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Check for Understanding-Whiteboard

Check for Understanding-Whiteboard

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Check for Understanding-Whiteboard

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  1. Check for Understanding-Whiteboard • Name the following: Cr2O3FeO NiCl2S3Cl6 HI HNO3 • Write the chemical formula for the following: Iron (III) Oxide Calcium Nitride Diphosphorushexaiodide Carbonic acid Hydrochloric acid

  2. Stoichiometry Ch. 3

  3. Lab A Partners • Group 1: • Group 2: • Group 3: • Group 4: • Group 5:

  4. Chapter 3 - Stoichiometry • Atomic Mass • The Mole • Molar Mass • Percent Composition • Empirical and Molecular Formulas • Balancing Chemical Equations • Stoichiometry—Limiting Reactants • Percent Yield

  5. Learning Objectives – Part 1 • I can calculate average atomic mass of an atom from mass spectrometry and isotopic data. • I can calculate the molar mass of a compound and the percent composition of an element in a compound. • I can convert between mass, moles, and numbers of representative particles (atoms, molecules, formula units) of a substance.

  6. Warm-up! • Balance the following equations. (a.) (NH4)2Cr2O7(s)  Cr2O3(s) + N2(g) + H2O(g) (b.) NH3(g) + O2(g)  NO(g) + H2O(g)

  7. The Mole • Definition: • The number equal to the number of carbon atoms in exactly 12 grams of pure C-12. This number has been determined (via Mass Spectrometry) to be 6.022 *10 23 (Avogadro’s Number). • One mole of something consists of 6.022 *10 23 units or particles of that substance. (Like a dozen or ream) • Video • The mass of 1 mole of an element is equal to its atomic mass in grams.

  8. Example – The Mole • Example Problem 1: Americium is an element that does not occur naturally. It can be made in very small amounts in a device known as a particle accelerator. Compute the mass in grams of a sample of americium containing six atoms. • Example Problem 2: A silicon chip used in an integrated circuit of a microcomputer has a mass of 5.68 mg. How many silicon atoms are present in the chip?

  9. Molar Mass • The mass in grams of one mole of the compound. • Obtained by summing the masses of the component atoms. • Example Problem 3: Juglone, a dye known for centuries, is produced from the husks of black walnuts. It is also a natural herbicide that kills off competitive plants around the black walnut tree but does not affect grass and other noncompetitive plants. The formula for juglone is C10H6O3. • (a.) Calculate the molar mass of juglone. • (b.) A sample of 1.56 x 10-2 g of pure juglone was extracted from black walnut husks. How many moles of juglone does this sample represent?

  10. Molar Mass Practice Time - 10 minutes

  11. Percent Composition • Mass percent of the elements can be determined from the formula of the compound by comparing the mass of each element present in 1 mole of the compound to the total mass of 1 mole of compound. • Hint: Always check that percentages add up to 100. • Example Problem 7: Calculate the mass percent of ethanol, C2H5OH.

  12. Percent Composition – Practice 10 minutes

  13. Empirical/Molecular Formulas • Empirical – Molecular Formula Determination • 1. Mass Percent to Grams (100 g sample) • 2. Convert each element from grams to moles • 3. Divide all mole values by the smallest. • 4. Subscripts obtained if whole number. = Empirical Formula • 5. Divide molar mass by empirical formula mass to obtain integer to multiple all subscripts in formula. “Method 1”

  14. Calculating Empirical Formulas Example Problem 10: Determine the empirical and molecular formulas for a compound that is given the following analysis (in mass percents): 71.65% Cl; 24.27% C; 4.07% H. The molar mass is known to be 98.96 g/mol.

  15. Check for Understanding- Whiteboard A sample compound weighing 83.5 g contains 33.4 g of sulfur. The rest is oxygen. What is the empirical formula and chemical name?

  16. Empirical/Molecular Formulas • Empirical – Molecular Formula Determination • 1. Mass Percent to Grams (100 g sample) • 2. Convert each element from grams to moles • 3. Divide all mole values by the smallest. • 4. Subscripts obtained if whole number. = Empirical Formula • 5. Divide molar mass by empirical formula mass to obtain integer to multiple all subscripts in formula. “Method 1”

  17. Example • A sample of 100g of acetic acid contains 39.9 g C, 6.7 g H, and 53.4 g O. The molar mass of the compound was determined by experiment to be 60.0 g/mol. Determine the molecular and empirical formula.

  18. Learning Objectives –Part 2 • I can write a balanced chemical equation and predict the amount of product formed from a given mass of reactant or the amount of reactant required to produce a desired amount of product. • I can identify limiting reactants, and calculate the amount of product formed when given the amounts of all the reactants present. • I can calculate the percent yield of a reaction.

  19. Chemical Equations • Chemical Reactions • Reorganization of the atoms in one or more substances. • Bonds are broken and new ones form. • Atoms are neither created nor destroyed. • All atoms present in the reactants must be accounted for among the products. = Balancing a chemical equation! • Reactants (initial substances) left of arrow • Products (substances produced) right of arrow • The relative number of reactants and products in a reaction are indicated by the coefficients in the balanced equation.

  20. Balancing Chemical Equations • ALWAYS check to make sure the equation you are working with is balanced!!! • Use element tables to help. • Writing and Balancing the Equation for a Chemical Reaction • 1. Determine what reaction is occurring. What are the reactants, the products, and the physical states involved? • 2. Write the unbalanced equation that summarizes the reaction described in step 1. • 3. Balance the equation by inspection, starting with the most complicated molecule(s). Determine what coefficients are necessary so that the same number of each type of atom appears on both reactant and product sides. Do NOT change the identities (formulas/subscripts) of any of the reactants or products.

  21. Example • Nitrogen and oxygen combine to form dinitrogenpentaoxide. • Calcium nitrate reacts with lithium sulfide forming calcium sulfide and lithium nitrate.

  22. Partner Practice • Solutions of lead (II) chloride and sodium chromate react to produce a precipitate of lead (II) chromate and a solution of sodium chloride. • Solid calcium reacts with solid sulfur to produce solid calcium sulfide. • Hydrogen gas reacts with fluorine gas to produce hydrogen fluoride gas.

  23. Warm-up Question Write the following balanced chemical reactions. • Liquid carbon disulfide reacts with oxygen gas to produce carbon dioxide gas and sulfur dioxide gas. • Aqueous solutions of sodium chloride and silver nitrate react to produce aqueous sodium nitrate and a precipitate of silver chloride. • Bubbling chlorine gas through a solution of potassium iodide gives elemental iodine and a solution of potassium chloride.

  24. Combustion Analysis - POGIL • Analyze substances for carbon and hydrogen. • A sample is burned in the presence of oxygen, which converts all its carbon to carbon dioxide and all its hydrogen to water. • The compounds are absorbed using appropriate materials and their amounts determined. • The results provide the mass of each type of element in the compound, which can be used to determine mass percent of each element and/or empirical formula.

  25. Peer activity- Combustion Analysis ALL GROUP Members participate in building understanding! Below are side jobs to learning and collaborating. • Getter 1 (1) • Group member that Collects and distributes materials for beginning of activity/day. • Group member that Collects and returns materials at the conclusion of activity/day. • Reporter (2) • Group member that shares thoughts with class; only member that can ask teacher question • Starter (3) • Group member that begins activities and keeps an eye on the time

  26. Combustion Analysis Activity Groups Getter StarterReporter Group 1: Jenna- Jacob- Kaitlyn Group 2: Kirsten-Tyler- Atul Group 3: Caleb- Tess- Emily Group 4: Aravin-Sura- Alyssa Group 5: Natalie- Rabab-Tim

  27. Check for Understanding -whiteboard with a partner!

  28. Hydrates • Example Problem 17:(Hydrates) Washing soda is a hydrate. Its formula can be written as Na2CO3∙xH2O. When a 2.558-g sample of washing soda is heated at 25°C, all the water of hydration is lost, leaving 0.948 g of Na2CO3. What is the value of x? [10; Na2CO3∙10H2O]

  29. Stoichiometry  Flowchart on page 111 • 1. Balance the equation for the reaction • 2. Convert the known mass of the reactant or product to moles of that substance. • 3. Use the balanced equation to set up the appropriate mole ratios. • 4. Use the appropriate mole ratios to calculate the number of moles of the desired reactant or product. • 5. Convert from moles back to grams if required by the problem. • Example

  30. Example Problem 18:(A template for solving stoichiometry problems!!) • (a) What mass of oxygen will react with 96.1 grams of propane? • (b) How many liters of carbon dioxide gas at STP (1 atm, 273K) would be formed? • (c) How many water molecules are produced?

  31. Stoichiometry Model • Example Problem 19: Solid lithium hydroxide is used in space vehicles to remove exhaled carbon dioxide from the living environment by forming solid lithium carbonate and liquid water. What mass of gaseous carbon dioxide can be absorbed by 1.00 kg of lithium hydroxide?

  32. Practice • Calculate the mass of oxygen produced if 2.50 g of potassium chlorate are completely decomposed by heating. • (0.976 g O2)

  33. Practice • If 20.0 g of magnesium react with excess hydrochloric acid, how many grams of magnesium chloride are produced? • (78.4 g MgCl2)

  34. Practice • How many grams of chlorine gas must be reacted with excess sodium iodide if 10.0 g of sodium chloride are needed? • (6.07 g Cl2)

  35. Practice • What mass of copper is required to replace silver from 4.00 g of silver nitrate? (Hint: Copper (II) product is formed!) • (0.747 g Cu)

  36. Limiting Reactants • Think of this as just a multi-part stoichiometry problem. • Example 21: • The Haber process (the AP Exam likes to ask questions about this…) is making ammonia for fertilizer production from the nitrogen in the air reacted with hydrogen gas. The hydrogen gas is obtained from the reaction of methane with water vapor. This process has saved millions from starvation. Suppose 25.0 kg of nitrogen reacts with 5.00 kg of hydrogen to form ammonia. What mass of ammonia can be produced? Which reactant is the limiting reactant? What is the mass of the reactant that is in excess?

  37. Practice • Zinc metal reacts with hydrochloric acid by the following reaction: Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) • If 0.30 moles of zinc is added to hydrochloric acid containing 0.52 molHCl, how many moles of H2 are produced?

  38. Practice • In a process for producing acetic acid, oxygen gas is bubbled into acetaldehyde, CH3CHO, containing manganese (II) acetate under pressure at 60C. • 2CH3CHO(l) + O2(g) 2HC2H3O2(l) • In a laboratory test of this reaction, 20.0 g of CH3CHO and 10.0 g of O2 were put into a reaction vessel. • How many grams of acetic acid can be produced by this reaction from these amounts of reactants? • How many grams of excess reactant remain after the reaction is complete?

  39. Percent Yield • Theoretical Yield • The amount of a product formed when the limiting reactant is completely consumed. • Percent Yield • The actual yield of product is often given as a percentage of the theoretical yield. • (Actual yield ÷ Theoretical Yield) x 100% = Percent Yield

  40. Percent Yield Model • Example Problem 24: Methanol (CH3OH), also called methyl alcohol, is the simplest alcohol. It is used as a fuel in race cars and is a potential replacement for gasoline. Methanol can be manufactured by combination of gaseous carbon monoxide and hydrogen. Suppose 68.5 kg CO(g) is reacted with 8.60 kg H2(g). Calculate the theoretical yield of methanol. If 3.57 x 104 g CH3OH is actually produced, what is the percent yield of methanol?

  41. HCl(aq) + NaOH(aq) NaCl(s) + H2O(l) Start with 5.0 grams of NaOH and after the reaction I have 6.3 grams of NaCl. What is the percent yield of the reaction?

  42. With A Partner • A student runs a reaction to prepare 40.0g of aspirin and yet recovers only 15.5g. What is the percent yield (please show/be able to explain your work)?

  43. CH4+ 2O2 CO2 + 2H2O If you react 5.00 g of methane, CH4, with 12.00g of oxygen, O2, what is the theoretical yield of H2O? After completing the reaction there was 4.60g of H2O, what is the percent yield?

  44. Quiz Unit 1 • Complete Individually! • Pre-lab for Lab B DUE TOMORROW! • Read through lab handout (in black binder) • Complete the 2 pre-lab questions • Write an abstract of Part 2 (omit Part 1 – sample data provided)

  45. Lab Safety • Lab Safety Video • Lab Safety Contract • Lab Expectations

  46. Labs over Ch. 3 • Lab A: Empirical Formula • Complete Lab Report • Due _______ • Lab B: ALUM Metal Hydrate Formula • Complete Pre-lab • Due _________

  47. Quiz Analysis • Locate a DIFFERENT colored pen/pencil. • 3 minutes reflect individually. • 3 minutes communicate with a peer. • Questions?!

  48. Lab Report Revision Process • Different colored pen/pencil or separate sheet ATTACHED to draft 1. • Review • Individually • With Peer • With Mrs. Bechtum • DUE: Tuesday, January 20th! • Questions?

  49. Lab B: Analysis of ALUM • Pre-Lab Questions • Safety • Part 1: Sample Data

  50. Alum Lab Pre-lab Questions • 1. When measuring a melting point, why is it necessary to raise the temperature very slowly in the vicinity of the melting temperature? • 2. Washing soda is a hydrated compound whose formula can be written Na2CO3x H2O, where x is the number of moles of H2O per mole of Na2CO3. When a 2.123 g sample of washing soda was heated at 130 C, all of the water of hydration was lost, leaving 0.787 g of anhydrous sodium carbonate. Calculate the value of x.