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Ch 7.1 Forming Ions

Ch 7.1 Forming Ions. Review…. Cations are Groups 1A, 2A, and 3A They have positive charges. Anions are Groups 5A, 6A, and 7A They have negative charges They end in “ide” Majority of elements in Groups 4A and 8A do not usually form ions. These are called monatomic ions!.

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Ch 7.1 Forming Ions

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  1. Ch 7.1 Forming Ions

  2. Review… • Cations are Groups 1A, 2A, and 3A • They have positive charges. • Anions are Groups 5A, 6A, and 7A • They have negative charges • They end in “ide” • Majority of elements in Groups 4A and 8A do not usually form ions. • These are called monatomic ions!

  3. Ionic compounds are composed of two different elements from select groups on the periodic table of elements. • All of the compounds listed below are composed of elements located in groups 1A-7A on the periodic table of elements. • Collectively, these elements are known as the main group elements.

  4. Observing the Patterns • Consider the following group of formulas and their corresponding names. Examine the formulas and names to identify patterns associated with naming ionic compounds. Chemical Formula Written Form NaBrsodium bromide AlPaluminum phosphide CaF2calcium fluoride K2Opotassium oxide Mg3N2 magnesium nitride Write down any patterns you observed.

  5. Observing the Patterns • Any patterns you notice when going from name to Chemical Formula? Chemical Formula Written Form NaBrsodium bromide AlPaluminum phosphide CaF2calcium fluoride K2Opotassium oxide Mg3N2 magnesium nitride

  6. Binary Compound: is composed of two elements and can be either ionic or molecular.

  7. Naming Binary Ionic Compounds • To name a binary ionic compound, place the cation name first, followed by the anion name. • Remember the anion ends in “-ide” • Examples: • Cs2O • NaBr Cesium Oxide Sodium Bromide

  8. Name the Following: 1. SrCl2 Strontium Chloride 2. BaO Barium Oxide 3. Na2S Sodium Sulfide 4. Cs3N Cesium Nitride 5. LiI Lithium Iodide 6. MgF2 Magnesium Fluoride 7. K3P Potassium Phosphide 8. CaSe Calcium Selenide

  9. Ions of Transition Metals • Some transition metals in Groups 1B-8B form more than one cation with different charges. • Examples: • Iron forms Fe2+ and Fe3+ • Two methods of naming: • Stock System: Iron (II) ion and Iron (III) ion • Classical name: Ferrous ion and Ferric ion *We will use the stock system.

  10. Name the Following: • MnF2 Manganese(II) Fluoride 2. Ni3P2 Nickel(II) Phosphide • PbO2 Lead(IV) Oxide 4. Cs2S Cesium Sulfide (not a transition metal)

  11. Other Transition Metal Ions • Some transition metals have only one charge and do not use a Roman numeral. • Examples • Silver: Ag+ • Cadmium: Cd2+ • Zinc: Zn2+

  12. Polyatomic Ions • The names of most polyatomic anions end in “-ite” or “-ate” • The “-ite” ending indicates one less oxygen atom than the “-ate” ending • State the cation first and then the anion just as you did in naming binary ionic compounds. • Remember to use parentheses if more than one is needed.

  13. Name the Following: • KNO3 K + NO3 Potassium Nitrate 2. Mg(ClO2)2 Mg + ClO2 Magnesium Chlorite

  14. Writing a Formula Write the formula that will form between Ba and Cl Solution: 1. Write the positive ion of metal first, and then negative ion Ba2+Cl 2. Do the charges equal zero? NO!! 3. Use Criss-Cross method – write subscripts Ba2+ Cl1BaCl2

  15. Examples Write the correct formula for the compounds containing the following ions: 1. Na+, S2- 2. Al3+, Cl- 3. Mg2+, N3- 4. Al3+, S2-

  16. Solution 1. Na+, S2- Na2S 2. Al3+, Cl- AlCl3 3. Mg2+, N3- Mg3N2 4. Al3+, S2- Al2S3

  17. Use the Criss Cross Method

  18. Use the Criss Cross Method

  19. Use the Criss Cross Method

  20. Use the Criss Cross Method

  21. Use the Criss Cross Method

  22. But what if you don’t have the charges already? • Write the symbol for each element with their charge,criss-cross the charges and then rewrite with only the needed subscripts, no charges should be written on final formula. If you have a transition metal, the roman numerals tell you the charge of the metal ion. 1. I 2. II 3. III 4. IV 5. V

  23. Examples • Calcium Chloride CaCl2 • Aluminum Oxide Al2O3 • Magnesium Sulfide MgS • Copper(II) Nitride Cu3N2 5. Lithium Phosphide Li3P • Potassium Fluoride KF • Barium Nitrate Ba(NO3)2 • Iron(III) Sulfate Fe2(SO4)3

  24. Observing the Patterns • Consider the following group of formulas and their corresponding names. Examine the formulas and names to identify patterns associated with naming covalent molecules. Chemical Formula Written Form NF3Nitrogentrifluoride NONitrogenmonoxide NO2Nitrogendioxide N2ODinitrogenmonoxide N2O4Dinitrogentetroxide • Write down any patterns you observed

  25. Naming Binary Molecular Compounds • Binary molecular compound: must be composed of two nonmetals • Use prefixes to indicate the number and kind of atom in the compound • Use the following general format: 1st name: prefix + element name 2nd name: prefix + element name + “ide” If there is only 1 of the 1st element, no prefix.

  26. Anyone want a cold glass of dihydrogen monoxide?

  27. Examples • CO • Carbon Monoxide • CO2 • Carbon Dioxide • N2O • Dinitrogen Monoxide • Cl2O8 • Dichlorine Octoxide • N2H4 • Dinitrogen Tetrahydride • OF2 • Oxygen Difluoride • SBr2 • Sulfur Dibromide • BCl3 • Boron Trichloride • XeF4 • Xenon Tetrafluoride • ClF3 • Chlorine Trifluoride • P4O3 • Tetraphosphorous Trioxide • CS2 • Carbon Disulfide

  28. Writing Formulas for Binary Molecular Compounds • Use the prefixes in the name to tell you the subscript of each element in the formula. • Then write the correct symbols for the two elements with the appropriate subscripts. • The least electronegative element is written first • Dinitrogen Tetroxide - N2O4

  29. Examples • Nitrogen Monoxide • NO • Carbon Tetrachloride • CCl4 • Diphosphorous Pentoxide • P2O5 • Carbon Dioxide • CO2 • Dihydrogen Monoxide • H2O • Phosphorus Triiodide • PI3 • Sulfur Dichloride • SCl2 • Boron Trifluoride • BF3

  30. Naming Acids • An acid is a compound that contains one or more hydrogen atoms and produces hydrogen ions (H+) when dissolved in water. • All acids begin with hydrogen • General Format: HnX • “X” represents a monatomic or polyatomic anion. • “n” represents the number of hydrogen ions

  31. 3 Rules for Naming Common Acids • If the name of “X” ends in -ate: ____________-ic acid • If the name of “X” ends in -ite: ____________-ous acid • If the name of “X” ends in -ide: hydro-__________-ic acid

  32. Name these acids • H2SO4 • HCl • H2S • HNO3 • HClO2 Sulfuric Acid Hydrochloric Acid Hydrosulfuric Acid Nitric Acid Chlorous Acid

  33. Writing Formulas for Acids • If the acid ends in –ic, then “X” ends in –ate • If the acid ends in –ous, then “X” ends in –ite • If the acid has hydro-______-ic, then “X” ends in –ide. • The subscript on hydrogen is equal to the charge of “X”.

  34. Write the Formula for the Following Acids HBr • Hydrobromic Acid • Carbonic Acid • Phosphoric Acid • Sulfurous Acid H2CO3 H3PO4 H2SO3

  35. Ch 7.3 Using Chemical Formulas

  36. The Mass of a Mole of an Element • Molar Mass: is the atomic mass of an element expressed in grams/mole (g/mol). • Carbon = 12.01 g/mol • Hydrogen = 1.01 g/mol • When dealing with molar mass, round off to two decimals. 12.011 g/mol  12.01 g/mol

  37. The Mass of a Mole of a Compound • You calculate the mass of a molecule by adding up the molar masses of the atoms making up the molecules. • Example: H2O • H = 1.01 g x 2 atoms = 2.02 g/mol • O = 16.00 g x 1 atom = 16.00 g/mol • Molar Mass of H2O = 2.02 g/mol + 16.00 g/mol = 18.02 g/mol • This applies to both molecular and ionic compounds

  38. Find the molar mass of PCl3 • P = 30.97 g x 1 atom = 30.97 g/mol • Cl = 35.45 g x 3 atoms = 106.35 g/mol • PCl3 = 30.97 g + 106.35 g = 137.32 g/mol • What is the molar mass of Sodium Hydrogen Carbonate (NaHCO3) ? • Na = 22.99 g x 1 atom = 22.99 g/mol • H = 1.01 g x 1 atom = 1.01 g/mol • C = 12.01 g x 1 atom = 12.01 g/mol • O = 16.00 g x 3 atoms = 48.00 g/mol • NaHCO3 = 22.99 + 1.01 + 12.01 + 48.00 = 84.01 g/mol

  39. Converting Moles to Mass • You can use the molar mass of an element or compound to convert between the mass of a substance and the moles of a substance. • Mass (g) = # of moles x mass (g) 1 mole Example: If molar mass of NaCl is 58.44 g/mol, what is the mass of 3.00 mol NaCl? Mass of NaCl = 3.00 mol x 58.44g = 1 1 mol 175 g NaCl

  40. Example 2: Moles to Mass • What is the mass of 9.45 mol of Aluminum Oxide (Al2O3)? • Find molar mass of Al2O3 = 101.96 g/mol • Mass = 9.45 mol Al2O3x 101.96 g Al2O3 1 1 mol Al2O3 = 964 g Al2O3

  41. Converting Mass to Moles • You can invert the conversion factor to find moles when given the mass. • Moles = mass (g) x 1 mole mass (g) Example: If molar mass of Na2SO4 142.05 g/mol, how many moles is 10.0 g of Na2SO4? Moles Na2SO4 = 10.0 g Na2SO4x 1 mol Na2SO4 = 1 142.05 g Na2SO4 = 0.0704 mol Na2SO4

  42. Example 2: Mass to Moles • How many moles are in 75.0 g of Dinitrogen Trioxide? • Find molar mass of N2O3 = 76.02 g/mol • Moles = 75.0 g N2O3x 1 mole N2O3 = 1 76.02 g N2O3 N2O3 0.987 mol N2O3

  43. Percent Composition • Percent Composition: the relative amount of the elements in a compound. • Also known as the percent by mass • It can be calculated in two ways: • Using Mass Data • Using the Chemical Formula % mass of element= mass of element x100% mass of compound

  44. Example • When a 13.60 g sample of a compound containing Mg and O is decomposed, 5.40 g O is obtained. What is the % composition of this compound? Mass of compound: 13.60 g Mass of oxygen: 5.40 g O Mass of magnesium: 13.60 g - 5.40 g = 8.20 g Mg % Mg = 8.20 g Mg x 100% = 13.60 g % O = 5.40 g O x 100% = 13.60 g 60.3% 39.7%

  45. Find the percent composition of Cu2S. • Find mass of Cu and S • Cu = 63.55 x 2 = 127.10 g • S = 32.07 g • Find mass of Cu2S • 127.10 g + 32.07 g = 159.17 g % Composition • Cu = 127.10 g x 100% = 159.17 g • S = 32.07 g x 100% = 159.17 g 79.85% 20.15%

  46. Find the percentage by mass of water in the hydrate Na2CO310H2O • Find the mass of 10H2O • H2O = 10 mol x 18.02 g/mol = 180.2 g • Find the mass of Na2CO310H2O • Na2CO3 = 1 mol x 105.99 g/mol =105.99 g • 105.99 g + 180.2 g = 286.2 g • % By Mass of Water • H2O = 180.2 g H2O x 100% = 286.2 g Na2CO310H2O 62.96%

  47. Ch 7.4Determining Chemical Formulas

  48. Empirical Formulas • Empirical Formula: shows the smallest whole-number ratio of the atoms of the elements in a compound. • Example: • The Empirical Formula for Hydrogen Peroxide (H2O2) is HO with a 1:1 ratio. • The Empirical Formula for Carbon Dioxide (CO2) is CO2 with a 1:2 ratio.

  49. Determining the Empirical Formula of a Compound • A compound is found to contain 25.9% Nitrogen and 74.1% Oxygen. What is the Empirical Formula of the compound? • 25.9 g N x 1 mol N = 14.01 g N • 74.1 g O x 1 mol O = 16.00 g O • N1.85O4.63 = N1O2.5 = N2O5 1.85 mol N 4.63 mol O

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