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Chapter 3 - Stoichiometry

This chapter covers the concepts of stoichiometry, including atomic masses, the mole, molar mass, percent composition, empirical and molecular formulas, chemical equations, balancing equations, and stoichiometric calculations.

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Chapter 3 - Stoichiometry

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  1. Chapter 3 - Stoichiometry

  2. 3.1 – Atomic Masses • Carbon-12, the relative standard • C-12 is assigned a mass of exactly 12 atomic mass units (amu) • Masses of all elements are determined in comparison to the 12C atom which is the most common isotope of carbon • Comparisons are made using a mass spectrometer • Atomic Mass (average atomic mass, atomic weight) • Atomic masses are the weighted average of the naturally occurring isotopes of an element • Atomic mass does not represent the mass of any actual atom • Atomic mass can be used to “weigh out” large numbers of atoms

  3. 3.2 – The Mole • Avogadro’s number • 6.022 x 1023 units = 1 mole • Measuring moles • An elements atomic mass expressed in grams contains 1 mole of that element • 12.01 grams of carbon = 1 mole of carbon • 12.01 grams of carbon = 6.022 x 1023 atoms of carbon • 1 mole of carbon = 6.022 x 1023 atoms of carbon

  4. 3.3 – Molar mass • Molar mass is the mass in grams of one mole of a compound • The sum of the masses of the component atoms in a compound is calculated using the PTE What is the molar mass of ethane (C2H6)? 2 moles of C = 2 x 12.01 g = 24.02 g 6 moles of H = 6 x 1.008 g = 6.048 g 30.07 g/mole ethane

  5. 3.4 – Percent Composition • A percentage is • “The part, divided by the whole, multiplied by 100” • Percent Composition • Calculate the percent composition of each element in the total mass of a compound (or sample) (# atoms of the element)(atomic mass of element) x 100 (molar mass of the compound) What is the % of sodium in sodium chloride? NaCl has a molar mass of 58.5 g/mol (23 /58.5) x 100 = 39%

  6. 3.5 – Determining the formula of a compound • Empirical formula – simplest whole number ratio of elements in a compound • Determine the percentage of each element in your compound • Treat % as grams and convert grams to moles using the mass from the PTE • Find the smallest whole number ratio of atoms (multiply by an integer to make them whole numbers). A compound contains 63.5% Silver, 8.2% Nitrogen and 28.2% Oxygen. What is the empirical formula for this compound? Ag = 63.5 g/ 108 g = .59 mol N = 8.2 g / 14 g = .59 mol O = 28.2 g / 16 g = 1.76 mol This is a 1:1:3 ratio Formula for this compound is AgNO3

  7. 3.5 – Determining the formula of a compound • Molecular formula – actual ratio of elements in a compound • Calculate the empirical formula mass • Divide the known molecular mass by the empirical formula mass deriving a whole number, n • Multiply the empirical formula by n to derive the molecular formula Ethane gas has an empirical formula of CH3 and a molecular mass of 30 g/mol. What is the molecular formula for ethane? Empirical formula mass is = 15 g/mol 30/15 = 2 2(CH3) = C2H6

  8. 3.6 – chemical equations • Chemical Reactions • Reactants are listed on the left had side of the arrow • Products are listed on the right hand side of the arrow • Atoms are neither created or destroyed • All atoms present in the reactants must be accounted for among the products, in the same number • No new atoms may appear in the products that were not present in the reactants REACTANTS yield PRODUCTS Zn (s) + 2 HCl (aq)  ZnCl2(aq) + H2 (g)

  9. 3.6 – chemical equations • The meaning of a chemical reaction • Solid - (s) • Liquid - (l) • Gas - (g) • Dissolved in water (aqueous solution) - (aq) • Relative numbers of reactants and products • Coefficients give atomic/molecular/mole ratios Zn (s) + 2 HCl (aq)  ZnCl2(aq) + H2 (g)

  10. 3.7 – Balancing Chemical Equations • Determine what reaction is occuring • It is sometimes helpful to write in word form • Write the unbalanced equation • Focus on writing correct atomic and compound formulas • Balance the equation using an atom inventory • It is often helpful to work systematically from left to right • Include phase information Hydrogen and oxygen gases combine to form liquid water

  11. 3.8 – Stoichiometric Calculations • Balance the chemical equation • Convert grams of reactant (or product) to moles, if required • Compare moles of the known value to moles of desired substance (mole ratio) • Convert from moles back to grams if required If 5 g of Zn reacts with excess hydrochloric acid, how many grams of zinc chloride are formed?

  12. 3.9 – Calculations involving a Limiting Reactant • A concept of limiting reactant • The limiting reactant controls the amount of product that can form • Solving limiting reactant problems • Convert grams of both reactants to moles • Use mole ratios to determine which reactant forms the least amount (moles) of desired product • Convert the least moles to grams (if required) If 5 g of Zn reacts with 5 g of hydrochloric acid, how many moles of hydrogen gas are formed?

  13. 3.9 – Calculations involving a Limiting Reactant • Calculating percent yield • Actual yield – what you got by actually doing the experiment • Theoretical yield – what stoichiometric calculations say the reaction should have produced. Actual yield (data table) x 100 = % yield Theoretical yield (Stoich) 5 g of Zn reacts with excess hydrochloric acid. 9.4 g of zinc chloride are formed. What is the percent yield for this reaction?

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