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Section 3.0

Learn the rules and guidelines for naming ionic compounds and polyatomic ions. Understand how to indicate the charge of ions and differentiate between different compounds.

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Section 3.0

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  1. Section 3.0 Compounds Form According to a Set of Rules NaCl C12H22O11

  2. Section 3.1 Naming Compounds • A compound occurs when two or more elements combine chemically to produce a new substance. • Each compound has its own chemical name and formula. • Chemical formula – identifies which elements and how many of each are in a compound

  3. For example: • water’s chemical formula is H2O • water has 2 hydrogen atoms and 1 oxygen. • The number 2 is a subscript.

  4. Nomenclature: • In 1787, Guyton de Morveau created a naming system for chemicals. • The metallic element name is written first followed by the non-metallic element. • Since 1920, the International Union of Pure and Applied Chemistry or IUPAC governs the naming of compounds.

  5. A compound’s physical state at room temperature is always given in brackets after the chemical formula: (g) = gas CH4(g) (l) = liquid H2O (l) (s) = solid NaCl (s) • For substances that can be dissolved in water, the subscript (aq) aqueous is used.

  6. Section 3.2 Ionic Compounds • An ionic compound is a pure substance formed by attraction between particles of opposite charges, called ions. • Ions are simply Atoms that have a charge. • properties of ionic substances: • high melting point • electrical conductivity when dissolved in water • distinct crystal shape • solids at room temperature

  7. When ionic substances dissolve in water, the metallic and non-metallic elements separate to become ions (electrically charged particles due to the gain or loss of electrons). • This allows ionic solutions to conduct electricity.

  8. Ion charges • Ion charges are shown by a superscript of either a plus sign or a minus sign. Cations (+): • Positive ions, called cations, have lost electrons • Metals tend to form cations. • If 1 electron is lost, the cation has a 1+ charge, if 2 electrons are lost, the cation has a 2+ charge, and so on. Anions (-): • Negative ions, anions, have gained electrons. Non-metals tend to form anions.

  9. How do you remember which ions are positive?

  10. How do you remember which ions are positive? • Just remember, “Cats have Pos”!

  11. By looking at the periodic table, a pattern in ion charges can be seen. Generally, all the elements in a group have the same charge. (See Fig 3.9, p. 147).

  12. Naming Ionic Compounds • Use the full name of the metal (cation) ion first • Put the name of the non-metal (anion) last and change the ending to –ide • Exception: if the anion is a polyatomic ion, its name remains unchanged • If the cation has more than 1 possible charge, indicate which ion is being used with roman numerals • eg. iron (III) oxide is a compound containing Fe3+

  13. Name the Following Ionic Compounds: • NaCl (s) • LiBr (s) • CaO (s) • Sr2Cl (s) • BaF2 (s) • K2S (s)

  14. Name the Following Ionic Compounds: • NaCl (s) sodium chloride • LiBr (s) lithium bromide • CaO (s) calcium oxide • Sr2Cl (s) strontium chloride • BaF2 (s) barium fluoride • K2S (s) potassium sulfide

  15. Polyatomic ions are groups of atoms that when together have a charge. (Some examples are on the next slide) • You will find the polyatomic ions on your periodic table.

  16. Polyatomic Ions (these groups of atoms tend to combine together but still have a charge) • Ammonium NH4+1 • Bicarbonate HCO3-1 • Carbonate CO3-2 • Chlorate CIO3-1 • Chromate CrO4-2 • Dichromate Cr2O7-2 • Hydroxide OH-1 • Nitrate NO3-1 • Nitrite NO2-1 • Permanganate MnO4-1 • Phosphate PO4-3 • Phosphite PO3-3 • Sulphate SO4-2 • Sulphite SO3-2

  17. Now name these Ionic Compounds: (these ones contain polyatomic ions) • Li2CO3 • KClO2 • CaSO4 • Ba(NO2)2

  18. Now name these Ionic Compounds: (these ones contain polyatomic ions) • Li2CO3 lithium carbonate • KClO2 potassium chlorite • CaSO4 calcium sulfate • Ba(NO2)2 barium nitrite

  19. If the cation can have more than one possible charge, you must specify which ion is being used. Example: Name Fe2O3

  20. If the cation can have more than one possible charge, you must specify which ion is being used. Example: Name Fe2O3 Fe3+ and Fe2+ are both possible ions of iron

  21. If the cation can have more than one possible charge, you must specify which ion is being used. Example: Name Fe2O3 Fe3+ and Fe2+ are both possible ions of iron Fe3+ + O2- = Fe2O3 Iron (III) ioxide

  22. If the cation can have more than one possible charge, you must specify which ion is being used. Example: Name Fe2O3 Fe3+ and Fe2+ are both possible ions of iron Fe3+ + O2- = Fe2O3 Iron (III) ioxide Fe2+ + O2- = FeOIron (II) ioxide

  23. If the cation can have more than one possible charge, you must specify which ion is being used. Example: Name Fe2O3 Fe3+ and Fe2+ are both possible ions of iron Fe3+ + O2- = Fe2O3 Iron (III) ioxide Fe2+ + O2- = FeOIron (II) ioxide

  24. Name the following compounds: HgF2 NiBr3 PbS2

  25. Name the following compounds: HgF2 Hg2+ Hg1+ F- NiBr3 PbS2

  26. Name the following compounds: HgF2mercury (II) fluoride Hg2+ Hg1+ F- NiBr3 PbS2

  27. Name the following compounds: HgF2mercury (II) fluoride Hg2+ Hg1+ F- NiBr3 nickel (III) bromide Ni2+ Ni3+ Br- PbS2

  28. Name the following compounds: HgF2mercury (II) fluoride Hg2+ Hg1+ F- NiBr3 nickel (III) bromide Ni2+ Ni3+ Br- PbS2lead (IV) sulfide Pb2+ Pb4+ S2-

  29. Using Ion Charges and Chemical Names to Write Formulas • Write the metal element symbol with its charge, next to it write the non-metal element symbol with its charge • Balance the ion charges so the net result is a charge of zero • Write the formula indicating how many atoms of each element are in it with a subscript. • If there is only 1 atom of an element, no subscript is needed

  30. For example: barium chloride Ba2+ Cl-

  31. For example: barium chloride Ba2+ Cl- How many of each ion do you need to balance the charges?

  32. For example: barium chloride Ba2+ Cl- How many of each ion do you need to balance the charges? Ba2+ Cl- Cl-

  33. For example: barium chloride Ba2+ Cl- How many of each ion do you need to balance the charges? Ba2+ Cl- Cl- BaCl2

  34. Write the formulas for the following Ionic Compounds: • Potassium chloride • Calcium chloride • Iridium oxide • Zirconium nitride • Cobalt (II) chloride

  35. Write the formulas for the following Ionic Compounds: • Potassium chloride K+ Cl- KCl (s) • Calcium chloride Ca2+ Cl- CaCl2 (s) • Iridium oxide Ir4+ O2- IrO2 (s) • Zirconium nitride Zr4+ N3- Zr3N4 (s) • Cobalt (II) chloride Co2+ Cl1- CoCl2 (s)

  36. Section 3.3 Molecular Compounds • Molecular compounds, or molecules, are formed when non-metals combine. • Some properties of molecular compounds are: - low melting and boiling points (forces between molecules are weaker) - poor conductors of electricity, good insulators - can be solids, liquids or gas at room temperature

  37. Naming Molecular Compounds • Many molecules are known by their common names, such as water, H2O, and ammonia, NH3. Others are named as follows:

  38. Use the full name of the first element (the most metal-like goes first) • Put the name of the second element last and change the ending to –ide • Use the correct prefix to indicate the number of each element • Exception: do not use the prefix mono when the first element only has 1 atom

  39. Number of AtomsPrefix • mono • di • tri • tetra • penta • hexa • hepta • octa • nona • deca

  40. So: To name molecules . . . Prefix + First Element, Prefix + Second Element (with –ide ending)

  41. Name the following molecular compounds: • CO2 • NO3 • N2O • NF3 • N2O3 • CO

  42. Name the following molecular compounds: • CO2 carbon dioxide • NO3 nitrogen trioxide • N2O dinitrogen monoxide • NF3 nitrogen trifluoride • N2O3 dinitrogen trioxide • CO carbon monoxide

  43. Remember: The number in subscript tells us how many of each element are in the compound.

  44. Comparing Ionic and Molecular Compounds • Ionic Compounds tend to have high melting and boiling points are therefore usually solids at room temp. • Ionic solutions conduct electricity. • Molecular compounds tend to have low melting and boiling points are therefore usually gases or liquids at room temp. • Molecular solutions do not conduct electricity

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