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Chemical Formulas and Chemical Compounds

Chemical Formulas and Chemical Compounds

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Chemical Formulas and Chemical Compounds

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  1. Chemical Formulas and Chemical Compounds Chapter 7

  2. Chemical Formulas • Combinations of symbols are used to represent compounds of two or more elements. • Also indicate the ratio of the number of atoms of each type of element in the compound. • H2O – means that there are 2 hydrogen atoms for every oxygen atom. • No subscript on O – means there is 1

  3. Chemical Formulas • Show either one molecule or one formula unit

  4. Organic Compounds • Written differently than other formulas • The shorthand shows how the atoms are joined, not just the number present. • Example – • CH3COOH, not C2H4O2

  5. Ions • Ion – charged atom or group of atoms • Monatomic Ions – single atom • Polyatomic Ions – more than one atom

  6. Monatomic Ions • Can be anions or cations • Transition elements can form more than one kind of ion • See table 7-1 on page 205 • You must memorize this table.

  7. Naming monatomic ions • Cations • Element’s name • Roman numerals are used when there are multiple ions • Anions • Drop the element name ending • Add -ide

  8. Binary compounds • Contain two different elements • When we write chemical formula for a compound, the charges must add up to zero. • Write the positive ion first.

  9. Example • Write a formula for a compound of tin (II) and Iodine. • Tin (II) is 2+ • Iodine is 1- • We need two iodines to cancel out the charge on the tin (II). • SnI2

  10. Nomenclature • Naming system • Works for most compounds

  11. Naming binary compounds • Write the name of the positive cation first. • Add the name of the negative anion • AlN – Aluminum nitride • KCl – potassium chloride

  12. The stock system • Elements with more than one possible charge • Cu2S – copper (I) sulfide • CuS – copper (II) sulfide • Note – in an older naming system the above could be written as cuprous sulfide and cupric sulfide

  13. Oxyanions • Polyatomic ions that contain oxygen • When there are two or more oxyanions formed from the same two elements, the most common has the ending –ate • The ion with one less oxygen than –ate ends in –ite • The ion with one less oxygen than –ite adds the prefix hypo- • The ion with one more oxygen than –ate adds the prefix per-

  14. Compounds with polyatomic ions • See table 7-2 on page 210 • They are written like binary compounds. • Except the ending isn’t changed to end in –ide • CuSO4 – copper (II) sulfate • Sn(SO4)2 – tin (IV) sulfate

  15. Discuss • Practice problems 7-1, 7-2, and 7-3 on pages 207, 209, and 211 • Practice

  16. Polyatomic ions you must memorize • Ammonium • Acetate • Chlorate • Chlorite • Hydroxide • Hypochlorite • Nitrate • Nitrite • Perchlorate • Permanganate • Carbonate • Peroxide • Sulfate • Sulfite • Phosphate

  17. Naming binary molecular compounds • Two systems – one will be covered in section 7-2 • Older system • Prefixes used – see table 7-3 on page 212 • CO – carbon monoxide • CO2 – carbon dioxide • SO2 – sulfur dioxide • SO3 sulfur trioxide

  18. Rules • List the less-electronegative element first. • Only has a prefix if there is more than one. • The second element • Has a prefix • Root of the element name • -ide ending • If the word begins with a vowel, drop the o or a at the end of the prefix (monoxide, not monooxide) • Order: C, P, N, H, S, I, Br, Cl, O, F

  19. Examples • PF5 • Phosphorus pentafluoride • N2O5 • Dinitrogen pentoxide • OF2 • Oxygen difluoride

  20. Acids • Have a different naming rules. • Some common ones are listed in table 7-5 on page 214 • You should know • Hydrochloric acid (HCl) • Sulfuric acid (H2SO4) • Acetic acid (CH3COOH) (vinegar)

  21. Salts • An ionic compound composed of a cation and the anion from an acid • Sometimes the salt keeps one or more hydrogen atoms from the acid • The prefix bi- or the word hydrogen is added to the anion name • HCO3- • Hydrogen carbonate ion or bicarbonate ion

  22. Discuss • Sample problem 7-4 on page 213 • Practice

  23. Discuss • www.dhmo.org/facts.html

  24. Oxidation numbers • Also called oxidation states • Assigned to atoms in molecules • Indicate the general distribution of electrons among the bonded atoms • Sort of like ionic charge

  25. Pure elements • Have oxidation numbers of zero • Single atoms – Na • Molecules of a pure substance • O2 • P4 • S8

  26. Like charges on ions • Shared electrons are assumed to belong to the more-electronegative atom • The more electronegative element gets a number equal to the negative charge it would have as an anion. • The less electronegative element gets a number equal to the positive charge it would have as a cation.

  27. Fluorine • Oxidation number of -1 • The most electronegative element

  28. Oxygen • Usually -2 • In peroxides, -1 • H2O2 • In compounds with halogens, +2 • OF2

  29. Hydrogen • +1 with more electronegative elements • -1 with metals

  30. Sum of oxidation numbers • In a neutral compound must be zero • In a polyatomic ion must equal the charge on the ion

  31. Ion • Can be assigned an oxidation number equal to the charge on the ion

  32. Example • Assign oxidation numbers to each atom in the following compound: • KClO4 • O is -2, which gives -8, since there are 4. • The charge on perchlorate is 1-, so Cl must be +7 • K must be +1 to cancel out the 1- • +1, +7, -2

  33. Example • Assign oxidation numbers to each atom in the following compound: • SO32- • O is -2, which gives -6, since there are 3. • The charge on sulfite is 2-, so S must be +4 • +4, -2

  34. You try • Assign oxidation numbers to each atom in the following compound: • CO2 • O is -2, which gives -4, since there are 2. • The charge is 0, so C must be +4 • +4, -2

  35. You try • Assign oxidation numbers to each atom in the following compound: • NO3- • O is -2, which gives -6, since there are 3. • The charge is 1-, so N must be +5 • +5, -2

  36. More oxidation numbers • See Appendix Table A-15 • There is also a pattern on the periodic table • Group 1 is usually +1 • Group 2 is usually +2 • Group 13 is usually +3 • Group 14 is usually +2 or +4 • Group 15 is usually -3 • Group 16 is usually -2 • Group 17 is usually -1

  37. The stock system • Can be used instead of prefixes for molecular compounds • Use the oxidation number • SO2 • Sulfur dioxide • Sulfur (IV) oxide • SO3 • Sulfur trioxide • Sulfur (VI) oxide

  38. Discuss • Name each of the following binary molecular compounds according to the stock system • CI4 • SO3 • As2S3 • NCl3

  39. Formula mass • The sum of the average atomic masses of all the atoms in a formula • For ions or molecules • Can also be called molecular mass for molecules

  40. Example • Find the formula mass of Na2SO3 • 126.05 amu

  41. Example • Find the formula mass of HClO3 • 84.46 amu

  42. You try • Find the formula mass of MnO4- • 118.94 amu

  43. You try • Find the formula mass of C2H6O • 46.08 amu

  44. Molar Mass • Chapter 3 • The mass in grams of one mole (6.022 x 1023 particles) of a substance • Example: H2O • The mass of two moles of hydrogen atoms and one mole of oxygen atoms

  45. Example • Find the molar mass of K2SO4 • 174.27 g/mol

  46. You try • Find the molar mass of (NH4)2CrO4 • 152.10 g/mol

  47. Formula mass and molar mass • Numerically equal • Only the units are different

  48. Discuss • How many moles of atoms of each element are there in one mole of ammonium carbonate, (NH4)2CO3 • 2 mol N, 8 mol H, 1 mol C, 3 mol O • Determine both the formula mass and the molar mass of ammonium carbonate • 96.11 amu, 96.11 g/mol

  49. Converting with molar mass • Relate mass in grams to number of moles • Relate mass in grams to number of particles

  50. Example • What is the mass in grams of 3.04 mol of ammonia vapor, NH3? • 51.8 g