulfur trioxide)

    • nitrogen(I) oxide N2O (dinitrogen oxide) nitrogen(II) oxide NO (nitrogen monoxide), nitrogen(IV) oxide NO2 (nitrogen dioxide) and nitrogen(V) oxide N2O5 (nitrogen pentoxide).

    • transition metal complex cations e.g.

      • diaquatetraamminecopper(II) ion, [Cu(H2O)2(NH3)4]2+

        • (water and ammonia are electrically neutral ligands attached to the central Cu2+ ion)

  • For elements (metal or non–metal) combined with oxygen or other more electronegative element, giving an anion, the ion name ends in ...ate with the prefix derived from the elements name. In such cases the oxygen carries the negative oxidation state of (–2) or chlorine (–1) e.g.

    • vanadate(V) ion, VO43–,

    • manganate(VI) ion, MnO42–, manganate(VII) ion, MnO4, (was called the permanganate ion)

    • sulfate(IV) ion, SO32– (sulphite) and sulfate(VI) ion,SO42– (sulfate)

    • nitrate(III), NO2 (nitrite) and nitrate(V), NO3 (nitrate)

      • Historic note: ...ite was the lower oxidation state, ...ate the higher.

    • chlorate(I), ClO, chlorate(VII), ClO4 etc. oxygen is more electronegative than chlorine.

      • (once called the hypochlorite ion and the perchlorate ion respectively)

    • transition metal complex anions e.g.

      • tetrachlorocuprate(II) ion, [CuCl4]2– (oxidation states Cu +2, Cl –1)


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