Student Version Notes: Unit 5 Moles & Stoichiometry

Transcription

1 Name: Regents Chemistry: Mr. Palermo Student Version Notes: Unit 5 Moles & Stoichiometry

2 Name: KEY IDEAS A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system. (3.1cc) Types of chemical formulas include empirical, molecular, and structural. (3.1ee) The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound. (3.3d) In all chemical reactions there is a conservation of mass, energy, and charge. (3.3a) A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. (3.3c) The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram formula mass) of a substance equals one mole of that substance. (3.3e) The percent composition by mass of each element in a compound can be calculated mathematically. (3.3f) Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement. (3.2b) VOCABULARY For each word, provide a short but specific definition from YOUR OWN BRAIN! No boring textbook definitions. Give an example if you can. Don t use the words given in your definition! Mole: Molar Mass (gfm): Reactants: Products: Coefficients (in reactions): Conservation of mass: Empirical Formula: Molecular Formula: Synthesis reaction: Decomposition reaction: Double replacement reaction: Single replacement reaction:

3 Name:

4 Unit 5 - Lesson 1: Moles and Molar Mass Objective: Calculate Molar Mass (gram formula mass) STOICHIOMETRY: The study of the amounts of reactants and products in chemical reactions (the study of chemical recipes) COUNTING ATOMS Subscripts: Coefficients: REVIEW OF Ex. Ca3(PO4)2 PRACTICE: How many oxygen atoms in each? 1. NH4NO3 2. C8H8O4 3. O3 4. C3H5(NO3)3 The MOLE (mol): Why we use it: 1 dozen eggs = 1 mol of eggs = Example: 3.5 dozen roses = 3.5 mol roses = 1

5 Unit 5 - Lesson 1: Moles and Molar Mass MOLAR MASS: **Element s molar masses are reported on the periodic table Calculating Molar Mass (Gram Formula Mass GFM) Molar Mass Examples: Elements 1. What is the molar mass of iron? 2. What is the molar mass of copper? Molar Mass Examples: Compounds 1. What is the molar mass of water? 2. What is the gram-formula-mass of CAF 2? PRACTICE: What is the molar mass (GFM) of Calcium(s)? PRACTICE: What is the GFM of Br2(g)? 2

6 Unit 5 - Lesson 1: Moles and Molar Mass PRACTICE: What is the GFM of CO2(g)? CHECK YOUR UNDERSTANDING 1: What is the GFM of Fe2O3 (s)? CHECK YOUR UNDERSTANDING 2: What is the GFM of Iron (II) oxide? 3

7 Unit 5 - Lesson 2: Calculating Moles Objective: Calculate the number of moles given the grams Calculate the number of grams given the moles Use MOLE FORMULA from Table T Given mass is the mass in grams How to Convert from Grams to Moles 1. Calculate the GFM for the compound. (round to nearest tenth) 2. Plug the given value and the GFM into the mole calculations formula and solve for the number of moles. EXAMPLE 1: How many moles are in 39.0 grams of LiF? EXAMPLE 2: What is the number of moles of potassium chloride present in 148 g? PRACTICE: Determine the number of moles in 8.0 grams of Boron 4

8 Unit 5 - Lesson 2: Calculating Moles PRACTICE: How many moles are in 168 g of KOH? How to Convert from Moles to Grams: Use the same mole formula to solve for grams by setting mole over 1 and cross multiplying CHECK: If more than 1 mol mass is GREATER than GFM If less than 1 mol mass is SMALLER than GFM EXAMPLE: What is the mass of 4.5 moles of KOH? EXAMPLE: What is the mass of 0.50 mol of CuSO4? 5

9 Unit 5 - Lesson 2: Calculating Moles PRACTICE: Calculate the mass of 6.70 moles of carbon. PRACTICE: Calculate the mass of 22.35g of MgSO4 CHECK YOUR UNDERSTANDING 2: Calculate the mass of moles of calcium phosphate. 6

10 Unit 5 - Lesson 3: Mole Ratios Objective: Calculate mole ratios in a chemical equation Mix + 2 Eggs + 1 cup water = Reactants Cake Products Just like a recipe in a chemical Rx you can double, triple, halve etc. the amounts of the ingredients to change how much product you make Example: 2 Mix = 2 Cakes 6 Eggs = 3 Cakes COEFFICIENTS: How many MOLES of the substance are needed in a reaction. 4Al + 3O2 2Al2O3 reactants products CALCULATING MOLE RATIOS: 1. Circle the substances involved 2. Set up a ratio of moles (proportion) of SUBSTANCES in the balanced equation to the ACTUAL MOLE values and solve for the unknown EXAMPLE: How many moles of oxygen are consumed when 0.6 moles of hydrogen burns to produce water? 2 H2(g) + O2(g) 2 H2O 7

11 Unit 5 - Lesson 3: Mole Ratios EXAMPLE: How many moles of nitrogen gas (N2) would be needed to produce 10 moles of ammonia (NH3) in the following reaction? N2 + 3H2 2NH3 PRACTICE: If 12 moles of C3H8 react completely, how many moles of H2O are formed in the reaction below? C3H8 + 5O2 3CO2 + 4H2O PRACTICE: If 20 moles of CO2 are formed, how many moles of O2 reacted in the reaction below? C3H8 + 5O2 3CO2 + 4H2O PRACTICE: If 3.75 moles of aluminum react, how many moles of sulfur are needed? 16 Al + 3S8 8Al2S3 CHECK YOUR UNDERSTANDING: If 2.5 moles of N2 react completely, how many moles of NH3 are formed? N2 + 3H2 2NH3 8

12 unit 5 - Lesson 4: Balancing Reactions Objective: Balance chemical reactions using coefficients CONSERVATION OF MASS **Matter and energy cannot be created nor destroyed, only changed from one form to another BALANCING EQUATIONS Balance by changing coefficients only **NOTE: STEPS FOR BALANCING EQUATIONS 1. Start with the element that is only found once on both sides. 2. Keep polyatomic ions together. Count as a unit if not broken up. 3. Coefficients must be smallest possible whole number (reduce if not) 4. Check each atom to see if its balanced 9

13 Unit 5 - Lesson 4: Balancing Reactions EXAMPLE: Al + CuCl2 Cu + AlCl3 EXAMPLE: N2 + H2 NH3 PRACTICE: H2 + O2 H2O PRACTICE: Li + O2 Li2O PRACTICE: BaCl2 + AgNO3 Ba(NO3)2 + AgCl CHECK YOUR UNDERSTANDING: Pb(NO3)2 + K2CrO4 PbCr2O4 + KNO3 10

14 Unit 5 - Lesson 4: Balancing Reactions 1. Synthesis: A + 2B AB2 2. Decomposition: AB2 A + 2B TYPES OF REACTIONS 3. Combustion: CH4 + O2 CO2 + H2O 4. Single Replacement: AB + C CB +A 5. Double Replacement: AB + CD AD + CB * Notice synthesis and decomposition are opposites. Also, combustion can have any carbon compound as a reactant. PRACTICE: Identify the type of reaction 1.) 4Li + O2 2 LI2O 2.) Al + CuCl2 AlCl3 + Cu 3.) BaCl2(aq) + Na2SO4(aq) 2NaCl (aq) + BaSO4 (aq) 4.) 2AuBr3 2Au + 3Br2 11

15 Unit 5 - Lesson 5: Determining empirical and molecular formulas Objective: Determine the empirical formula from the molecular formula Determine the molecular formula from the empirical formula EMPIRICAL FORMULA: Example: B10G20 Example: C2H6 MOLECULAR FORMULA: Whole # multiple of empirical If the empirical formula is CH4 a molecular formula could be CH4, C2H8, C3H12 etc. Examples: 12

16 Unit 5 - Lesson 5: Determining empirical and molecular formulas Divide subscripts by the greatest common factor DETERMINE THE EMPIRICAL FORMULA Example: molecular formula = C4H10 Divide by 2 (greatest common factor) Answer: PRACTICE: Determine empirical formula from molecular formula. 1. C6H12O6 2. N2O4 3. BaCl2 4. C2H6 5. CH3 CALCULATING MOLECULAR FORMULA FROM EMPIRICAL FORMULA 1. Calculate Gram Formula Mass of the EMPIRICAL FORMULA 2. Divide molecular mass given by the empirical formula mass 3. Multiply all of the subscripts in the empirical formula by the number (multiple) you calculated in step 2 EXAMPLE: The empirical formula for ethylene is CH2. Find the molecular formula if the molecular mass is 28.1 g/mol? 13

17 Unit 5 - Lesson 5: Determining empirical and molecular formulas PRACTICE: What is the molecular formula of a compound that has an empirical formula of NO2 and molecular mass of 92.0 g? PRACTICE: A compound has an empirical formula of HCO2 and a molecular mass of 90 grams per mole. What is the molecular formula of this compound? PRACTICE: If a compound has a mass of 45 g/mol and an empirical formula of CH3, what is the molecular formula? CHECK YOUR UNDERSTANDING: What is the molecular formula of a compound that has a molecular mass of 48g and an empirical formula of CO2? 14

18 Unit 5 - Lesson 6: Calculating Percent Composition Objective: Calculate Percent Composition Calculate Percent of Water in a hydrate PERCENT COMPOSITION: Formula located on Table T CALCULATING PERCENT COMPOSITION 1. Calculate GFM of compound (mass of whole) 2. Plug into % composition formula EXAMPLE: What is the percent composition of Calcium in CaCl2? PRACTICE: What is the percentage by mass of carbon in CO2? PRACTICE: What is the percent by mass of nitrogen in NH4NO3? 15

19 Unit 5 - Lesson 6: Calculating Percent Composition CHECK YOUR UNDERSTANDING 1: What is the percent by mass of magnesium in MgCl2? HYDRATES Hydrate- It is written like this: Ionic Compound s Formula n H2O (n) is a whole number Ex. CuSO4. 5H2O Anhydride(anhydrate)- CALCULATING PERCENT OF WATER IN A HYDRATE 1. Calculate GFM of HYDRATE including the water (mass of whole) 2. Plug into % composition formula T EXAMPLE: What is the percentage by mass of water in sodium carbonate crystals (Na2CO3 10H2O)? Step 1- Calculate GFM of Hydrate Step 2- Plug values into Formula 16

20 Unit 5 - Lesson 6: Calculating Percent Composition PRACTICE: What is the percent by mass of water in BaCl2 2H2O? EXAMPLE: (Calculating % mass of water in the lab) A 12.2g sample of a hydrate was heated to a constant mass of 10.2 grams. What is the percent by mass of the water in the hydrate? Water is heated (evaporated) off so the decrease in mass is the mass of the water CHECK YOUR UNDERSTANDING: A gram sample of hydrated crystal is heated to a constant mass of 8.72 grams. This means all of the water has been evaporated (lost) by the heat. a) Calculate the mass of water that was lost to evaporation b) Calculate the % mass of water in the hydrate. 17