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Chapter 3

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Chapter 3

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  1. Chapter 3 Chemical Reactions and Earth’s Composition

  2. The Development of Earth • Matter in the Universe condensed into planets. • The planets closer to the sun have different chemical compositions than the rest of the universe. • Many of the volatile chemicals were lost from these planets.

  3. The Composition of Compounds • The law of definite proportions states that a specific chemical compound obtained from any source always contains the same proportion by mass of its elements • H+ + OH- ---> H2O • 2 H2 + O2 ---> 2 H2O

  4. The Composition of Compounds • The law of multiple proportions states that the masses of element Y that combine with a fixed mass of elements X to form two or more different compounds are in the ratios of small whole numbers. • Examples: NO, NO2, N2O, N2O5, etc.

  5. Chemical Equations • Reactants ↔ Products

  6. Chapter 4 Definitions Chemical Reaction - A process in which substances are changed into other substances through rearrangement, combination, or separation of atoms. Chemical Equation - A written representation of a chemical reaction, showing the reactants and products, their physical states, and the direction of the reaction. Reactants - The starting material in a chemical reaction or equation. Products - The substances formed in a chemical reaction or equation. Physical States - solids (s), liquids (l), gases (g) and aqueous (aq). Balanced Chemical Equation - A written representation of a chemical reaction that gives the relative amounts of the reactants and products, their physical states, and the direction of the reaction.

  7. Balanced Chemical Equations • The reactants appear on the left and the products appear on the right. The adjoining arrow shows the direction of the reaction. • 2. The phase of the reactant or product is written after the chemical symbol and is in parentheses. • 3. An integer precedes the chemical formula of each substance. This number, known as the stoichiometric coefficient, is the smallest integer that allows the equation to be balanced. • 4. Matter and charge are conserved in balanced chemical equations.

  8. Chemical Reactions

  9. Avogadro number and the mole. • NA = 6.0221367 E23 mol-1 • A mole is the amount of a substance that contains as many elementary particles (atoms, molecules, or whatever) as there are in exactly 12 g of the carbon-12 isotope. • 1 mol = 6.0221367 E23 particles • Problem • Express the following estimates for the year 2010 in nanomoles. • USA. 298. million • China. 1.34 billion • c: US Debt (to China): 3 Trillion

  10. Mass ↔ moles ↔ molecules ↔ atoms • Problem 39 • Aluminum, silicon, and oxygen form minerals known as aluminosilicates. How many moles of aluminum are in 1.50 moles of: • pyrophyllite, Al2Si4O10(OH)2 • b. Mica, KAl3Si3O10(OH)2

  11. Problem 44. How many moles of O2- ions are in 0.55 mol of Aluminum oxide? What is their mass in grams? Answers: 1.65 mol O2-, 26.4 grams

  12. Conversions • Converting between a number of particles and an equivalent number of moles (or vice versa) is a matter of dividing (or multiplying) by Avogadro’s number.

  13. Molar Mass • The average mass of an atom of helium is 4.003 amu, and the mass of a mole of helium (6.022 x 1023 atoms of He) is 4.003 g. • The molar mass (M) of helium is 4.003 g/mol.

  14. Molar Mass • A substance’s molar massis the mass in grams of one mole of the compound. CO2 = 44.01 grams per mole C + 2O 12.01 + 2(16.00) = 44.01g/mol

  15. Problem. Calculate the molar masses of the following: a. sucrose, C12H22O11 Answer: 342.31 d. fructose, C6H12O6 Answer: 180.158

  16. Mole Calculations • How many moles of Ca atoms are present in 20.0 g of calcium? • How many Cu atoms are present in 15.0 g of copper?

  17. Mole Calculations • How many grams are present in 3.40 moles of nitrogen gas (N2)? • How many molecules are present in 5.32 moles of chalk (CaCO3)? • How many oxygen atoms are present in this sample?

  18. Mole Calculations The uranium used nuclear fuel exists in nature in several minerals. Calculate how many moles of uranium are found in 100.0 grams of carnotite, K2(UO2)2(VO4)2•3H2O. ans: 902.176 g/mol; 0.4434 moles of U.

  19. 100 gram of K2(UO2)2(VO4)2•3H2O

  20. The law of conservation of mass statesthat the sum of the masses of the reactants of a chemical equation is equal to the sum of the masses of the products. Law of Conservation of Mass

  21. Chemical reactions follow the law of conservation of mass. Chemical Change

  22. Balanced Chemical Equations • Chemical equations should be balanced to follow the law of conservation of mass. • Total mass of each element on the reactant side must equal the total mass of each element on the product side. • Total charge of reactant side must equal the total charge of product side.

  23. Combustion Reactions • Reactions that occur between oxygen (O2) and another element in a compound. • When the other compound is a hydrocarbon, the products of complete combustion are carbon dioxide and water vapor. • Hydrocarbons are molecular compounds composed of only hydrogen and carbon and are a class of organic compounds.

  24. Practice Balance the following equations for the following combustions reactions. C3H8 + O2 CO2 + H2O C5H10 + O2 CO2 + H2O

  25. Stoichiometric Calculations • Calculating the mass of a product from the mass of a reactant requires: • A balanced chemical reaction • Molar mass of the reactant • Molar mass of the product

  26. Example How much carbon dioxide would be formed if 10.0 grams of C5H12 were completely burned in oxygen? C5H12 + 8 O2 ---> 5 CO2 + 6 H2O

  27. Percent Composition • Mass percent of an element in a compound Mass % = mass of element in compound x 100% mass of compound • Practice: Calculate the percent of iron in iron(III) oxide, (Fe2O3).

  28. Mass Percent and Empirical Formulas Problem . A sample of an iron compound is 22.0% Fe, 50.2% oxygen, and 27.8% chlorine by mass. What is the empirical formula of this compound. Answer: FeO8Cl2

  29. Simplified Carbon Cycle

  30. Problem Sodium carbonate (105.988 g/mol) reacts with hydrochloric acid (36.461) to produce sodium chloride, water, and carbon dioxide. How much hydrochloric acid is required to produce 10.0 g of carbon dioxide (44.01g/mol)?

  31. Combustion Analysis CaHb + excess O2 ---> a CO2(g) + b/2 H2O The percent of carbon and hydrogen in CaHb can be determined from the mass of H2O and CO2 produced.

  32. Percent Composition and Empirical Formulas • Assume there is 100 g of the sample, so the percent composition will equal the number of grams of each element. • Convert the grams of each element into the moles of each element with their molar mass. • Divide the smallest number of moles of an element into the moles of each element present. • Convert the fractional ratios for each element into whole numbers by multiplying all the ratios by the same number. • The resulting numbers are the subscripts for the each element in the empirical formula.

  33. Example Asbestos was used for years as an insulating material in buildings until prolonged exposure to asbestos was demonstrated to cause lung cancer. Asbestos is a mineral containing magnesium, silicon, oxygen, and hydrogen. One form of asbestos, chrysotile (520.27 g/mol), has the composition 28.03% magnesium, 21.60% silicon, 1.16% hydrogen. Determine the empirical formula of chrysotile.

  34. Mass Spectrometry and Molecular Mass • All Mass spectrometers separate atoms and molecules by first converting them into ions and then separating those ions based on the ratio of their masses to their electric charges. • Mass spectrometers are instruments used to determine the mass of substances.

  35. Mass Spectrometer

  36. Mass Spectra

  37. Determining the Molecular Formula • The molecular formula can be determined from the percent composition and mass spectral data.

  38. Example A combustion analysis of an unknown compound indicated that it is 92. 23% C and 7.82% H. The mass spectrum indicated the molar mass is 78 g/mol. What is the molecular formula of this unknown compound?

  39. Limiting Reactants During photosynthesis a reaction mixture of carbon dioxide and water is converted to a molecule of glucose.

  40. Ham and Cheese as “limiting” reagents…

  41. Limiting Reagents • The limiting reactant is completely consumed in the chemical reaction. • The amount of product formed depends on the amount of the limiting reagent available.

  42. Example 10.0 g of methane (CH4) is burned in 20.0 g of oxygen (O2) to produce carbon dioxide (CO2) and water (H2O). • What is the limiting reactant? • How many grams of water will be produced?