AP Chemistry Notes

1. AP Chemistry Notes Molecules, Ions, & Inorganic Nomenclature

2. Molecules and Molecular Compounds: • A molecule consists of two or more atoms bound together. H2 H2O NH3 CH4

3. Molecules & Chemical Formulas: • Each molecule has a chemical formula. • The chemical formula indicates: • Which atoms are found in the molecule, and • In what proportion they are found. • Compounds composed of molecules are molecular compounds. • These contain at least two types of atoms. • Different forms of a compound which have different chemical formulas are known as allotropes. • Allotropes differ in their chemical and physical properties.

4. Molecular & Empirical Formulas: • Molecular formulas – • Give the actual numbers and types of atoms in a molecule. • Examples: H2O2; C6H12O6; N2H4 • Empirical formulas – • Give the relative numbers and types of atoms in a molecule (they give the lowest whole-number ratio of atoms in a molecule). • Examples: HO; CH2O; NH2

5. Picturing Molecules: • Molecules occupy three-dimensional space. • However, we often represent them in two dimensions. • The structural formula gives the connectivity between individual atoms in the molecule. • If the structural formula does not show the shape of the molecule, then either a perspective drawing, ball-and-stick model, or space-filling model is used. • Perspective drawings use dashed lines and wedges to represent bonds receding and emerging from the plane of the paper. • Ball-and-stick models show atoms as contracted spheres and the bonds as sticks. The angles in the ball-and-stick model are accurate. • Space-filling models give an accurate representation of the relative sizes of the atoms and the 3D shape of the molecule.

6. Picturing Molecules:

7. Ions and Ionic Compounds: Remember… ELECTRONS are gained and lost … nothing happens to the protons! • If electrons are added or removed from a neutral atom, an ion is formed.

8. Ions and Ionic Compounds: • When an atom or molecule loses electrons, it becomes positively charged. • Positively charged ions are called cations. • When an atom or molecule gains electrons, it becomes negatively charged. • Negatively charged ions are called anions. • In general, metal atoms tend to lose electrons and nonmetal atoms gain electrons. • When molecules lose electrons, polyatomic ions are formed.

9. Predicting Ionic Charges: • An atom or molecule can lose more than one electron. • Many atoms gain or lose enough electrons to have the same number of electrons as the nearest noble gas (group 18). • The number of electrons an atom loses is related to its position on the periodic table.

10. Predicting Ionic Charges: • Group 1  +1 • Group 2  +2 • Group 13  +3 • Group 14  +4 / -4 • Group 15  -3 • Group 16  -2 • Group 17  -1 • Group 18  STABLE

11. Ionic Compounds: • A great deal of chemistry involves the transfer of electrons between species. • In general, ionic compounds are usually combinations of metals and nonmetals, whereas molecular compounds are generally composed of nonmetals only. • Writing formulas for ionic compounds: • You need to know the ions of which it is composed. • The formula must reflect the electrical neutrality of the compound. • You must combine cations and anions in a ratio so that the total positive charge is equal to the total negative charge.

12. Ionic Compounds:

13. Nomenclature: • Inorganic binary compounds • Ionic & covalent bonds • Compounds composed of two elements

14. Nomenclature: • Binary ionic compounds (type 1) • Cation written first • Monatomic cation named from element • Anion second • Monatomic anion named from the root of the element and changing the ending to -ide

15. Nomenclature: • Binary ionic compounds (type 2) • When cations form more than one type of positively charged ion, the charge on the ion is indicated by Roman numerals • Example – iron (III) & iron (II) • ALSO – for metals with only two ions – the ion with the higher charge ends in –ic and the ion with the lower charge ends in –ous • Example – ferric & ferrous (aka: iron)

16. Nomenclature:

17. Nomenclature: • Binary covalent compounds • First element in the formula is named first using the full element name • Second element is named like anions (root of element & -ide) • Prefixes are used to denote numbers of atoms present in the compound • Exception: mono- is not used with the first element

18. Nomenclature: • Prefixes: • Mono- = 1 • Di- = 2 • Tri- = 3 • Tetra- = 4 • Penta- = 5 • Hexa- = 6 • Hepta- = 7 • Octa- = 8 • Nona- = 9 • Deca- = 10

19. Nomenclature: • Another exception: • Some binary covalents are referred to by their common names NOT their systemic names • Examples: water, ammonia, phosphine…

20. Nomenclature:

21. Nomenclature: • Ionic compounds with polyatomic ions • Polyatomic names aren’t changed when naming compounds • Oxyanions can vary in the number of oxygen atoms composing the polyatomic ion • With variation - name changes BUT CHARGE DOES NOT CHANGE

22. Nomenclature: • When there are two possible oxyanions, most oxygen atoms ends in –ate & fewest oxygen atoms ends in –ite • When there are more than two possible oxyanions: • Most oxygen atoms also ends in –ate AND it is given the prefix per- • Fewest oxygen atoms also ends in –ite AND it is given the prefix hypo-

23. Nomenclature: • Examples: • Sulfate (SO42-) & sulfite (SO32-) • Perchlorate (ClO4-), chlorate (ClO3-), chlorite (ClO2-), & hypochlorite (ClO-)

24. Nomenclature: • Acids • Acids w/oxyanions • Acid name formed from the root of the anion name and the ending is –ic or –ous depending on whether the oxyanion is an –ate or an –ite • ate becomes –ic & -ite becomes -ous • Acids w/o oxyanions • The prefix hydro- is used and the ending of the anion is -ic

25. Nomenclature:

26. Mixed Practice • Dinitrogen monoxide • Potassium sulfide • Copper (II) nitrate • Dichlorine heptoxide • Chromium (III) sulfate • Ferric sulfite • Calcium oxide • Barium carbonate • Iodine monochloride • N2O • K2S • Cu(NO3)2 • Cl2O7 • Cr2(SO4)3 • Fe2(SO3)3 • CaO • BaCO3 • ICl

27. Mixed Practice • BaI2 • P4S3 • Ca(OH)2 • FeCO3 • Na2Cr2O7 • I2O5 • Cu(ClO4)2 • CS2 • B2Cl4 • Barium iodide • Tetraphosphorus trisulfide • Calcium hydroxide • Iron (II) carbonate • Sodium dichromate • Diiodine pentoxide • Cupric perchlorate • Carbon disulfide • Diboron tetrachloride

28. Nomenclature: • A HYDRATE is an salt that has water molecules trapped within its crystals. • Every hydrate has a certain number of water molecules associated with each formula unit of the ionic compound. The number of water molecules is indicated by using prefixes. • Mono- for one; di- for two; tri- for three; etc.

29. Hydrate Examples: • WRITING FORMULAS: copper(II) fluoride tetrahydrate CuF2·4 H2O calcium nitrate trihydrate Ca(NO3)2 ·3 H2O • WRITING NAMES: MgSO4· 7 H2O magnesium sulfate heptahydrate FeCl2 · H2O iron (II) chloride monohydrate

30. Nomenclature:

31. Nomenclature: • Complex Ion Formation: • IUPAC Rules for Naming Coordination Compounds – • 1. The cation is always named before the anion, with a space between the names. • 2. In naming a complex ion, the ligands are named before the metallic ion. • 3. Ligands have special endings. • Notice that an –o ending is used in place of any –ide ending for anions. • For anions ending in –ate, the –o is substituted for the –e. • For neutral ligands, the name of the molecules is used, with the exception of H2O, NH3, CO, and NO.

32. Nomenclature: • 4. Ligands are named first, in alphabetical order, followed by the metal name. • The Greek prefixes (di = 2, tri = 3, tetra = 4, penta = 5, hexa = 6, etc.) are used to denote the number of simple ligands present. • For more complicated ligands, the prefixes bis (twice), tris (thrice), tetrakis (four times), pentakis (five times), and hexakis (six times) are used. • NOTE: Prefixes do not affect the alphabetical order • 5. Many metals exhibit variable oxidation numbers. • The oxidation number of the metal is designated by a Roman numeral in parentheses following the name of the complex ion or molecule. • The Roman numeral is omitted if a metallic cation does not exhibit variable oxidation numbers. • Example: Silver is always +1 in a compound or complex ion, so no Roman numeral is used.

33. Nomenclature: • 6. For anionic complex ions (those having negative charges), the suffix –ate is added to the name of the metal. The English name for the metal is used, except in the following cases where the Latin name is used to identify the metal. The suffix –ate indicates that the complex ion is an anion. Neutral and cationic complexes do not use suffixes.

34. Nomenclature: • 7. In the case of complex-ion isomerism, the names cis or trans may precede the formula of the complex ion name to indicate the spatial arrangement of the ligands. • Cis means the ligands occupy adjacent coordination positions • Trans means opposite positions

35. Nomenclature: • Examples: • [Cr(H2O)4Cl2]+ = tetraaquodichlorochromium(III) • [Cr(H2O)2Cl4]- = diaquotetrachlorochromate(III) • K2[CoCl4] = potassium tetrachlorocobaltate(II) • Na[Al(OH)4] = sodium tetrahydroxoaluminate • [Zn(NH3)4]Cl2 = tetraamminezinc chloride • [Pt(NH3)4][PtCl6] = tetraammineplatinum(II) hexachloroplatinate(IV)