Welcome ! • Review : Atomic structure and the periodic table. • Today: the mol, the law of conservation of mass, stoichiometry Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–1 Learning Objectives Today we will: Gain an understanding of where major Chemistry strands spiral throughout the curriculum. Explore the states of matter, the mol, and the Law of Conservation of Mass. Measure and analyze percent composition of a substance in a compound. Explore where an understanding of stoichiometry can be useful in all content areas. 2 Definition of Matter The two properties of matter are: •Mass – amount of matter •Volume – amount of space occupied •Therefore…matter is anything that has mass and takes up space. States of Matter • Solid – definite shape and volume • Liquid – definite volume, and takes the shape of the container • Gas – takes the shape and volume of the container • Plasma – a collection of high energy ions and electrons. • Go to lecture on the ideal gas law….. Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–5 Crushed Can Expt • When the can is heated, what occurs? • When the water vaporizes in the can, what happens to the volume of the can? • What phase change occurs to the water vapor in the can when the can is immersed in the cold water? The water ____________. • What happens to the pressure of the water vapor in the can when it condenses? • Would the same result occur if the can were immersed into the water with the open side upward? Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–6 Now we can introduce the Mole A mole is a unit of quantity. A mole is 6.02 x 1023 things. 6.02 x 1023 is known as Avogadro’s constant (NA) Number of atoms, molecules or fundamental units ÷ 6.02 x 1023 Number of moles (mol) × 6.02 x 1023 Why the Mole? Consider one molecule of water How many molecules in 2000 mL of water? 6.7 x 1025 molecules 6.7 x 1025 molecules is not a manageable number. Consider: We count eggs by the dozen We measure long periods of time in centuries. We measure long distances in our universe using light years. The Mole There are many ways of measuring large quantities that utilize large units. The mole is one such unit. The mole is the SI unit for chemical quantity used to count the particles in a sample of pure substance. One mole = 6.02x1023 particles. “One mole of anything = 6.02x1023 units of that thing” The Mole How many molecules of water in 2000 mL? 6.7 x 1025 molecules Or 111 mol meaning 111 moles of water molecules. This is a much more manageable number. Molar Mass Are you sitting comfortably? Mass and the Mole: Why 6.02×1023? 1 mole = 6.02×1023 was chosen because this was the number of carbon-12 atoms that has a mass of 12 grams. Pure carbon (a mix of isotopes) has a mass of 12.01 g per mole. We call this value MOLAR MASS Consider 1 atom of sulfur AR = 32.07 = 5.326×10-23 grams Consider 1 mole of sulfur = 6.02×1023 atoms = 32.07 grams 6.02×1023 is a REALLY BIG number Consider 6.02×1023 sheets of paper stacked. How many round trips to the Moon would this stack of paper be equivalent to? 8×1010 (eighty billion ) roundtrips. Create, pick, find a mole analogy of your own. Show the math. Language issue gram atomic mass gram molecular mass gram formula mass Aluminum MOLAR MASS AR = 26.98 gram atomic mass = 26.98 g mol-1 Carbon dioxide (CO2) MR = 44.00 gram molecular mass = 44.00 g mol-1 Sodium chloride (NaCl) MR = 58.44 gram formula mass = 58.44 g mol-1 We weigh chemical quantities in grams. The molar mass value for a substance allows us to determine the number of moles from a measured mass. Consider 5.68 g of MgCl2 Molar mass of MgCl2 = 95.21 g mol-1 mass number of moles = molar mass 5.68 g n =0.0597 mol -1 92.51 g mol Number of ÷ 6.02 x 1023 atoms, molecules or fundamental × 6.02 x 1023 units × molar mass MOLES MASS ÷ molar mass Molarity • Used to describe solutions • M = moles/L Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–19 Activity – Determine the Molar Mass of Butane Review safety concerns and protocols !!!!! 1. Identify at least 4 possible sources of experimental error 2. The accepted molar mass of butane is 58.0 g/mole. Determine percentage error for your data. 3. Use your data, assume a molar mass of 58.0 g/mol, and derive the value of R, the gas constant. 4. The molar mass of the gas is 58.0 g/mol. Determine the volume of one mole of butane at STP based on your data. What is the percentage error between your value and 22.4 L (the accepted value). Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–20 An Introduction… • C6H12O6 + O2 H20 + CO2 + Energy • What is wrong with this equation? • How did you learn to balance an equation? 21 Chemical Reactions: Equations • Writing chemical equations – A chemical equation is the symbolic representation of a chemical reaction in terms of chemical formulas. – For example, the burning of sodium and chlorine to produce sodium chloride is written 2Na Cl 2 2NaCl – The reactants are starting substances in a chemical reaction. The arrow means “yields.” The formulas on the right side of the arrow represent the products. Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–22 Chemical Reactions: Equations • Writing chemical equations – In many cases, it is useful to indicate the states of the substances in the equation. – When you use these labels, the previous equation becomes 2Na(s ) Cl 2 (g ) 2NaCl(s ) Balancing chemical reactions video link Balancing an equation activity….. Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–23 Law of Conservation of Mass Activity 24 Law of Conservation of Mass: • This experiment verifies the Law of Conservation of Matter: Matter is neither created or destroyed as a result of chemical changes but may be changed in form. The balanced equations are as follows: • 2NaOH (aq) + CuSO4 (aq) -----> Na2SO4 (aq) + Cu(OH)2 (s) • 4NH3 (aq) + CuSO4 (aq) -------> Cu(NH3)4SO4 (s) • Na2CO3 (aq) + CuSO4 (aq) -------> Na2SO4 (aq) + CuCO3 (s) 25 Law of Conservation of Mass: Content Blast • Matter is neither created nor destroyed in any process. • Changes in matter are always accompanied by changes in energy. • Energy changes operate under the Law of Conservation of Energy. • Atoms are neither created nor destroyed in a chemical reaction. • For mass to remain constant during a chemical reaction, the number of atoms of each element must be the same before and after a chemical reaction. • Equations are balanced using coefficients before the formulas for reactants and products. 26 Bubble Gum Activity An Introduction to Stoichiometry • Complete Task 1 and Procedures 1-3. – Take a 15 minute break while chewing the gum. • Return to your groups and complete the remainder of the activity. • NOTE the review hints. • Track and record content, pedagogical issues/examples, or diversity issues that may appear in the classroom. 27 Bubble Gum Activity Report Out Summary The manufacturer of one common type of bubble gum indicates on the package that each piece of bubble gum has a mass of 5 grams. The sugar content for each piece is 4 grams. Determine the percent by mass of sugar in this type of gum. How do your results compare to the manufacturer’s stated content? Using your data collected calculate the total mass of one piece of gum and the mass of sugar in one piece of gum. How many moles of sugar are present in one piece of gum (assume that the sugar is sucrose)? How many molecules of sugar are present in one piece of gum if the sugar is sucrose? Were any of the questions you or your group had answered by this activity? If so, which ones? 28 Bubble Gum Activity Report Out Summary • What questions did your group have prior to the activity? • What challenges did you experience while completing the activity? • Were your questions answered? If not, what adjustments could be made to the activity to help it better address your questions? 29 Bubble Gum Activity Content Blast • Percent composition is the mass of each component relative to the total mass of the substance. • Molar mass is the mass in grams of one mole of a substance. • The mole establishes a relationship between the atomic mass unit and the gram. • The mass in grams of 1 mole of a substance is numerically equal to its atomic mass or formulas mass in atomic mass units. • The number of molecules in a mole of any molecular compound is 6.02 x 1023. 30 Stoichiometry • Using a balanced chemical equation to determine how much reactant or product is consumed or produced in a chemical reaction. • 2 H2 (g) + O2 (g) 2 H2O (l) – How many moles of H2O are produced from the reaction of 2 moles of H2? – How many moles of O2 are required to produce 4 moles of H2O? Chemical equation = recipe Chemical Equations • 2 H2 (g) + O2 (g) H2O (l) 4 moles molecules + 2 moles molecules 4 moles molecules Coefficients in Chemical Equation represents mole ratios http://www.chemistry.ohio-state.edu/betha/nealChemBal/ Stoichiometry “Chemical Equation = recipe” • 2 bread slices + 1 cheese slice 1 sandwich – Suppose need to make as many cheese sandwiches as possible for a party and have 20 slices of bread...how many slices of cheese are needed? • By inspection of recipe (chemical equation) • Or use coefficient ratio (mole ratio) from equation: 2 bread slices = 1 cheese slice • 1cheese _ slice cheese slices = 20 bread slices x 2bread _ slices = 10 cheese slices needed Image source: http://www.fotosearch.com/clip-art/sandwich.html Stoichiometry • 4 Fe + 3 O2 2 Fe2O3 • How many moles of Fe2O3 will be produced from the reaction of 1.50 mol of iron? 2molFe 2O 3 • ….mol Fe2O3 = 1.50 mol Fe x = 0.75 mol 4molFe Fe2O3 Mole ratio is coefficient ratio Stoichiometry – using grams • Suppose the given is grams, not moles… • Consider the unbalanced reaction: Al (s) + Fe2O3 (s) Al2O3 (s) + Fe (l) If you react 4.0 g of Al with excess Fe2O3, how many moles of Fe will be produced?? • First balance the equation: 2 Al (s) + Fe2O3 (s) Al2O3 (s) + 2 Fe (l) • Chemical equation speaks to us in moles, not grams! So to use the mole ratio it gives us (2mol Al: 2 mol Fe), – first need to convert 4.0 g Al to moles of Al – Then take moles of Al and “convert” to moles of Fe using mole ratio from balanced equation (coefficient ratio) as was done on previous slides. Stoichiometry – using grams • Road map: Stoichiometry – using grams • Suppose we wanted to solve for grams, not moles 2 Al (s) + Fe2O3 (s) Al2O3 (s) + 2 Fe (l) • Question: If you react 4.0 g of Al with excess Fe2O3, how many grams of Fe will be produced? • Road map is same as before, except one last step to convert mol Fe to g Fe. – Remember that to do any stoichiometry problem, you must be in the language of moles since the chemical equation (source of relationship or conversion factor) only speaks to us in moles, not grams! Stoichiometry – using grams • Road map: g → mol → mol → g Demonstration – Limiting Reagents Mg + 2 HCl (aq) MgCl2(aq) + H2 (g) 1. 100 mL of 1.0 M HCl would have how many mol HCl present? 2. From the balanced equation, we can do a calculation that will show 0.1 mole HCl requires exactly how many Mg to completely react? 0.1 mol HCl x 1 mole Mg x 24.3 g Mg = 1.2 g Mg 2 mol HCl 1 mole Mg This is the amount in the Red balloon! Therefore, both the HCl and the Mg in the 2nd flask are completely consumed! Cool! 3. How much H2 should be produced in the Red balloon? 1.20 g Mg x 1 mole Mg 24.3 g Mg x 1 mole H2 1 mole Mg = 0.050 mol H2 39 Demonstration – Limiting Reagents Mg + 2 HCl (aq) MgCl2(aq) + H2 (g) 4. How much hydrogen should be produced by the Blue balloon flask? Since there was 0.60 Mg in the blue balloon flask, the Mg will limit the amount of hydrogen produced. Therefore, use the amount of Mg to determine the amount of H2 produced. 0.60 g Mg x 1 mole Mg 24.3 g Mg x 1 mole H2 1 mole Mg = 0.025 mol H2 (Blue balloon) This is half the amount of gas formed in the red balloon. This explains why the blue balloon is half the size of the red balloon. Demonstration – Limiting Reagents Mg + 2 HCl (aq) MgCl2(aq) + H2 (g) 5. How much HCl would be needed to react with the 0.60 g Mg? 0.60 g Mg x 1 mole Mg 24.3 g Mg x 2 mole HCl 1 mole Mg x ___1 L___ = 50 mL! 1 mole HCl That means 50 mL of HCl are left over un-reacted in the 1st flask! 6. What about the 3rd flask with the Yellow balloon? You can do similar calculations to those shown previously. Content Specific Activities Report Out Summary • What questions did your group have prior to the activity? • What challenges did you experience while completing the activity? • Were your questions answered? If not, what adjustments could be made to the activity to help it better address your questions? 42 Day Two Learning Objectives Review Today we: Gained an understanding of where major Chemistry strands spiral throughout the curriculum. Explored the Law of Conservation of Mass. Measured and analyzed percent composition of a substance in a compound. Explored stoichiometry 43