IB Chemistry - Stoichiometry

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Chemical formula are written to show the relative numbers of atoms or ions in the simplest formula unit of the compound. It helps to become familiar with the formulae of the common compounds of chemistry. This saves time during exams and helps in the general understanding of chemistry.

Syllabus reference

Structure 2.1.2 - The ionic bond is formed by electrostatic attractions between oppositely charged ions.

  • Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions.
  • Binary ionic compounds are named with the cation first, followed by the anion. The anion adopts the suffix “ide”.
  • Interconvert names and formulas of binary ionic compounds.

Guidance

  • The following polyatomic ions should be known by name and formula: ammonium NH4+, hydroxide OH–, nitrate NO3–, hydrogencarbonate HCO3–, carbonate CO32–, sulfate SO42–, phosphate PO43–.

Tools and links

  • Reactivity 3.2 - Why is the formation of an ionic compound from its elements a redox reaction?
  • AHL Structure 2.2 - How is formal charge used to predict the preferred structure of sulfate?
  • AHL Reactivity 3.1 - Polyatomic anions are conjugate bases of common acids. What is the relationship between their stability and the conjugate acid’s dissociation constant, Ka?

The formulae of the compounds reflect the valency rules described in section 3.22

Ionic compounds

Ionic compounds have positive ions arranged in a giant lattice with negative ions. Every positive ion is surrounded by negaitive ions and every negative ion is surrounded by positive ions. The overall structure has no charge, therefore the number of positive charges is exactly cancelled out by the number of negative charges. This is also the case in the simplest formula unit.

Sodium ion surrounded by chloride ions
Chloride ion surrounded by sodium ions
Name formula   name formula
metals with a valency of I+ (example sodium)
metals with a valency of III+ (example aluminium)
sodium chloride NaCl   aluminium fluoride AlF3
sodium nitrate NaNO3 aluminium nitrate Al(NO3)3
sodium sulfate Na2SO4 aluminium sulfate Al2(SO4)3
sodium sulfite (sulfate IV) Na2SO3 aluminium sulfite (sulfate IV) Al2(SO3)3
sodium hydroxide NaOH aluminium hydroxide Al(OH)3
sodium carbonate Na2CO3 aluminium carbonate Al2(CO3)3
sodium oxide Na2O aluminium oxide Al2O3
metals with a valency of II+ (example magnesium)
 ammonium compounds (the NH4+ ion has a valency I+)
magnesium chloride MgCl2   ammonium chloride NH4Cl
magnesium nitrate Mg(NO3)2 ammonium nitrate NH4NO3
magnesium sulfate MgSO4 ammonium sulfate (NH4)2SO4
magnesium hydroxide Mg(OH)2 ammonium hydroxide NH4OH
magnesium carbonate MgCO3 ammonium carbonate (NH4)2CO3
magnesium oxide MgO ammonium dichromate (NH4)2Cr2O7
       
other compounds of interest
potassium manganate(VII) KMnO4   lead(IV) oxide PbO2
potassium dichromate K2Cr2O7 manganese(IV) oxide MnO2
potassium chromate K2CrO4 sodium thiosulfate Na2S2O3

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Covalent compounds

In covalent compounds all of the atoms in each molecule are held to one another by bonds comprising electron pairs. In the case of double bonds there are two electron pairs involved in the bond. There are no full charges in covalent molecules unlike their ionic counterparts, but there may be partial charges caused by dipoles between atoms having different electronegativities, such as oxygen and hydrogen.

Name formula   name formula
organic compounds
inorganic covalent compounds
Methane CH4 carbon monoxide CO
Ethane CH3CH3 carbon dioxide CO2
Propane CH3CH2CH3 sulfur(IV) oxide SO2
Butane CH3CH2CH2CH3 sulfur(VI) oxide SO3
Ethene CH2=CH2

sulfur hexafluoride

 SF6
Propene CH2=CHCH3 ammonia NH3
Methanol CH3OH nitrogen monoxide NO
Ethanol CH3CH2OH nitrogen dioxide NO2
Methanal HCHO phosphorus trichloride PCl3
Ethanal CH3CHO phosphorus pentachloride PCl5
Propanone CH3COCH3 chlorine trifluoride ClF3
Methanoic acid HCOOH xenon tetrafluoride XeF4
Ethanoic acid CH3COOH xenon hexafluoride XeF6
Methyl methanoate HCOOCH3 hydrogen peroxide H2O2
Ethyl ethanoate CH3COOCH2CH3 hydrogen chloride HCl
Chloromethane CH3Cl sulfuric acid H2SO4
1,1-dibromoethane CH2BrCH2Br nitric acid HNO3
Aminoethane CH3CH2NH2 phosphoric acid H3PO4
Ethanamide CH3CONH2    
Benzene C6H6    
Benzenecarboxylic acid C6H5COOH    
Phenol C6H5OH    
Nitrobenzene C6H5NO2    
Phenylamine C6H5NH2    
Glucose C6H12O6      

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Hydrated compounds

Ionic compounds in which water molecules has been used to build the crystal lattice (water of crystallisation) are called 'hydrated' salts

When many substances are crystallised from aqueous solution, water molecules form part of the crystal lattice and become an integral part of the final crystal structure. These molecules are called 'water of crystallisation' and when the compound is weighed out they must be taken into account.

Example: Cobalt(II) chloride crystals contain two molecules of water for every cobalt ion.

The formula of the crystals must be written CoCl2.2H2O showing the two water molecules.

Any mass of cobalt chloride weighed out also contains the water molecules.

Example: Calculate the mass of copper sulfate pentahydrate, CuSO4.5H2O, that must be weighed out to prepare 1dm3 of 1 molar solution.

The solution contains 1 x 1 = 1 mole of solute.

Relative formula mass = 63.5 + 32 + (4 x 16) + [5 x (2 + 16)] = 249.5

Therefore 249.5g must be weighed out.

If a compound can be prepared with or without water of crystallisation, then the terms 'hydrated' and 'anhydrous' are used to differentiate between the two types of compound.

Water of crystallisation can often be 'driven off' by heating the hydrated crystal, giving the anhydrous salt.

CuSO4.5H2O CuSO4.H2O + 4H2O
blue crystals white powder

Some compounds form several different hydrated crystals.

Efflorescence

Certain hydrated compounds lose some of their water of crystallisation when left in the open air. This is known as efflorescence. A case in point is that of sodium carbonate decahydrate Na2CO3.10H2O. The crystals develop a powdery layer on the outside as water of crystallisation is lost to the atmosphere. Salts that effloresce cannot be used as standards for accurate preparation of solutions, as the exact composition of the crystals cannot be known.

Hygroscopy

Some compounds, typically ionic salts, absorb water from the atmosphere and increase in mass on exposure to air. Once again, the exact composition of the compound cannot be known and such salts cannot be used as primary standards. Sodium nitrate behaves in this way and, as such, is unsuitable for use in gunpowder, potassium nitrate being preferred instead.

Deliquescence

This is a special case of hygroscopy, where the salt ends up dissolving in the water absorbed from the air. A pellet of sodium hydroxide will turn into a small pool of sodium hydroxide solution when left in the air for long enough.

Table of some hydrated compounds (with water of crystallisation)
cobalt(II) chloride CoCl2.2H2O   dihydrate di = 2 water molecules
copper(II) sulfate CuSO4.5H2O pentahydrate pent = 5 water molecules
magnesium sulfate MgSO4.7H2O heptahydrate hepta = 7 water molecules
barium chloride BaCl2.2H2O dihydrate di = 2 water molecules
sodium carbonate Na2CO3.10H2O decahydrate deca = 10 water molecules
copper(II) chloride CuCl2.2H2O dihydrate di = 2 water molecules
copper(II) nitrate Cu(NO3)2.3H2O trihydrate tri = 3 water molecules
sodium thiosulfate Na2S2O3.5H2O pentahydrate penta = 5 water molecules
ColSol Testing

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