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Naming inorganic compounds

Naming inorganic compounds. Two principal groups:. Naming inorganic compounds. Two principal groups: (1) metal with nonmetal (or nonmetal group) use Stock system. Naming inorganic compounds. Two principal groups: (1) metal with nonmetal (or nonmetal group)

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Naming inorganic compounds

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  1. Naming inorganic compounds Two principal groups:

  2. Naming inorganic compounds Two principal groups: (1) metal with nonmetal (or nonmetal group) use Stock system

  3. Naming inorganic compounds Two principal groups: (1) metal with nonmetal (or nonmetal group) use Stock system (2) nonmetal with nonmetal use prefix system

  4. Stock system

  5. Stock system Oxidation number: is the charge that an atom in a compound would have if the electrons in each bond belonged entirely to the more electronegative atom.

  6. Stock system Oxidation number: is the charge that an atom in a compound would have if the electrons in each bond belonged entirely to the more electronegative atom. Electronegativity: relative attraction of an atom for electrons.

  7. Oxidation numbers: In most cases the oxidation number of an element in an ionic compound is the same as the formal charge on the ions present. Example: in the ionic compound KCl, which is composed of K+ ions and Cl- ions, the oxidation number of potassium is +1 and the oxidation number of the chlorine is -1.

  8. For a neutral compound the sum of the oxidation numbers of the elements = 0. Example: for KMnO4 for which the oxidation numbers are K (+1), Mn (+7), and O (-2), the sum of the oxidation numbers = 1 + 7 + 4x(-2) = 0.

  9. For a cation or an anion, the sum of the oxidation numbers of the elements = the charge on the species. Example: For MnO4- the sum of the oxidation numbers = 7 + 4x(-2) = -1, where -1 is the charge on the anion.

  10. The cation is named first.

  11. The cation is named first. • The oxidation number is given next (using Roman numerals in parentheses) – if it is needed.

  12. The cation is named first. • The oxidation number is given next (using Roman numerals in parentheses) – if it is needed. The oxidation number is given only if the element commonly has more than one oxidation state.

  13. The cation is named first. • The oxidation number is given next (using Roman numerals in parentheses) – if it is needed. The oxidation number is given only if the element commonly has more than one oxidation state. • The anion is named second.

  14. Examples: FeCl3 Iron forms two common cations, Fe2+ and Fe3+ , hence it will be necessary to specify the oxidation number of the Fe. The name is iron (III) chloride

  15. Examples: FeCl3 Iron forms two common cations, Fe2+ and Fe3+ , hence it will be necessary to specify the oxidation number of the Fe. The name is iron (III) chloride CoPO4 Cobalt forms two common cations, Co2+ and Co3+ , hence it will be necessary to specify the oxidation number of the Co. The name is cobalt (III) phosphate

  16. LiClO3 The cation has only one common oxidation state (which is +1), so it is not necessary to give this as part of the name. The name is lithium chlorate You need to be able to name in both directions: formula name name formula

  17. Give the formula for the following: Iron (II) phosphate The two ions are Fe2+ and PO43- which can be put together Fe2+ PO43- Fe3(PO4)2

  18. calcium sulfate The two ions are Ca2+ and SO42- which can be put together directly as CaSO4 (Note: there are no subscripts of 2 on the calcium and the sulfate ions in the final formula. Keep in mind that we are working with empirical formulas for ionic compounds. The one odd exception to this that you will encounter are the mercury (I) salts.)

  19. To use the Stock system, you need to know the charges on the common cations and anions.

  20. For the anions, pay attention to the endings. Within a given series, the name of the anion with more oxygen atoms usually ends with an “ate” ending and that with fewer oxygen atoms ends with an “ite” ending. Examples: ClO- hypochlorite ClO2- chlorite ClO3- chlorate ClO4- perchlorate

  21. Prefix system Use this for nonmetal nonmetal compounds.

  22. Examples: N2O dinitrogenmonoxide (also called dinitrogen oxide)

  23. Examples: N2O dinitrogenmonoxide (also called dinitrogen oxide) CO carbon monoxide

  24. Examples: N2O dinitrogenmonoxide (also called dinitrogen oxide) CO carbon monoxide Cl2O7dichlorineheptaoxide

  25. Examples: N2O dinitrogenmonoxide (also called dinitrogen oxide) CO carbon monoxide Cl2O7dichlorineheptaoxide P4O10tetraphosphorousdecaoxide

  26. Examples: N2O dinitrogenmonoxide (also called dinitrogen oxide) CO carbon monoxide Cl2O7dichlorineheptaoxide P4O10tetraphosphorousdecaoxide PF5 phosphorous pentafluoride

  27. Examples: N2O dinitrogenmonoxide (also called dinitrogen oxide) CO carbon monoxide Cl2O7dichlorineheptaoxide P4O10tetraphosphorousdecaoxide PF5 phosphorous pentafluoride N2dinitrogen

  28. Chemical Reactions and Chemical Equations – An introduction Chemical reaction: The transformation of one or more chemicals into different compounds.

  29. Examples: 2 Na + 2 H2O 2 NaOH + H2 (sodium) (water) (sodium dihydrogen hydroxide)

  30. Examples: 2 Na + 2 H2O 2 NaOH + H2 (sodium) (water) (sodium dihydrogen hydroxide) Cu + 4 HNO3 Cu(NO3)2 + 2 NO2 (copper) (nitric acid) (copper (II) nitrate) (nitrogen dioxide) + 2H2O

  31. Examples: 2 Na + 2 H2O 2 NaOH + H2 (sodium) (water) (sodium dihydrogen hydroxide) Cu + 4 HNO3 Cu(NO3)2 + 2 NO2 (copper) (nitric acid) (copper (II) nitrate) (nitrogen dioxide) + 2H2O Zn + 2HCl(aq) ZnCl2 + H2

  32. Reactants: The starting substances in a chemical reaction. E.g. Na and H2O in the first reaction.

  33. Reactants: The starting substances in a chemical reaction. E.g. Na and H2O in the first reaction. Products: Substances as a result of a chemical reaction. E.g. Cu(NO3)2, NO2, and H2O in the second reaction.

  34. Reactants: The starting substances in a chemical reaction. E.g. Na and H2O in the first reaction. Products: Substances as a result of a chemical reaction. E.g. Cu(NO3)2, NO2, and H2O in the second reaction. Balanced equation: The number of atoms of each element (free or part of a compound) is the same on both sides of the equation.

  35. The three equations given are all balanced. An example of a non-balanced equation is: H2 + O2 H2O The commonly employed definition of balanced equation is the smallest set of integers that leads to a balanced equation.

  36. So, H2 + ½ O2H2O and 4 H2 + 2 O2 4 H2O are not balanced equations. The balanced equation is: 2 H2 + O22 H2O

  37. The coefficients tell us how many molecules of each species are reacting.

  38. Measurements and Units Most physical quantities we encounter in chemistry have units. Old system: The cgs system – centimeter, gram, second were some of the key units in use – hence the abbreviation cgs

  39. The SI system:

  40. The SI system: Basic Quantity Name of Unit Symbol

  41. The SI system: Basic Quantity Name of Unit Symbol length meter m

  42. The SI system: Basic Quantity Name of Unit Symbol length meter m mass kilogram kg

  43. The SI system: Basic Quantity Name of Unit Symbol length meter m mass kilogram kg time second s

  44. The SI system: Basic Quantity Name of Unit Symbol length meter m mass kilogram kg time second s temperature kelvin K

  45. The SI system: Basic Quantity Name of Unit Symbol length meter m mass kilogram kg time second s temperature kelvin K amount of substance mole mol

  46. Derived Units: These are obtained from the basic units just given. Examples: The SI unit of volume is derived as follows: volume = length3(think about the volume of a cube). Since the unit of length is m, then the unit of volume is m3

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