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Stoichiometry!

Stoichiometry!. Formulas. Indicate the chemical composition of a compound. Subscripts – small numbers located after the element symbol in a formula that indicate how many of that element are in the formula. H 2 O - the 2 is a subscript

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Stoichiometry!

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  1. Stoichiometry!

  2. Formulas • Indicate the chemical composition of a compound. • Subscripts – small numbers located after the element symbol in a formula that indicate how many of that element are in the formula. • H2O - the 2 is a subscript • Coefficients – the large number in front of a formula that indicates the number of molecules or formula units. • 3 H2O The 3 is a coefficient!

  3. Formula Mass • The sum of all the atomic masses in a compound or formula. • Example: H2O 2 atoms of H (mass of 1) = 2 x 1 = 2 1 atom of O (mass of 16) = 1 x 16 = 16 18 amu How many atoms of Mg, O and H are in 3Mg(OH)2?

  4. Molar Mass • The molar mass (in g/mol) of any substance is always numerically equal to its formula weight in amu. • Molar mass is the mass of one mole of that substance.

  5. Mass Ratios No matter how many atoms of HO there are, as long as the number of H = the number of O, the ratio will always stay the same. 3 HO = 3 H and 3 O, the ratio is 3:48 or 1:16 The ratio between the masses that are present in a compound. Example: HO Mass of H = 1 Mass of O = 16

  6. Percent Composition What is the percent of an element in a compound? (by mass.) It is the mass of part of the compound divided by the mass of the whole compound.

  7. Example: 14 17 X 100 = 82.3% N = • What percent of ammonia is nitrogen (by mass)? • NH3 N = 1 x 14 = 14 H = 3 x 1 = 3 17 Nitrogen has a mass of 14 out of a total mass of 17 amu

  8. Types of Chemical Formulas • I. Molecular Formulas • They indicate the actual number and type of atoms in a molecular compound. These are both: Qualitative Quantitative and How many What element

  9. II. Empirical Formulas • The simplest whole number ratio of atoms in a formula • Empirical formulas are always used in ionic compounds. Molecular formula Empirical formula C2H6 CH3 C8H18 C4H9 C5H10 CH2 C6H12O6 CH2O

  10. Structural Formulas • Formulas that indicate the type of element present, how many atoms are present and how they are arranged. (Molecular) H H H C C H H H

  11. Determining Chemical Formula • If you know the compound’s percent composition, you can find the formula. • 1. Divide each percent value by that element’s atomic mass. • 2. Select the lowest value of all, and divide each number by that value.

  12. Example A substance is 40% sulfur and 60% oxygen. What is the chemical formula? S 40% ÷ 32 g = 1.25 O 60% ÷ 16 g = 3.75 1.25 ÷ 1.25 = 1 3.75 ÷ 1.25 = 3 SO3

  13. Try this! • Ascorbic acid (Vitamin C) contains 40.92% C, 4.58% H and 54.50 % O by mass. What is the empirical formula? • C = 40.92g/12.01g/mol = 3.407 mol C / 3.407 = 1 • H = 4.58 g/1.008g/mol = 4.54 mol H / 3.407 = 1.33 • O = 54.50g/16.00g/mol = 3.406 mol O / 3.407 = 1 • You need a whole number ratio! Multiply 1.33 x 3 = 4. Multiply the others by 3 as well. • The formula is C3H4O3

  14. Determining Molecular Formulas • If given the empirical formula and a molecular mass of a compound, you can determine the molecular formula. Molecular weight Empirical formula weight Whole number multiple =

  15. Example The empirical formula of a compound is CH2 and its molecular mass is 70g. What is the molecular formula? Step 1: C = 12 g H2 = 2 g ---------- 14 g Step 2: 70 g ÷ 14 g = 5 C5H10 CH2

  16. Try this! • Mesitylene, a hydrocarbon that occurs in small amounts in crude oil, has an empirical formula of C3H4. The experimentally determined molecular weight is 121 amu. What is the molecular formula? 121 amu 40 amu so, the formula is C9H12. = 3.02

  17. Let’s Talk Moles! • 1 mole is 6.02 x 1023 OR 602,000,000,000,000,000,000,000

  18. This unit of measurement, 6.02 x 1023, is called Avogadro’s number. It’s named after the Italian scientist Amedeo Avogadro who published it in 1811.

  19. So . . . How much does a mole weigh? It depends on what it is!

  20. Molar Mass • Molar mass is the mass of 1 mole of something (anything!) It is the same as the gram formula mass or the formula mass.

  21. Molar Mass • 1 mole of sodium (Na) = 6.02 x 1023atoms = 23 g • 1 mole of gold (Au) = 6.02 x 1023atoms = 197 g • 1 mole of lead (Pb) = 6.02 x 1023atoms = 207 g • 1 mole of carbon (C) = 6.02 x 1023atoms = 12 g • 2 moles of carbon(C) = 12.04 x 1023atoms = 24 g

  22. Examples • A sample of aluminum (Al) has a mass of 35 grams. • a. How many moles of Al is this? 27 g Al 35 g Al --------- = ---------- 1 mol Al x mol Al 27x = 35 ----- ---- x = 1.3 mol Al 27 27 Cross multiply

  23. Example 1, continued b. How many atoms would be in this sample? 6.02 x 1023 atoms x atoms ----------------------- = ------------------ 1 mol Al 1.3 mol Al (1.3)(6.02 x 1023) = 7.8 x 1023 atoms

  24. Example 2 • A sample of zinc (Zn) contains 0.65 moles. • a. How many grams of zinc is this? 65 g Zn x g Zn ---------- = ---------- 1 mol 0.65 mol (65)(0.65) = 42 g Zn

  25. Example 2, continued b. How many atoms of zinc are in this sample? 6.02 x 1023 atoms x atoms ----------------------- = ------------------ 1 mol Zn 0.65 mol Zn (6.02 x 1023)(0.65) = 3.9 x 1023 atoms Simple ratios!

  26. Helpful Hints Only use 2 variables at a time! (like moles and mass, or moles and atoms) Remember, moles stay underground!

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