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Chemical Formulas and Chemical Compounds

Chemical Formulas and Chemical Compounds. Chapter 7. 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.

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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

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