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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 Chemistry chapter 7

  3. Chemical Formulas • Show either one molecule or one formula unit Chemistry chapter 7

  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 Chemistry chapter 7

  5. Ions • Ion – charged atom or group of atoms • Monatomic Ions – single atom • Polyatomic Ions – more than one atom Chemistry chapter 7

  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. Chemistry chapter 7

  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 Chemistry chapter 7

  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. Chemistry chapter 7

  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 Chemistry chapter 7

  10. Nomenclature • Naming system • Works for most compounds Chemistry chapter 7

  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 Chemistry chapter 7

  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 Chemistry chapter 7

  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- Chemistry chapter 7

  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 Chemistry chapter 7

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

  16. Polyatomic ions you must memorize • Ammonium • Acetate • Chlorate • Chlorite • Hydroxide • Hypochlorite • Nitrate • Nitrite • Perchlorate • Permanganate • Carbonate • Peroxide • Sulfate • Sulfite • Phosphate Chemistry chapter 7

  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 Chemistry chapter 7

  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 Chemistry chapter 7

  19. Examples • PF5 • Phosphorus pentafluoride • N2O5 • Dinitrogen pentoxide • OF2 • Oxygen difluoride Chemistry chapter 7

  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) Chemistry chapter 7

  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 Chemistry chapter 7

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

  23. Discuss • www.dhmo.org/facts.html Chemistry chapter 7

  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 Chemistry chapter 7

  25. Pure elements • Have oxidation numbers of zero • Single atoms – Na • Molecules of a pure substance • O2 • P4 • S8 Chemistry chapter 7

  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. Chemistry chapter 7

  27. Fluorine • Oxidation number of -1 • The most electronegative element Chemistry chapter 7

  28. Oxygen • Usually -2 • In peroxides, -1 • H2O2 • In compounds with halogens, +2 • OF2 Chemistry chapter 7

  29. Hydrogen • +1 with more electronegative elements • -1 with metals Chemistry chapter 7

  30. Sum of oxidation numbers • In a neutral compound must be zero • In a polyatomic ion must equal the charge on the ion Chemistry chapter 7

  31. Ion • Can be assigned an oxidation number equal to the charge on the ion Chemistry chapter 7

  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 Chemistry chapter 7

  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 Chemistry chapter 7

  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 Chemistry chapter 7

  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 Chemistry chapter 7

  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 Chemistry chapter 7

  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 Chemistry chapter 7

  38. Discuss • Name each of the following binary molecular compounds according to the stock system • CI4 • SO3 • As2S3 • NCl3 Chemistry chapter 7

  39. Formula mass • The sum of the average atomic masses of all the atoms in a formula • For ionic compounds or molecules • Can also be called molecular mass for molecules Chemistry chapter 7

  40. Example • Find the formula mass of Na2SO3 • 126.05 amu Chemistry chapter 7

  41. Example • Find the formula mass of HClO3 • 84.46 amu Chemistry chapter 7

  42. You try • Find the formula mass of MnO4- • 118.94 amu Chemistry chapter 7

  43. You try • Find the formula mass of C2H6O • 46.08 amu Chemistry chapter 7

  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 Chemistry chapter 7

  45. Example • Find the molar mass of K2SO4 • 174.27 g/mol Chemistry chapter 7

  46. You try • Find the molar mass of (NH4)2CrO4 • 152.10 g/mol Chemistry chapter 7

  47. Formula mass and molar mass • Numerically equal • Only the units are different Chemistry chapter 7

  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 Chemistry chapter 7

  49. Converting with molar mass • Relate mass in grams to number of moles • Relate mass in grams to number of particles Chemistry chapter 7

  50. Example • What is the mass in grams of 3.04 mol of ammonia vapor, NH3? • 51.8 g Chemistry chapter 7

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