Naming IONS & formulas for Ionic Compounds

1. Sections 9.1 & 9.2 Naming IONS & formulas for Ionic Compounds

2. Type of Ion Formed • Remember that the representative elements behave in a predictable manner when determining the type of ion they will form. • The group # (1A-8A) tells how many valence electrons each has (except for helium).

3. Forming an Ion

4. Forming an Ion • Remember that ions are different in size from the atomic form of the element. This affects crystalline structure, but not ratios in formulas.

5. Review of Naming Cations • For metallic elements, naming cations that are formed is easy. The name of the ion is the same as the name of the atom, just followed by the word “ion” or “cation.”

6. Transition Metal Cations • Transition metals are not easy to predict. Several will readily form more than one type of ion. Roman numerals also are part of their names. • Others to know are silver (Ag+), gold (Au+), zinc (Zn2+) and cadmium (Cd2+).

7. Review: Naming Anions • The atom name endings (suffix) is changed to –ide, and it is followed by either “ion” or “anion.” • Fluorine atom becomes a fluoride ion, sulfur atom becomes sulfide ion, and phosphorus atom becomes phosphide ion. • Group 4A and 8A tend not to form ions.

8. Transition Metal Ion Use • The transition metal compounds are frequently very bright in color, so they are frequently used as pigments (to color other materials). • For example: • chromium compounds are used to make yellow, orange, red or green paints. • Cadmium compounds produce colors from yellow to red to maroon.

9. Practice Write the symbol (including the charge) for the ion formed by each element and then name the ion and classify it as an anion or a cation: Arsenic Beryllium Astatine Gallium

10. Mono- versus Poly- • All of these ions we’ve discussed are called monatomic ions – ions formed when one atom forms an ion. • Ions also exist that are polyatomic ions – these are tightly bound groups of atoms that behave as a unit and carry a charge. • Most polyatomic ions are anions and usually end in “ite” or “ate;” however, there are some exceptions to this rule.

11. Polyatomic Suffixes • The two exceptions included in the table are cyanide and hydroxide. • You may also encounter a couple of polyatomic cations, which end in “ium”: the ammonium ion (NH4+) and hydronium ion (H3O+). • You should study your ion flashcards again this chapter.

12. Polyatomic Suffixes • One thing you may notice about polyatomic anions ending in “ite” or “ate” is that they always refer to anions containing oxygen. • The “ite” suffix indicates that the ion has one less oxygen than a polyatomic ion with the “ate” ending (but it doesn’t tell you how many that is).

13. Similar Polyatomics • Notice also that some of these ions are similar, they just have an H as part of the formula – think of this as having a hydrogen ion (H+) combined with the polyatomic ion (notice that the anionic charge is one less when H is part of the formula – from adding a +1 to the charge). • Look at PO43-, HPO42- and H2PO4-.

14. Classical Naming • One other thing you should be aware if is that initially their naming system differed from the system we use today. • In the “classical” system, the elements that had more than one possible charge were named using “ous” or “ic” as suffixes, with “ous” being used on the cation with the lower charge, as shown in the table:

15. Naming Compounds • Scientists used to make up their own names for compounds so its name described its properties, source or use: • Baking soda – used to make cakes rise when baked • Plaster of paris – plaster used to make face masks

16. Naming Systematically • Antoine Lavoisier and other chemists decided to come up with a systematic method for naming compounds (so they wouldn’t have to memorize all the names). • The easiest ones to name are binary ionic compounds – these consist of two types of ions joined together: the name of the compound is just the combination of the ion names, with the cation name always given first.

17. Naming Systematically • Before we practice naming binary ionic compounds, first recall that ions will combine in a way that the result has no charge (the positive and negative charges will balance). Sn4+ bound to F- would be: SnF4 This is called: tin(IV) fluoride

18. Naming Systematically • Also, you should know that if you wish to indicate that there are 2 or more units of a polyatomic ion in one formula unit of the compound, use parentheses around the polyatomic ion formula. • It would take 2 nitrate cations to balance with one calcium ion in calcium nitrate: Ca2+ bound to NO3- would be: Ca(NO3)2

19. Practice Naming Binary Ionic: NaF MgCl2 Cs2O FeCl2 MnF3 CaSO4

20. Practice Naming Binary Ionic: Na2CO3 NaHCO3 Cu(NO2)2 KC2H3O2 Zn(OH)2NH4Cl

21. Taking it the other way: • If you know the name of a binary ionic compound, you can write the formula – just remember to make sure that your formula has the proper number of each type of ion to insure that the charges add to give zero. • Start with the formulas of the ions involved along with their charges. Then “crisscross” the charges to get the subscripts – finish by making sure the ions are in a simplified ratio in the formula.

22. Using the Crisscross Method • Iron(III) oxide • Iron(II) oxide • Cobalt(II) phosphate • Silver sulfite • Calcium dihydrogen phosphate

23. Practice Write formulas for compounds formed from these pairs of ions, then name the compounds: Ba2+, S2- Li+, O2- Ca2+, CN- NH4+, I-

24. Practice Write formulas for these compounds: Sodium iodide Stannous chloride Potassium sulfide Plumbic fluoride Lithium hydrogen sulfate Chromium(III) nitrite

25. Review: • Three things must be considered when naming ionic compounds: • The identity of ions • The order of the names (cation first) • The possibility that an element may form cations with more than one charge (so you must use roman numerals) • Name FeN, Mg(NO3)2and (NH4)2C2O4.

26. References • http://www.calzim.com/online/online2_1/class_material/unit1/ion.gif • http://www.roymech.co.uk/images14/lewis_elements.gif • http://www.middleschoolchemistry.com/img/content/multimedia/chapter_4/lesson_6/lewis_dot_table.jpg • http://www.kanescience.com/_images/chem_ionic/sodium_ions.png • http://www.calzim.com/online/online2_1/class_material/unit1/unit1.htm • http://www.personal.kent.edu/~cearley/ChemWrld/compounds/ionictable.gif • http://www.chem.uwec.edu/Chem103_F08_F0F/pages/resources/media/transitions_metal_ions_Silberberg_table_2.4.jpg • http://wps.prenhall.com/wps/media/objects/1053/1078773/tools/HPMP.table.2.4.gif • http://www.behsscience.com/apchem/chapter7/ion_size.jpg • http://www.chemicalregister.com/storefront/cr/3099/f-Pigments.gif • http://3.bp.blogspot.com/-tM9zAA_mvZ0/Td-V4G0HG_I/AAAAAAAAAXE/x_Zx-i-OICo/s1600/baking-795089.JPG • http://images.melissaesplin.com/wp-content/uploads/2009/03/_dsc1749sm.jpg • http://reich-chemistry.wikispaces.com/file/view/HSantoin.jpg/44989145/HSantoin.jpg