1 / 22

Stoichiometry

Stoichiometry. A measure of the quantities consumed and produced in chemical reactions. Atomic Mass. Is measured in either amu or in grams / mol 1 amu is exactly 1/12 the mass of a Carbon - 12 atom so it serves as the standard for all other atoms

ojal
Download Presentation

Stoichiometry

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. Stoichiometry A measure of the quantities consumed and produced in chemical reactions

  2. Atomic Mass • Is measured in either amu or in grams / mol • 1 amu is exactly 1/12 the mass of a Carbon - 12 atom so it serves as the standard for all other atoms • There are no atoms in nature that have the mass stated on the periodic table • These are averages of the most stable isotopes of that atom

  3. Calculating the Average Mass of an Element • There are two isotopes of copper, 63copper (62.93 amu and 69.09 % of copper atoms) and 65Cu (64.93 amu and 30.91 % of copper atoms). What is copper’s average atomic mass? • Practice Do 22 p. 123 (207.2 amu)

  4. The Mole • Avogadro’s number • 6.022 x 1023 of anything • Like a dozen • Relationship between the mole and amu to get grams / mole

  5. The Mole • Determining mass of a sample of atoms • Determining the number of moles of atoms from mass • Calculating number of atoms from mass

  6. Molar Mass • The mass of one mole of a substance • Just add together the elements mass • What is the mass of C10H6O3 • Practice – # 33 p. 123 • Practice - 1 g (1.0 x 10-6 ) of C7H14O2 contains how many molecules and how many atoms of carbon?

  7. Percent Composition • (Part / Whole) x 100 • What is the percent of Carbon in table sugar (C12H22O11)? • Practice – Calculate the mass percent of each element in C10H14O

  8. Determining Formulas • Empirical Formula • The simplest whole-number ratio of atoms in a compound • 3 steps • Determine the number of moles of each element • Divide all elements by the lowest mole number • Multiply all moles by a number that will make all whole if necessary

  9. Determining Formulas cont… • Empirical Formulas cont… • What is the empirical formula of a compound made from 43.64% phosphorus and 56.36 % oxygen by mass? • Practice - # 64 and 66 p. 125

  10. Determining Formulas cont… • Molecular Formula • The actual formula of a compound • Two things are needed to calculate it • The empirical formula of a compounds • The molar mass of the compound

  11. Determining Formulas cont… • Molecular formulas cont…. • SNH molar mass = 188.35 g/mol • NPCl2, 347.64 g / mol • CoC4O4, 341.94 g / mol • A compound is 41.39% C, 3.47% H, and the rest is oxygen, the molar mass is 116 g / mol • Practice – Do 68 p. 125

  12. Determining Formulas cont… • A compound consists of only C, H, and O. Combustion of 0.1156 grams of this compound produces 0.1638 g of CO2 and 0.1676 g of H2O, what is the empirical formula of this compound? • Practice – Do 71, and 73 p. 125 • Do Free Response # 3 2006 part (a)

  13. The meaning of a chemical equation • Gives you the number of each reactant required and the number of each product produced • Physical states are represented by (s), (l), (g), (aq) • The number of each type of atom must be the same on both sides of the equation.

  14. Balancing Chemical Equations • Whenever you see an equation the first thing you should ask yourself is whether or not it is balanced. • Most equations can be balanced by inspection • Balance all element other than H and O first, then do H, and finally O

  15. Balancing Chemical Equations cont… • Hints: • Make the odd one even • FeO + O2→ Fe2O3 • If the last element can be balanced by multiplying by a half, do so, then multiply all reactants and products by 2. • NH3 + O2→ NO + H2O • C2H6 + O2 → CO2 + H2O • Practice – do 79 p. 126

  16. Stoichiometric Calculations • Calculating masses of reactants or products from known amount • Balance equation • Convert known mass to moles of that same substance • Use mole ratio to convert to moles of other substance • Covert moles of other substance to grams if necessary

  17. Stoichiometric Calculations cont… • Sample: What masses of iron III oxide and aluminum must be used to produce 15.0 g of Fe from the following equation for the thermite reaction? • Fe2O3(s) + 2Al(s) → • 2Fe(l) + Al2O3(s) • Practice – do # 88 p. 127

  18. Calculations involving a limiting reactant • Unless the starting mole amounts of reactant are the exact right ratio (which is rare), then one of the reactants will run out before the other reactant(s) does. • This is the limiting reactant • It is the reactant we base all answers on • The reactant left over is known as the reactant in excess

  19. Calculations involving a limiting reactant • If you are given a starting amount of more than one reactant than you have a limiting reactant problem. • If the problem says anything like reacts with excess of another reactant it is not a limiting reactant problem

  20. Calculations involving a limiting reactant • Example of determining which reactant will run out first • Example: # 93 p. 127 • Practice – do 94 p. 127 • Example - # 95 p. 127 • Practice - # 96 p. 127

  21. Percent Yield • (Actual Yield / (Theoretical Yield) x 100 • Actual is the amount actually produced (given to you in the problem) • Theoretical is the mathematical calculation we do to determine how much we should get

  22. Percent Yield cont… • Sample # 99 p. 128 • Practice # 100 p. 128

More Related