1 / 45

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!.

elizabethprice
Download Presentation

Ch 7.1 Forming Ions

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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. 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.

  4. Common Metal Ions with More than One Ionic Charge

  5. Common Metal Ions with More than One Ionic Charge

  6. Other Transition Metal Ions • Some transition metals have only one charge and do not use a Roman numeral. • Examples (Write on your periodic table!) • Silver: Ag+ • Cadmium: Cd2+ • Zinc: Zn2+

  7. 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 • Remember to use parentheses if more than one is needed.

  8. 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

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

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

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

  12. 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 • CuO Cesium Oxide Sodium Bromide Copper(II) Oxide

  13. Naming Compounds withPolyatomic Ions • State the cation first and then the anion just as you did in naming binary ionic compounds. • KNO3 • Mg(ClO2)2 Potassium Nitrate Magnesium Chlorite

  14. 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.

  15. Anyone want a cold glass of dihydrogen monoxide?

  16. Examples • Name the following • CO • CO2 • N2O • Cl2O8 Carbon Monoxide Carbon Dioxide Dinitrogen Monoxide Dichlorine Octoxide

  17. 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

  18. Examples • Write formulas for the following: • Nitrogen Monoxide • Carbon Tetrachloride • Diphosphorous Pentoxide NO CCl4 P2O5

  19. 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

  20. 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

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

  22. 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”.

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

  24. Homework • 7.1 pg 251 #14, 15(a-d), 41 • Look at page 226 and 230 for help with acids.

  25. Ch 7.3 Using Chemical Formulas

  26. 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

  27. 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

  28. 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

  29. 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

  30. 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

  31. 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

  32. 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

  33. Homework • 7.3 pg 253 #30-31

  34. 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

  35. 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%

  36. 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%

  37. 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%

  38. Homework • 7.3 pg 253 #32-33, 47

  39. Ch 7.4Determining Chemical Formulas

  40. 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.

  41. 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

  42. Molecular Formulas • Molecular Formula: tells the actual number of each kind of atom present in a molecule of a compound • Example: • The Molecular Formula for Hydrogen Peroxide is H2O2. • The Molecular Formula for Carbon Dioxide is CO2 • It is possible to find the Molecular Formula using the Empirical Formula if you know the molar mass of the compound.

  43. Finding the Molecular Formula • Calculate the molecular formula of a compound whose molar mass is 60.0 g/mol and empirical formula is CH4N • Step 1: Find the empirical formula molar mass • 12.01 + (4 x 1.01) + 14.01 = 30.06 g/mol • Step 2: Divide molar mass by EF molar mass • 60.0 g/mol = 1.99  2 30.06 g/mol • Step 3: Multiply empirical formula by 2 • CH4N x 2 = C2H8N2

  44. Homework • 7.4 pg 253 #36-38

More Related