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Chapter 3 Stoichiometry

AP Chemistry. Chapter 3 Stoichiometry. Finding atomic, molecular and formula masses. Atomic mass is the mass of an atom of an element These are the “red numbers” on the “buck-o-five” card. Molecular mass is the sum of the atomic masses of the atoms present in a covalent molecule.

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Chapter 3 Stoichiometry

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  1. AP Chemistry Chapter 3 Stoichiometry

  2. Finding atomic, molecular and formula masses. Atomic mass is the mass of an atom of an element These are the “red numbers” on the “buck-o-five” card. Molecular mass is the sum of the atomic masses of the atoms present in a covalent molecule. Formula mass is the sum of the atomic masses of the atoms present in a formula unit of an ionic compound. All three are expressed in atomic mass units (amu).

  3. Example #1 What is the atomic mass of hydrogen? Look on the periodic table and the red number for hydrogen is 1.0079 amu We’ll always round to 4 digits for mass, so the answer is 1.008 amu.

  4. Example #2 What is the molecular mass of water? The formula for water is H2O, this represents two hydrogens and one oxygen. hydrogen’s mass is 1.008 x 2 = 2.016 oxygen’s mass is 16.00 x 1 = 16.00 molecular mass of water 18.016

  5. Example #3 What is the formula mass of sodium hydroxide? The formula is NaOH. One sodium, one oxygen and one hydrogen. sodium’s mass is 22.99 x 1 = 22.99 hydrogen’s mass is 1.008 x 1 = 1.008 oxygen’s mass is 16.00 x 1 = 16.00 formula mass of NaOH 39.998

  6. Percent composition Percent composition is the % mass of each element in a compound. The formula is: mass you’re interested inx 100% total mass

  7. Example 4 What is the percent composition of water? First determine the molecular mass. hydrogen’s mass is 1.008 x 2 = 2.016 oxygen’s mass is 16.00 x 1 = 16.00 molecular mass of water 18.016 Then divide each by the total and change it to a percentage.

  8. Example 4 cont’d % hydrogen is 2.016x 100% = 11.2 % 18.016 % oxygen is 16.00x 100% = 88.8 % 18.016 Therefore, the percent composition of water is 11.2 % H and 88.8% O.

  9. The Mole Chemists rarely work with single atoms or molecules and can seldom measure anything as small as an amu in the lab. A larger unit was needed to make the masses more “user-friendly”. The mole was developed to be that unit. It was based upon Amedeo Avogadro’s theory about equal volumes of gases having the same number of molecules.

  10. The Mole cont’d Volumes of gases that had the same mass in grams as their atomic mass in amu were found to occupy 22.4 dm3. The number of molecules of gas in that volume was calculated to be 6.02 x 1023. If a mole is present, then the mass in amu is equal to the mass in grams. If the mass in grams is equal to the mass in amu, then a mole is present.

  11. The Mole cont’d An atom of sodium has the mass of 22.99 amu and a mole of sodium atoms has the mass of 22.99 g. Same number, different unit. A water molecule has the mass of 18.016 amu and 6.02 x 1023 water molecules have a mass of 18.016 g. Masses in grams are called the gram atomic, gram formula or gram molecular masses.

  12. Example 5 What is the mass in grams of 1 mole of carbon dioxide gas molecules? The formula is CO2. The mass is found by adding the masses. Carbon 12.01 x 1 = 12.01 Oxygen 16.00 x 2 = 32.00 44.01 amu One mole is therefore 44.01 g.

  13. mass to moles & moles to mass X MASS (g) MOLES ÷ By the red numbers • One mole is the same number as amu, just expressed in grams. • One molecule of water is 18.016 amu. • One mole of water molecules is 18.016g. • If you have 18.016g of water, you have a mole of water. Moles to different, multiply. Different to moles, divide.

  14. Example 6 How many grams of carbon dioxide are in 2.0 moles of carbon dioxide? The molecular mass of CO2 is44.01 amu. Since we’re working in moles, 1 mole of CO2 is 44.01 g. 2.0 moles CO2 x 44.01g CO2= 88.02g CO2 1mole CO2

  15. Example 7 How many moles of oxygen molecules are in 96.0 g of O2? The molecular mass of O2 is32.00 amu. Since we’re working in moles, 1 mole of O2 is 32.0 g. 96.0 g O2 x 1 mole O2 = 3.00 moles O2 32.00 g O2

  16. particles to moles & moles to particles X • One mole is 6.02 x 1023 anything (Avogadro’s Number). • One mole of water is 6.02 x 1023 water molecules • If you have 6.02 x 1023 molecules of water, you have a mole of water. particles MOLES ÷ By Avogadro’s number Moles to different, multiply. Different to moles, divide.

  17. Example 8 How many sodium atoms are in 2.0 moles of sodium? 2.0 moles Na x 6.02 x 1023 Na atoms = 1.2 x 1024 Na atoms 1 mole Na atoms How many moles of hydrogen molecules are in 3.01 x 1023 hydrogen molecules? 3.01 x 1023 H2 x 1 mole H2 molecules = 0.5 mole H2 6.02 x 1023 H2 molecules molecules

  18. volumes to moles & moles to volumes X • One mole of gas is 22.4 dm3 . • One mole of water is 22.4 dm3 water vapor. • If you have 22.4 dm3 water vapor, you have a mole of water vapor. volumes MOLES ÷ By 22.4 dm3 Moles to different, multiply. Different to moles, divide.

  19. Example 9 How many dm3 of nitrogen gas are in 2.0 moles of nitrogen? 2.0 moles N2 x 22.4 dm3 N2 = 44.8 dm3 N2 1 mole N2 How many moles of hydrogen molecules are in 67.2 dm3 hydrogen molecules? 67.2 dm3 H2 x 1 mole H2 = 3.0 moles H2 22.4 dm3 H2

  20. Empirical formula and molecular formula Empirical formula is the smallest whole number ratio of atoms of each element in a formula. Ionic compounds are always empirical. Molecular compounds can be empirical. If the real formula isn’t empirical then it is called the molecular formula. More info is needed to determine a molecular formula, it will be given.

  21. How to determine empirical formula If the analysis is given in percent, assume a 100 g sample and change to grams. If the analysis is given in grams continue. Divide the mass of each element by its gram atomic mass to get moles. Divide each by the smallest molar value to get the smallest ratio. If necessary, multiply to get whole numbers for the ratios.

  22. E.F. Example What is the empirical formula for a compound that is 88.2% oxygen and 11.2% hydrogen? Change % to grams (just change % to g). 88.2 g O x 1 mole O = 5.51 moles O 16.00 g O 11.2 g H x 1 mole H = 11.1 moles H 1.008 g H

  23. E.F. Example cont’d Divide each by the smallest: 5.51 moles O = 1.00 mole O 5.51 11.1 moles H = 2.01 moles H 5.51 The ratio is 2 H : 1 O, the formula is H20.

  24. AP Level E.F. Example A 3.489 g sample of a compound containing C, H, and O yields 7.832 g of carbon dioxide, and 1.922 grams of water upon combustion. (a) Calculate the number of moles of carbon produced. (b) Calculate the number of moles of hydrogen produced. (c) Calculate the number of moles of oxygen in the compound. (d) Calculate the empirical formula of the compound

  25. AP Level E.F. Example, cont'd (a) Calculate the number of moles of carbon produced: CxHyOz + O2 ----> CO2 + H2O 7.832 g CO x 1 mol CO2 x 1 mol C 44.01 g CO2 1 mol CO2 = 0.1780 mol C

  26. AP Level E.F. Example, cont'd (b) Calculate the number of moles of hydrogen produced. 1.922 g H O x 1 mol H2O x 2 mol H 18.02 g H2O 1 mol H2O = 0.2134 mol H

  27. AP Level Example, cont'd (c) Calculate the number of moles of oxygen in the compound. Mass oxygen = mass compound – mass hydrogen – mass carbon 0.1780 mol C x 12.01 g C = 2.138 g C 1 mol C 0.2134 mol H x 1.008 g H = 0.2151 g H 1 mol H

  28. AP Level Example, cont'd (c) Mass oxygen = 3.489 g –2.138 g-0.2151 g =1.136 g O 1.136 g O x 1 mol O = 0.07100 mol O 16.00 g O

  29. AP Level Example, cont'd (d) Calculate the empirical formula of the compound. 0.1780 mol C = 2.5 x 2 = 5 0.07100 mol O 0.2134 mol H = 3 x 2 = 6 0.07100 mol O 0.07100 mol O = 1 x 2 = 2 0.07100 mol O

  30. AP Level E.F. Example, cont'd (d) C5H6O2

  31. Writing a chemical equation. FOR A CHEMICAL EQUATION TO BE VALID: THE EQUATION MUST REPRESENT THE KNOWN FACTS. THE EQUATION MUST CONTAIN THE CORRECT FORMULAS OF THE REACTANTS AND PRODUCTS. THE LAW OF CONSERVATION OF MATTER MUST BE SATISFIED.

  32. When balancing: BALANCE ATOMS THAT APPEAR ONLY ONCE PER SIDE FIRST. TREAT POLYATOMIC IONS AS UNITS IF THEY APPEAR ON BOTH SIDES OF THE EQUATION. BALANCE ANY REMAINING NON-H OR O ATOMS NEXT. BALANCE H NEXT. BALANCE O LAST.

  33. DON’T CHANGE SUBSCRIPTS!!!! AVOID MISTAKES-USE THE CORRECT CHEMICAL FORMULA!

  34. Three Types of Equations The first type of Chemical Equation is the word equation Word equations are expressed as you would speak a sentence. For example: HYDROGEN PLUS OXYGEN REACT TO FORM WATER.

  35. Second type: The skeleton or unbalanced chemical equation Correct chemical formulas and symbols are substituted for the words. For the word equation: HYDROGEN PLUS OXYGEN REACT TO FORM WATER We write H2 + 02→ H2O

  36. The third type is the balanced chemical reaction • The law of conservation of matter must be obeyed. For every atom of oxygen on the reactant side, an atom of oxygen must be represented on the product side. To achieve this you add coefficients in front of the chemical formulas you need more of. • DO NOT CHANGE SUBSCRIPTS

  37. H2 + O2→ H20 is not balanced • Both sides have two H, but the product side only has one 0. We can't change the subscripts but we can have more than one H2O. Add another H2O then the 0 are balanced but we have two H on the reactant side and four on the product side. We balance the H by adding another H2. This gives the balanced chemical equation…

  38. 2H2 + O2→ 2H2OThis is a balanced chemical equation!

  39. COMMON SYMBOLS USED IN CHEMICAL EQUATIONS It is often useful to know what state of matter or phase a reactant or product is in. They are not to be balanced. (aq) aqueous dissolved or made up in a water solution (I) exists as a liquid during this rxn (s) exists as a solid during this rxn (cr) exists as a crystal solid during the rxn (g) exists as a gas during the rxn (↓), precipitates falls out of the solution as a solid product (↑) gas product leaves the solution as a gas product → yields separates the reactants from the products indicates the direction of the rxn

  40. ↔ reversible rxn can go forwards or backwards

  41. Five types of chemical reactions. Synthesis, composition or combination Analysis or decomposition Single replacement or displacement Double replacement or displacement (metathesis) Combustion

  42. Synthesis, composition or combination A + X → AX RXN OF ELEMENT WITH OXYGEN → OXIDE RXN OF 2 NONMETALS → COVALENT COMPOUND RXN OF METAL WITH NONMETALS, NOT OXYGEN → SALTS RXN OF OXIDES AND WATER → HYDROXIDES

  43. Analysis or decomposition AX → A + X DECOMP OF BINARY→ ELEMENTS DECOMP OF METAL CARBONATE → METAL OXIDE AND CO2 DECOMP OF METAL HYDROXIDE → METAL OXIDE AND H2O DECOMP OF METAL CHLORATES → METAL CHLORIDE AND 02 DECOMP OF ACIDS → NONMETAL OXIDES AND H2O

  44. Single replacement or displacement A + BX → AX + B or Y + BX → BY + X REPLACE. OF A METAL BY A MORE REACTIVE METAL REPLACE. OF H2 IN H20 BY A METAL → METAL HYDROXIDES AND H2 REPLACE. OF H2 IN AN ACID BY A METAL → SALT AND H2. REPLACE. OF HALOGENS

  45. Double replacement or displacement AY + BX → AX + BY FORMATION OF PRECIP.-ANIONS OF ONE REACTANT COMBINE WITH CATIONS OF THE OTHER REACTANT TO FORM AN INSOLUBLE OR SLIGHTLY INSOLUBLE COMPOUND. p 920 FORMATION OF GAS-WHEN THE PRODUCT PRODUCES A GAS WHICH IS INSOLUBLE IN THE NEW SOLUTION. THIS GAS WILL BUBBLE OUT OF SOLUTION.

  46. Combustion CxHx + O2→ CO2 + H2O Hydrocarbon or carbohydrate and oxygen yields carbon dioxide and water

  47. Reactions that form precipitates Complete ionic equation - break the chemical formulas into the ions they are formed from. Be sure to include symbols. Start with: AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s) Break it down into the aqueous ions: Ag+1(aq) + NO3-1(aq) +Na+1(aq) + Cl-1(aq) → Na+1(aq) + NO3-1(aq) + AgCl(s)

  48. Coefficients in a balanced chemical equation can be used to represent: 1. Moles 2. Particles (Molecules, Atoms or Formula Units) 3. Volumes of Gases BUT NEVER MASS!!!

  49. If the problem involves mass, CONVERT TO MOLES FIRST! If you are given atoms and the answer needs to be in volumes, CONVERT TO MOLES FIRST! If you are given mass and asked for molecules, CONVERT TO MOLES FIRST!

  50. For the rxn: 2H2 + 02 → 2H2O The coefficient 2 in front of the H2 can represent any of these. 2 Moles of Hydrogen Molecules 2 Molecules of Hydrogen 2 Volumes of Hydrogen Gas

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