ChemH Unit 7 Packet 2025 PDF

Name:________________________________________ Period:______ Score: _____/6 ChemH Unit 7 Packet - Stoichiometry Student Teacher Description Score ​ ​ WS#1: Percent composition (p. 2)...

Name:________________________________________ Period:______ Score: _____/6 ChemH Unit 7 Packet - Stoichiometry Student Teacher Description Score ​ ​ WS#1: Percent composition (p. 2) ​ ​ WS#2: Percentage composition and Empirical formula WS (p. 3-4) ​ ​ WS#3:Reaction Stoichiometry (p. 5-6) ​ ​ WS#4: Intro to Limiting and excess reactant stoichiometry (p. 7) ​ ​ WS#5: Limiting reactant and percent yield practice (p. 8-9) ​ ​ WS#6: Word Equations and Stoichiometry (p. 10-11) U7 Lecture Notes (p. 12-18, not graded) ***SHOW ALL WORK FOR FULL CREDIT*** *If you did your work on another paper, attach it to the back of this packet. DATE WARM-UP/RETRIEVAL PRACTICE 1 Name: ________________________________ Class Per: _____ WS#1 Percent composition - SHOW ALL OF YOUR WORK FOR CREDIT Find the percent compositions of all of the elements in the following compounds: 1) CuBr2 Cu: ___________ Br: ___________ 2) (NH4)2S N: ___________ H: ___________ S: ___________ 3) Mg(NO3)2 Mg: ___________ N: ___________ O: ___________ 4) (NH4)3PO4 N: ___________ H: ___________ O: ___________ P: ___________ 5) Al2(SO4)3 Al: ___________ S: ___________ O: ___________ 2 Name: ________________________________ Class Per: _____ WS#2: Percentage composition and Empirical formula WS - SHOW YOUR WORK FOR CREDIT 1.​ What is the percent composition by mass for each element in CO2? 2.​ What is the percent composition by mass for each element in Na3PO4? Done in class 3.​ What is the percent composition by mass for each element in glucose, C6H12O6? Done in class 4.​ A chemist determines that 1.26g of iron reacts with 0.54g of oxygen to form rust. What is the percent composition of each element in the new compound? 5.​ The following compounds are used in the production of fertilizers. Determine which has the highest percentage by mass of nitrogen: NH3, CO(NH2)2, (NH4)2SO4. 3 Name: ________________________________ Class Per: _____ 6.​ Give the empirical formulas of the following compounds: a.​ C2H2 b.​ C6H12O6 c.​ C12H22O11 d.​ C8H18 e.​ C8H14 7.​ A sample of a compound contains only the elements sodium, sulfur and oxygen. It is found by analysis to contain 0.979g Na, 1.365g S, and 1.021g O. Determine its empirical formula. 8.​ A street drug has the following composition: 83.89% C, 10.35% H, 5.76% N. Determine its empirical formula. 9.​ A compound with an empirical formula of C2OH4 and a molecular mass of 88 g/mol. What is its molecular formula? 10.​Benzene is a hydrocarbon, a compound of carbon and hydrogen only. It is found to contain 7.74% H by mass. a.​ Determine the empirical formula of the compound. b.​ The molar mass of the compound is 78.10 g/mol. Determine the molecular formula of the compound. 11.​A weak acid has a molar mass of 162 g/mol. Analysis of a 0.8821g sample showed the composition by mass is 0.0220g H, 0.3374g P, and the remainder was O. Determine its empirical and molecular formulas. 4 Name: ________________________________ Class Per: _____ WS#3: Reaction Stoichiometry - SHOW YOUR WORK FOR FULL CREDIT 1.​ For this reaction: _____NH3 + _____O2 → _____NO + _____H2O a.​ How many moles of oxygen (O2) are needed to react with 3.24 moles of ammonia (NH3)? b.​ How many moles of water (H2O) are produced from 12.8 moles of oxygen (O2)? 2.​ For this reaction: ______Fe3O4 + ____CO → _____Fe + _______CO2 a.​ How many moles of carbon dioxide (CO2) are produced from 2.87 moles of Fe3O4? b.​ How many moles of carbon monoxide (CO) are needed to react with 8.25 moles of Fe3O4? 3.​ _____NH3 + ______O2 → ______N2O + ______H2O If 80.0 grams of O2 are reacted in the above reaction, how many grams of N2O will be produced? 4.​ _______MnO2 + ______HCl → ______Cl2 + ______MnCl2 + ______H2O Given 145.7 grams of manganese (IV) oxide (MnO2), how many grams of hydrochloric acid (HCl) is needed to use up the MnO2 completely? 5.​ _______Pb(NO3)2 + ______NaCl → ______PbCl2 + _______NaNO3 How many grams of lead (II) nitrate (Pb(NO3)2) are needed to produce 49.35 g of sodium nitrate (NaNO3)? 5 Name: ________________________________ Class Per: _____ 6.​ ________Fe + _____H2O → _________Fe2O3 + _______H2 If 42.84 grams of iron are reacted with water (H2O), how many grams of Fe2O3 are produced? 7.​ ______F2 + ______AlBr3 → _______Br2 + ______AlF3 If 8.4 grams of AlBr3 is reacted with F2, how many grams of AlF3 are produced? 8.​ _______HCl + ______Cu → ______CuCl2 + ______H2 How many moles of H2 are produced when 67.12 g of HCl are reacted with Cu? 9.​ _______NaClO3 → ______NaCl + _____O2 How many molecules of O2 are produced in the decomposition reaction of 145.73 g of NaClO3? 10.​_______Cu + ______AgNO3 → _______Ag + _______Cu(NO3)2 How many grams of silver are produced when 0.75 moles of Cu displaces silver from AgNO3 in the above reaction? 6 Name: ________________________________ Class Per: _____ WS#4: Intro to Limiting and excess reactant stoichiometry - SHOW YOUR WORK FOR CREDIT 1.​ Consider the following reaction: 2Al + 6HBr → 2AlBr3 + 3H2 a.​ When 3.22 moles of Al react with 4.96 moles of HBr, how many moles of H2 are formed? b.​ What is the limiting reactant? c.​ How much of the excess reactant is left over? 2.​ Consider the following reaction: 2 CuCl2 + 4 KI → 2 CuI + 4 KCl + I2 a.​ When 0.56 grams of CuCl2 reacts with 0.64 grams of KI, how many grams of I2 are formed? b.​ What is the limiting reactant? c.​ How many grams of the excess reactant is left over? 3.​ Consider the following reaction: 4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2 a.​ When 26.62 grams of FeS2 reacts with 5.44 grams of O2, how many grams of SO2 are formed? b.​ How many grams of the excess reactant is left over? 7 Name: ________________________________ Class Per: _____ WS#5: Limiting reactant and percent yield practice - SHOW ALL WORK FOR FULL CREDIT 1.​ Consider the following reaction: P4 + 6 Cl2 → 4 PCl3 a.​ When 15.2 moles of P4 react with 44.50 moles of Cl2, how many moles of PCl3 are formed? b.​ What was the limiting reactant? c.​ How much of the excess reactant was left over? 2.​ Consider the following reaction: 4 Al + 3 O2 → 2 Al2O3 a.​ When 58.5g Al and 98.0g O2, how many grams of Al2O3 are formed? b.​ What is the limiting reactant? c.​ How many grams of the excess reactant is left over? 3.​ Consider the following reaction: 2 H2O + 2 Na → 2 NaOH + H2 a.​ If 10.0g H2O reacts with 4.5g Na, how many grams of NaOH are formed? b.​ What is the limiting reactant? c.​ How many grams of the excess reactant is left over? 4.​ Calculate the percent yield for the reaction: P4 + 6 Cl2 → 4 PCl3 if 75.0 g of P4 reacts with excess Cl2 to produce 111.0 g of PCl3 8 Name: ________________________________ Class Per: _____ 5.​ 3Fe(OH)2 (aq) + 2 H3PO4 (aq) → Fe3(PO4)2 (s) + 6 H2O (l) a.​ If 3.20 g of Fe(OH)2 is reacted with 2.50 g of H3PO4, what is the limiting reagent and what is the reactant in excess? b.​ How many grams of Fe3(PO4)2 precipitate can be formed? c.​ If 3.99 g of Fe3(PO4)2 is actually obtained, what is the percent yield? 6.​ When octane (C8H18) is burned in oxygen, carbon dioxide and water are produced. If 320g of octane is burned and 392g of water is recovered, what is the percent yield of the experiment? The equation is: 2 C8H18 + 25 O2 →16 CO2 + 18 H2O. 7.​ A student places an iron nail with a mass of 2.32g into a flask of CuSO4. The nail reacts completely, leaving a quantity of copper metal in the bottom of the flask. The student finds the mass of the recovered copper to be 2.51g. The equation for this reaction is Fe + CuSO4 → FeSO4 + Cu. a. What is the expected yield? b. What is the percent yield? 9 Name: ________________________________ Class Per: _____ WS #6 Word Equations and Stoichiometry SHOW ALL WORK FOR FULL CREDIT 1.​ Hydrogen gas reacts in a synthesis reaction with oxygen gas to produce water. a.​ Write out the balanced equation. b.​ If 10.0 grams of hydrogen gas completely reacts, how many grams of oxygen gas were required? c.​ How many grams of water were produced? 2.​ Magnesium hydroxide decomposes into magnesium oxide and water. If 98.03 g of Magnesium hydroxide decompose, how many molecules of water are produced? 3.​ Aluminum sulfate reacts in a double replacement reaction with calcium hydroxide. How many grams of aluminum hydroxide are produced if 165.7 grams of Aluminum sulfate are fully reacted? 4.​ Lead (II) nitrate reacts with sodium chloride in a double replacement reaction. How many moles of Lead II nitrate are needed to produce 49.35 g of sodium nitrate? 5.​ Chlorine gas reacts in a single replacement reaction with sodium bromide. If 23.5 grams of Chlorine gas fully reacts, how many grams of each product will be produced? 10 Name: ________________________________ Class Per: _____ 6.​ Calcium oxide reacts with carbon dioxide to produce calcium carbonate in a combination reaction. a.​ If 0.50 moles of calcium oxide reacts with 0.40 moles of carbon dioxide, how many moles of calcium carbonate are formed? 7.​ Nitrogen gas reacts with hydrogen gas in a combination reaction to produce 10.5 g of ammonia (NH3). Determine the percent yield for the reaction if 15.0g of nitrogen gas is reacted with 15.0g of hydrogen gas. 8.​ Determine the percent yield for a double replacement reaction between 2.80g of aluminum nitrate and excess sodium hydroxide if 0.996g of aluminum hydroxide is recovered. 9.​ The % yield of the reaction between solid aluminum and phosphoric acid (H3PO4) to produce aluminum phosphate and hydrogen gas is 73.7%. Theoretically, how many grams of Al must be used to react with an excess of H3PO4 to give 30.0g of aluminum phosphate with that percent yield? 11 Name: ________________________________ Class Per: _____ ChemH Unit 7 Notes Packet Unit 7 Learning Objectives: ​ Apply first semester learnings by converting word equations into balanced chemical equations, including recognizing ionic and covalent compound names, the 7 elements that form diatomics (H2, N2, O2, F2, Cl2, Br2, I2), the names and charges of common polyatomic ions, recognizing the 5 main types of reactions, and balancing chemical equations. ​ Use chemical formulas to calculate the percent composition of each type of atom in a molecule. ​ Calculate the empirical formula of a substance based on given mass quantities or percent composition. ​ Calculate the molecular formula of a substance based on given mass quantities or percent composition and molar mass. ​ Calculate the mass of reactants required to produce a given mass of products and vice versa (gram to gram conversion), based on the mole ratios in a balanced chemical equation. ​ Identify the limiting reactant and excess reactant in a chemical reaction based on the mole ratios in the balanced chemical equation and given mass or mole quantities. ​ Calculate the percent yield of a chemical reaction based on the mole ratios in the balanced chemical equation, theoretical yield, and experimentally determined actual yield. Unit 7 Lecture Notes: ​ Review: Moles ○​ Activity: Go to a lab station with a partner and count exactly 8.0 x 1022 molecules of sugar or salt (depending on what is at your station!) ​ Circle the substance you have:​ Sugar​ Salt ​ Write the chemical formula for this substance: ________________________ ​ Make a plan: How will you count this number of molecules of this substance? ​ Show your collected data and calculations below: ​ Talk to another group that had the other substance. Did they have the same mass for the same number of molecules? ___________ ​ Announcement: During 2nd semester, you will be required to apply what you learned in 1st semester by translating word equations into balanced chemical equations: ○​ Recognize names of ionic and covalent compounds, and convert them into chemical formulas, including the 7 elements that form diatomics (H2, N2, O2, F2, Cl2, Br2, I2) and the names and charges of the transition metals and all of the polyatomic ions on your sheet ○​ Recognize the 5 types of reactions ○​ Balance chemical equations 12 Name: ________________________________ Class Per: _____ ​ Diatomics ○​ Diatomic molecules consist of __________ atoms bonded together covalently in a molecule. ○​ 7 pure elements naturally form diatomics (H2, N2, O2, F2, Cl2, Br2, I2). ​ Mnemonic: “I Bring Clocks For Our New Home” ○​ For example, when you see “X reacts with hydrogen to form…”, that translates to “X + H2 → …” ​ Percent Composition ○​ Review: Writing chemical formulas ​ Write the chemical formula for the compound formed when Magnesium and Chlorine get together: ________________ What type of compound is this (circle one): ionic or covalent? What is it named? ______________________________ ​ Summarize how you determined the chemical formula of the ionic compound made from magnesium and chloride ions. ○​ Question: How much of a sample of this ionic compound is made up of magnesium, and how much of chlorine? ​ Calculate the percent composition by mass! Show all work: ​ ​ Practice: Calculate the percent composition of each of the following. Show all work. ​ Na2O ​ NaOH ​ Na3PO4 ​ C5H10O5 ​ C6H12O6 13 Name: ________________________________ Class Per: _____ ​ Empirical and Molecular Formula ○​ Question: How can we determine the formula of a compound from experimental data? ​ In a 100g sample, you determine that 26.492g is Mg 73.508g is Cl. Determine the formula for this compound. Show your work: ​ ​ ○​ Practice: Show all work ​ What is the empirical formula of a compound that has 18.7g of C and 6.28g of H? ​ A sample of an unidentified compound contains 29.84 g sodium, 67.49 g chromium, and 72.67 g oxygen. Find the empirical formula. ​ A sample of a compound has mass percentages of 40.00% C, 6.71% H and 53.28% O. What is the formula? ○​ By converting the element masses in a substance to moles, you can find only the ________________ ratio for the formula. This is called the _________________________. Multiple compounds can have the same empirical formula. ○​ In order to know what the true formula of the compound is, we need to know the ______________________. ​ Example: The molar mass of a substance with empirical formula CH2O is 180 g/mol. Determine the molecular formula. Show all work. 14 Name: ________________________________ Class Per: _____ ​ Stoichiometry - Be sure to SHOW ALL WORK! ○​ Review: Writing balanced chemical equations from a word equation: ​ Aluminum metal reacts with oxygen to produce aluminum oxide. Write the balanced chemical equation below: ​ ​ If we have 24 atoms of aluminum, how many molecules of aluminum oxide will we get? ​ ​ How many moles of aluminum oxide can we produce using 8 moles of aluminum? ​ ○​ Mass to mass conversions using mole ratios from a balanced chemical equation: ​ Example: How many grams of aluminum oxide can we produce using 10.0 grams of Aluminum? (show your work and don’t forget sig figs!) ​ Practice: Nitrogen and hydrogen react to form ammonia. ​ Write the balanced chemical equation: ○​ ​ When 38g of N2 are used, how many grams of NH3 are produced? Show all work. ○​ ​ ​ Limiting Reactants ○​ Chemical equations are like __________________. They tell you how many molecules and/or atoms of each reactant you need to make the products. ○​ Example: Making a cheese sandwich. Fill in the vocabulary in the blanks using this model. ​ Given: 1 sandwich requires 2 slices of bread and 1 slice of cheese ​ Write the Balanced Equation: ​ Question: How many sandwiches can we make with 8 slices of bread and 3 pieces of cheese? ​ What is the limiting reactant?_____________________ ​ What was in excess (and by how much)? ________________ ​ Question: A loaf of bread from the grocery store had 28 slices of bread. 1 pack of cheese slices had 16 slices. How many sandwiches can we make using this material? ​ What is the limiting reactant?_____________________ ​ What was in excess (and by how much)? ________________ 15 Name: ________________________________ Class Per: _____ ○​ When we run out of the ___________________________, the reaction stops. It doesn’t matter that we have _______________ of the other reactant, it has nothing to react with! ○​ Example: Methane (aka carbon tetrahydride), the primary fuel found in natural gas, is undergoing combustion. ​ What type of compound is methane? __________________ ​ Write the balanced chemical equation. ​ What could limit this reaction? _________________________ ​ Question: If we have 4 moles of methane and 2 moles of oxygen: ​ What is the limiting reactant? __________________ ​ What is the excess reactant? __________________ ​ How much of the excess reactant is left over? _________________ ​ How much carbon dioxide is formed? _________________ ​ How much water is formed? ____________________ ○​ Process: Determining limiting and excess reactants 1.​ Write the balanced chemical equation. 2.​ Pick one of the reactants and use dimensional analysis to figure out how much of the other reactant you would need. 3.​ If you have less than what you need of the second reactant, then the second reactant is limiting. 4.​ If you have more than what you need of the second reactant, then the second reactant is in excess. ○​ Practice: 1.​ Hydrogen reacts with oxygen to form water. If we have 5 moles of of hydrogen and 2 moles of oxygen: ​ Write the balanced chemical equation: ​ What is the limiting reactant? __________________ ​ What is the excess reactant? __________________ ​ How much of the excess reactant is left over? _________________ ​ How much carbon dioxide is formed? _________________ ​ How much water is formed? ____________________ 2.​ Nitrogen and hydrogen react to form ammonia (nitrogen trihydride). If we have 6 moles of nitrogen and 21 moles of hydrogen: ​ Write the balanced chemical equation: 16 Name: ________________________________ Class Per: _____ ​ What is the limiting reactant? __________________ ​ What is the excess reactant? __________________ ​ How much of the excess reactant is left over? _________________ ​ How much carbon dioxide is formed? _________________ ​ How much water is formed? ____________________ 3.​ Now try it with masses (show all work!)! If we have 6.0 g of nitrogen and 18.0 g of hydrogen: ​ What is the limiting reactant? __________________ ​ What is the excess reactant? __________________ ​ How much of the excess reactant is left over? _________________ ​ How much carbon dioxide is formed? _________________ ​ How much water is formed? ____________________ ​ Percent Yield ○​ So far we have assumed that all of the reactants are used in a reaction, and the full amount of product is produced. This is called __________________ yield. However, this never happens in real life! What we collect in a lab is called the _____________ yield, and is less than we expected. We can use these 2 values to calculate the ___________ yield, which tells us how effective we were in collecting the product! ​ Percent yield = Actual yield * 100 Theoretical yield ○​ Example: A student reacted 6.0 g of N2 with 18.0 g of H2, and obtained 6.5g of NH3 in their experiment. What was their percent yield? Show all work. 17 Name: ________________________________ Class Per: _____ ○​ Practice: When 24.0 g of potassium nitrate was heated with lead, 13.8 g of potassium nitrite was formed. Calculate the percent yield of potassium nitrite. Show all work. 18