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Unit 7

Unit 7. Compound Stoichiometry. Terminology. Atomic mass: Mass used for elements and is found on the periodic table Molecular mass: Mass used for covalent compounds Formula mass: Mass used for ionic compounds Molar mass: Mass used as a general term for all of the above masses.

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Unit 7

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  1. Unit 7 Compound Stoichiometry

  2. Terminology • Atomic mass: • Mass used for elements and is found on the periodic table • Molecular mass: • Mass used for covalent compounds • Formula mass: • Mass used for ionic compounds • Molar mass: • Mass used as a general term for all of the above masses.

  3. Steps to finding molar mass • 1) Write a correct formula for the compound • 2) Round the average atomic mass (found on the periodic table) of each element to two decimal places • 3)Multiply the rounded average atomic mass of each element by the number of atoms (each atom’s subscript) • 4) Add the total rounded atomic masses of all elements. Label with amu or g/mol

  4. Practice #1 • Find the formula mass of calcium phosphate

  5. Practice #2 Find the formula mass for ammonium sulfate

  6. Find the molecular mass of dichlorineheptaoxide Practice #3

  7. Terms • What is a mole? • A counting number: just as one dozen equals 12, one mole = 6.02x1023 • What is Avogadro's number? • 6.02x1023 particles

  8. What is the difference between a mole of elephants, a mole of feathers, and a mole of water?

  9. How many moles are in 28.7 grams of lithium nitrate? When working with moles, molar mass can be used as a conversion factor. We will use this while performing calculations with dimensional analysis

  10. How many moles are there in 100.0 grams of Al2(SO4)3?

  11. If I tell you to measure out 1.34 moles of water, how many grams will you pour?

  12. How many grams are in 2.8 moles of oxygen (diatomic)?

  13. If 256 grams of a pure monatomic element is known to be 8.26 moles of that element, what element is it?

  14. One mole of a monatomic element (Mg) = • One mole of a diatomic element (O2) or covalent compound (H2O) = • One mole of any ionic bond (NaCl) = • Restate Avogadro’s number and its meaning: 6.02x1023 particles: the number of atoms in one mole of substance

  15. 1. How many formula units are there in 7.55 moles of sodium sulfate? 2. How many moles are in 8.24x1023 molecules of water

  16. 3. How many formula units are there in 0.75 moles of sodium bromide? 4. How many moles of carbon dioxide are present in 4.55x1024 molecules of carbon dioxide?

  17. Mole Map!!!

  18. Two step Dimensional Analysis The first thing you must do when doing any conversion is convert to moles • How many atoms are present in a pure iron nail that weighs 13.2 grams

  19. How many grams of magnesium chloride does 6.25x1025 formula units contain?

  20. Practice #1 How many formula units are in 20.0 grams of sodium chlorate?

  21. Practice #2 If there are known to be 8.24x1026mlcl of water present, would it weigh more or less than a pound?

  22. Practice #3 If you know that a sample of mercury (II) chloride contains 7.32x1024 formula units, how much does the same weigh in pounds?

  23. Representative Particles • 1) elements (except diatomics) = • 2) diatomic elements = • 3) Molecular compound = • 4) Ionic compound =

  24. Example 1 How many carbon atoms are there in 16.0 grams of glucose (C6H12O6)

  25. Example 2 If a sample of aluminum sulfite is known to contain 4.55x1023 atoms of aluminum, what is the mass of the entire sample of aluminum sulfite in grams?

  26. Example 3 How many nitrogen atoms are there in 45.8 grams of ammonium nitride?

  27. Percent Composition • With any percentage problem you are comparing part to the whole. • In chemistry we do this based on mass, not by number of atoms. To find the percentage of certain elements in a compound you will compare the mass of the element to the mass of the entire compound

  28. Percent composition steps • 1) Write the correct formula • 2) Find the formula or molecular mass of the entire compound • 3) Divide total mass of the element whose percentage you are looking for by the total mass of the entire element. Multiply the result by 100 to convert to a percentage • 4) Since there are no significant digits given in the problem, round your answer to 2 decimal places and add the % sign as the label.

  29. Example What is the complete percentage composition (by mass) of potassium carbonate?

  30. Practice • What is the percent of calcium in calcium phosphide?

  31. What is the percent of oxygen in barium chlorate?

  32. How many grams of pure magnesium could be recovered from the decomposition of magnesium fluoride? • Method 1: • Method 2:

  33. How many grams of calcium are present in 156.8 grams of chalk (calcium carbonate)? • Method 1: • Method 2:

  34. Practice • How many grams of pure iron can be recovered from 50.0 grams of an ore known to be 32% iron (II) nitrate? • Method 1: • Method 2:

  35. Practice • Limestone is known to be 95% calcium carbonate. How many grams of pure calcium could you expect to obtain from a piece of limestone which weighs 100 Kg? • Method 1: • Method 2

  36. Terms • What is a molecular formula? • True formula of a molecule • What is an empirical formula? • Lowest whole number ratio of atoms: can represent several different atoms • Ionic compounds will only have an empirical formula

  37. Application • What do the subscripts represent in the molecular formula? • The actual number of atoms in the molecule • How do the subscripts differ for the empirical formula? • Smallest whole number ratio

  38. If the molecular formula is C6H12O6, and the empirical formula is CH2O, how many “empirical formulas” (empirical formula units (EFU)) are needed to make the molecular formula. • What is the molar mass for the empirical formula? How does it relate mathematically to the molar mass of the molecule?

  39. Empirical Formula Steps • 1) Percent to mass: assume 100 grams. Drop the percent label and add the grams label (42% becomes 42 grams) • 2) Mass to mole: convert the masses found in the first step to moles of each element. Leave at least 4 sig figs • 3) Divide by smallest: Whichever element has the smallest amount of moles (found in step 2) divide all moles by this amount. • 4) Multiply ‘til whole: you must end in a whole number.

  40. Empirical Steps • Memorize the saying : • Percent to mass, mass to mole, divide by smallest, multiply ‘til whole.

  41. Example • A compound is found to contain 34.39% Zinc, 14.82% nitrogen, and 50.79% Oxygen. What is the empirical formula of the compound?

  42. Practice • A 200 gram sample of compound which contains only carbon, hydrogen, and oxygen is found to contain 94.74 grams of carbon, 21.05 grams of hydrogen, and 84.21 grams of oxygen. What is the empirical formula of the compound?

  43. Molecular Formula Steps • 1) Find the molar mass of the empirical formula (MMef) • 2) Divide the given molecular mass of the entire molecule (MMmlcl) by the molecular mass of the empirical formula (MMef). This number gives you the number of EFUs(empirical formula units). • 3) Use the whole number result to multiply all of the subscripts of the empirical formula by and write the molecular formula

  44. Example • The empirical formula and molar mass for a substance were determined to be C2H4O2 and 132.16 g/mol respectively. Determine the molecular formula for the substance.

  45. Practice • If the empirical formula for a carbohydrate is CH2O, and its molecular mass is known to be 240 g/mol, what is the molecular formula for the carbohydrate? • If the empirical formula for a hydrocarbon is known to contain 92.3% carbon, and its molecular mass is known to be 78 g/mol, what is the molecular formula for the compound?

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