Chemical Nomenclature (Naming Compounds)

1. Chemical Nomenclature (Naming Compounds) MOLECULAR COMPOUNDS / IONIC COMPOUNDS

2. Naming Inorganic Compounds • All compounds can be written as a full name or as a chemical formula. A formula provides a short hand method of writing the compound. • sodium chloride  NaCl(s)

3. Naming Binary Molecular Compounds 1) The full name of the first element in the compound is used. 2) The ending of the name of the second element in the compound is changed to “ide” (both names written in lower case) 3) A system of prefixes is used to indicate the number of each type of atom in the compound.

4. Naming Binary Molecular Compounds • 1 mono- 6 hexa- • 2 di- 7 hepta- • 3 tri- 8 octa- • 4 tetra- 9 nona- • 5 penta- 10 deca- GREEK PREFIX SYSTEM [ NONMETAL WITH NONMETAL ]

5. Naming Binary Molecular Compounds 4) The symbol for the elements are written in the same order as they appear in the name, followed by a subscript number to indicate the number of each atom in the compound

6. Examples • carbon dioxide CO2 • carbon monoxide CO • iodine tribromide IBr3 • silicon tetrahydride SiH4 • diphosphorus pentasulfide P2S5

7. Naming Binary Ionic Compounds • The chemical formula of an ionic compound represents the lowest whole number ratio of the ions in the crystal and it is called a formula unit.

8. Naming Binary Ionic Compounds 1) Use the full name of the positive ion first. 2) Put the name of the negative ion last and change the ending to “ide” 3) The name of an ionic compound does not use prefixes to show the number of ions present

9. Naming Binary Ionic Compounds 4) When writing formulas of binary ionic compounds the symbol for the elements are written in the same order as they appear in the name. Subscript numbers are used to indicate the ratio of the ions in the compound.

10. Naming Binary Ionic Compounds 5)The charges on the ions must balance in the chemical formula, since ionic compounds are electrically neutral.

11. Examples of ionic compounds • sodium chloride NaCl • magnesium chloride MgCl2 • sodium phosphide Na3P • silver sulfide Ag2S • zinc nitride Zn3N2

12. Criss-cross method • Al3+ and Cl- • AlCl3 • Cu2+ and O2- • CuO

13. Naming ionic compounds formed by transition elements • Some of the transition elements can form more than one common type of ion.

14. 26 Atomic # Fe3+ ion charge symbol iron (III) Stock name Fe2+ iron (II)

15. Stock System • The stock system uses Roman numerals after the name of the positive (metal) ion to indicate the charge on that ion. • Example: iron (II) oxide • iron (III) oxide

16. Stock system examples • Copper can be either Cu2+ or Cu+. In combination with oxygen each ion will form a different compound. • copper (II) oxide  CuO • copper (I) oxide  Cu2O

17. The Classical System • The classical system uses different suffixes(endings) to indicate the possible charges on the ion.

18. The Classical System • The ending “ic” indicates the higher of the two possible charges. • The ending “ous” indicates the lower of the two possible charges. • chromic refers to Cr3+ • chromous refers to Cr2+ • These endings are added to the name of the transition element

19. Examples • Write the formulas of chromic nitride and chromous sulfide. • CrN & CrS • Write the names for Fe3P2 and Mn3P4 • iron (II) phosphide • manganese (IV) phosphide

20. Polyatomic ions • Polyatomic ions consist of a group of atoms combined together which have an overall charge.

21. Naming polyatomic ions • Polyatomic ions, being negative, are always written last in the name of the formula, except for ammonium (+), it is written first.

22. Naming polyatomic ions • When writing the formula for compounds containing more than one of a polyatomic ion, the symbol for the ion must be written in brackets. • A subscript number is added to indicate the number of polyatomic ions present in the compound.

23. Examples of polyatomic ions • Ammonium sulfate NH4+ SO42- • (NH4)2SO4 • Tin (II) nitrate Sn2+ NO3- • Sn(NO3)2

24. Hydrated compounds • ionic compounds that have water molecules associated with their crystal structure. • anhydrous compound: • hydrates that lose water; usually a different colour.

25. Hydrated molecular formula copper sulfate pentahydrate copper sulfate

26. Naming Hydrates • When writing the formula for hydrated crystal of an ionic compound a dot is placed between the formula and the number of water molecules. • Greek prefixes used to indicate how many water molecules are present. • copper(II) sulfate pentahydrate • CuSO4●5H2O

27. Acids and Bases

28. Acids • Acids are compounds that contain hydrogen atoms covalently bonded to other atoms to form a molecular compound.

29. Bases • Bases are ionic compounds that are able to release hydroxide ions (OH-) in solution. • Bases react with acids to produce a neutral solution of salt and water. • HCl(aq) + NaOH(aq)―>H2O(l) + NaCl(aq)

30. Naming Acids • All acids will be considered as being dissolved in water, ie. In aqueous solution and this must always be indicated by placing the subscript (aq) after the acid formula. • Acids can be named in two ways. Both systems are equally acceptable to the IUPAC.

31. IUPAC system • The IUPAC system places the word aqueous in front of the name of the acid, named as if it were an ionic compound. • Example: HCl (g) is named as hydrogen chloride, but dissolved in water it becomes an acid called aqueous hydrogen chloride, and its formula is HCl (aq).

32. The classical system • The names of hydrogen compounds when named as ionic compounds will end in ide, ate or ite. • There are three naming rules based on each of these three possible endings.

33. Rule 1 • When the name of the negative ion ends in -ide the acid begins with the prefix hydro- and the stem of the negative ion is given the ending –ic, in place of –ide. This is followed by the word acid. • HCl(aq) would be named as hydrogen chloride and so becomes hydrochloricacid.

34. Examples • HF(aq) • hydroflouric acid • HBr(aq) • hydrobromic acid • HI(aq) • hydroiodic acid

35. Rule 2 • When the negative ion ends in –ate the acid name is the stem of the negative ion given the ending –ic , in place of –ate , followed by the word acid. • HClO3(aq) would be named hydrogen chlorate and so becomes chloricacid.

36. Examples • H2CO3(aq) • carbonic acid • HNO3(aq) • nitric acid

37. Rule 3 • When the negative ion ends in –ite the acid name is the stem of the negative ion given the ending –ous , in the place of –ite, followed by the word acid. • Example: HNO2(aq) would be named hydrogen nitrite and so becomes nitrousacid.

38. Examples • H2SO3(aq) • sulphurous acid • HClO2(aq) • chlorous acid

39. Writing names of acids • Writing names of the acids is simply the reverse of the rules above.

40. Exceptions to the rules • Acids which have sulf- or phosph- as their roots have extra letters added to make the name sound better. • Sulf- has -ur added so that it becomes sulfuric or sulfurous. • Phosph- has –or added to it to become phosphoric and phosphorous.

41. Exceptions to the rules • When the formula for the acid contains COO, the hydrogen atom is placed at the end of the formula after the COO. • ethanoic acid is written as CH3COOH • Oxalic acid is written H2OOCCOO or • HOOCCOOH / HO2C2O2H condensed • Benzoic acid HC6H5COO / C6H5COOH

42. Always remember, when writing the formula of an acid, to indicate that it is an acid by placing the subscript (aq) after the formula.