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Chapter 7: Naming Compounds

Chapter 7: Naming Compounds. Naming Ionic Compounds. Review: Chemical formula indicates the name and number of atoms in a formula Ex: C 8 H 18 Al(SO 4 ) 3. Monoatomic ions-ions with just one atom Ex: Li + Naming Monoatomic Ions: Cations are simply the element’s name

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Chapter 7: Naming Compounds

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  1. Chapter 7: Naming Compounds

  2. Naming Ionic Compounds • Review: Chemical formula indicates the name and number of atoms in a formula • Ex: C8H18 • Al(SO4)3

  3. Monoatomic ions-ions with just one atom • Ex: Li+ • Naming Monoatomic Ions: • Cations are simply the element’s name • Anions-drop the ending of the element’s name and add –ide • Ex: F- Fluoride

  4. You Try It! • Name the following monoatomic anions: • Carbon • Nitrogen • Oxygen • chlorine

  5. Binary Ionic compounds • Compounds composed of two different compounds • The bond must be between a negative and a positive ion and the charges must balance to a zero

  6. Writing Binary Compound Formulas • Writing the formula: • Write each symbol with its charge • Write the cation first and then cross the charges • Al+3 O-2 • Al2O3

  7. You Try It! • Zinc and iodine • Zinc and sulfur • Potassium and iodine • Sodium and sulfur

  8. Naming Binary Ionic Compounds • Write the name of the cation first and then the anion (remember that the ending to the anion is changed) • Ex: Al2O3 • Aluminum Oxide

  9. You Try It! • AgCl • ZnO • CaBr2 • SrF2 • BaO • CaCl2

  10. Stock System of Nomenclature • Some cationic elements can have more than one correct charge so a Roman Numeral in parentheses is used after the symbol to show the correct charge • Ex: Iron (II) Fe+2 • Iron (III) Fe+3

  11. Naming an ionic Compound using the Stock System • CuCl2 • Copper (II) Chloride

  12. You Try It! • Write the formula and name the following compounds Using the Stock System: • Cu+2 and Br- • Fe+2 and O-2 • Pb+2 and Cl- • Hg+2 and S-2 • Sn+2 and F- • Fe+3 and O-2

  13. Polyatomic Ions • Pg. 210 Table 7-2 • A group of atoms with a common charge • Oxyanions-polyatomic ions that contain oxygen • Ex: OH-

  14. Some have the same two elements but are different • Most common ion is given the ending –ate • One less oxygen is given the ending –ite • One less oxygen than the –ite is given the prefix hypo- • One more oxygen than the –ate is given the prefix per-

  15. Examples • Hypochlorite ClO- • Chlorite ClO2- • Chlorate ClO3- • Perchlorate ClO4-

  16. Naming Binary Molecular Compounds • There are two methods to naming binary molecules • The older is the prefix system • The newer is the Stock System

  17. Prefix System • Table 7-3 • Rules for naming: • The less electronegative element is first. Prefix given only if it has more than one atom • The second word is begun with a prefix for the number of atoms, followed by the root word of the element, and ending with an –ide (indicating it has only 2 elements)

  18. Examples • P4O10 • Tetraphophorus decoxide • SO3 • Sulfur trioxide

  19. You Try It! • Write the name of the compound using prefixes: • ICl3 • PBr5 • N2O3

  20. Write the formulas for the compounds using the prefix system: • Carbon tetraiodide • Phosphorus trichloride • Nitrogen monoxide

  21. Acids • Binary acids-acids that consist of two elements, usually hydrogen as the cation and another element, dissolved in water. • Oxyacids-acids that contain hydrogen, oxygen, and a third element • Look at some examples on pg. 214 and Table 7-5

  22. How acids are named Compound Acid Name -ide hydro- -ic acid -ite -ous acid -ate -ic acid

  23. Examples • HClO3 Hydrogen Chlorate Acid: chloric acid • HClO2 Hydrogen chlorite Acid: Chlorous Acid

  24. HCl Hydrogen chloride Acid: hydrochloric acid

  25. Salt • An ionic compound composed of a cation and the anion from an acid

  26. Acid Salts • Ionic compounds that still contain an acidic hydrogen, such as NaHSO4 • Sodium hydrogen sulfate

  27. Hydrocarbons • Contain only carbon and hydrogen and are the simplest organic compounds • Alkanes are the simplest hydrocarbons and are only carbon-carbon single bonds.

  28. Alkanes (intro to organic chemistry) • Straight chain alkanes are where the carbons are all linked in a straight chain (no branching) • CnH2n+2

  29. Examples • CH4 Methane • C2H6 Ethane • C3H8 Propane • C4H10 Butane • C5H12 Pentane • C6H14 Hexane

  30. C7H16 Heptane • C8H18 Octane • C9H20 Nonane • C10H22 Decane

  31. Oxidation Numbers-Rules • A pure element has an oxidation number of zero • The more electronegative element is assigned the negative charge. The least electronegative element is assigned the number of its positive charge

  32. Fluorine is -1 always • Oxygen is almost always -2 • Hydrogen has +1 in compounds with elements more electronegative than it. It has -1 with metals • The sum of all oxidation numbers must equal zero • The oxidation number for polyatomic ions is equal to the sum of its charges

  33. Assigning Oxidation Numbers • UF6 Oxidation number for F is -1 (times 6 because there are 6 atoms= -6) So, to equal 0, the oxidation number of U must be +6

  34. H2SO4 • Most electronegative? Oxygen, so it has -2 (times 4 for a total of -8) • Hydrogen is least electronegative so it is +1(times 2 for a total of +2) • If we subtract +2 from -8 we get -6 so Sulfur must have +6

  35. You Try It! • Assign oxidation numbers: • HCl HF CI4 • CF4 H2O ClO2- • HNO3 H2CO3 IO3- • P4O10 PI3

  36. Table 7-6 gives oxidation numbers for some nonmetals that can have more than one. • These are also used in the Stock System learned earlier (roman numeral) • SO2 SO3 • Sulfur (IV) oxide and Sulfur (VI) oxide

  37. Formula Masses • Sum of the average atomic masses of all atoms in the formula • H2O • 2H X 1.01 = 2.02 • 1 O X 16.00= 16.00 • 16.00 + 2.02 = 18.02 amu for H2O

  38. You Try It! • Find the formula masses: • KClO3 • H2SO4 • Ca(NO3)2 • PO4-3 • MgCl2

  39. Molar Mass • Grams/mole for a molecule • Numerically equal to the formula mass • Ex: H2O is equal to 18.02 g/mol

  40. Conversion Factors • What is the mass in grams of 2.5 mol of Oxygen gas (O2)? • Amount in moles X molar mass = mass in grams • 2.5 mol X 32.00g = 80 g

  41. Sample Problem 7-9 • a. Grams to moles • B. Moles to molecules • C. Moles to moles to grams

  42. You Try It! • Practice Problems pg. 226

  43. Percent Composition • Percentage of one element in a compound • Mass of element in a compound X100 % Mass of a sample of compound

  44. Find the percent composition of copper(I) sulfide, Cu2S • Mass of copper = 127.1 g • Mass of Sulfur = 32.07 g • Molar Mass = 159.2 g • % Cu (127.1g/159.2g) X 100= 79.84% • % S (32.07 g/159.2 g) X 100= 20.14%

  45. You Try It! • Find the percent compositions of the following: • PbCl2 • Ba(NO3)2

  46. Empirical Formulas • Consists of symbols for the elements combined in a compound, with subscripts showing the smallest whole-number mole ratio of the different atoms in the compound • Ex: B2H6 • Empirical formula BH3

  47. How to Calculate Empirical Formula • BH3 • 78.1% B • 21.9% H • 78.1 g B X (1mol B/10.81 g B) =7.22 mol B • 21.9 g H X (1 mol H/1.01 g H) = 21.7 mol H

  48. These numbers give us numbers of moles but not smallest whole number ratios so divide each by the smallest number: • 7.22 mol B/7.22 = 1 B • 21.7 mol H/7.22 = 3.01 H • So, the formula is BH3

  49. You Try It! • Pg. 231 Practice Problems 1-3

  50. Molecular Formula • Actual formula of a molecule instead of the smallest whole-number formula • X(empirical formula)=molecular formula • Diborane’s formula mass is 27.67 amu and the empirical formula mass is 13.84 amu

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