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Understanding Chemical Equations and Reactions

This text explains the fundamentals of chemical equations and reactions, including balancing equations, different types of reactions, and stoichiometry. Learn how to predict the products and quantities involved in chemical reactions.

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Understanding Chemical Equations and Reactions

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  1. Chemical equations describeCHEMICAL REACTIONS. • During a chemical reaction, the ways in which atoms are joined together are changed. • OLD bondsare broken andNEW bondsare formed asREACTANTS are converted into PRODUCTS. • A reactionSTARTS with substances known asREACTANTS (always written on theLEFTside of the equation). • The substances that areFORMED during a reaction are known asPRODUCTS (always written on theRIGHTside of the equation).

  2. Chemical Equations • Chemical Equations are used (as chem. shorthand) to represent what is occurring during a chemical reaction. • Ex.: Butane burns in oxygen to produce carbon dioxide gas and water vapor. • C4H10 + O2 --->CO2 + H2O ∆

  3. Why Balance Equations? • When 9.386 g Ca reacts completely with 7.514 g S, 16.90 g of CaS are formed. Ca(s) + S(s) ---> CaS(s)

  4. Balancing Equations • Make sure the chemical formulas are CORRECT(you cannot change the correct formula of a substance). • I2 Br2 Cl2 F2 O2 N2 H2 • I Bring Clay For Our New House • Use COEFFICIENTS (numbers in front of formulas) to change the number of atoms of an element. • Coefficients represent MULTIPLES of the formulas

  5. Balance the following • N2 (g) + H2 (g) ---> NH3 (g) • FeCl3 (s) ---> Fe(s) + Cl2 (g) • NaOCl + KI + HC2H3O2 ---> I2 + NaCl + KC2H3O2 + H2O

  6. Balance the following • C3H8 + O2 --> CO2 + H2O • C3H6 + O2 --> CO2 + H2O

  7. Combination Decomposition Single Replacement Double Replacement Neutralization Combustion RedOx A + B --> AB AB --> A + B A + BC --> AC + B AC + BD --> AD + BC A + O2 --> AO CxHy + O2 --> x CO2 + y/2 H2O Types of Chemical Reactions

  8. Measuring MATTER Technique - Necessary information Counting - Avogadro’s # Weighing - Molar Mass Measuring Volume - Molar Volume

  9. Types of Chemical Particles • Atoms - represented by the symbol of an element (C, Ag, N, Na, Cl, Fe, Pb, S, etc.) • Molecules - represented by multiple symbol(s) of NONMETAL atoms/elements (CO2, H2O, O2, CH4, etc.) • Ions - represented by the symbol/formula of an ion • Formula Units - represented by the formula of an IONIC compound (NaCl, AgNO3, Fe2O3, etc.) • 1 mole = 6.02 x 1023 particles (a conversion factor)

  10. Avogadro’s Number = 6.02 x 1023 • This number can be determined experimentally several ways: • Measurement of crystal structure • Ti - body-centered unit cell (2 atoms/unit cell) • # of atoms/mol = 6.02 x 1023atoms

  11. MOLAR MASS • The number of grams of a substance equivalent to the sum of all its average atomic mass units (amu) is known as the molar mass. • One mole of particles is equal to its molar mass in grams.

  12. MOLAR VOLUME • The volume, 22.4 L, of anygas at STP is known as the molar volume. • STP = Standard Temperature and Pressure • 273 K (O˚C, 32 ˚F) and 1 atm (101.3 kPa, 760 mm Hg, 29.92 in Hg)

  13. Molar Conversions

  14. Molar Conversions

  15. STOICHIOMETRY -think recipes! • Balanced chemical equations can be used to predict the QUANTITATIVE amounts ofREACTANTSandPRODUCTS. • N2(g)+ 3H2 (g) ---> 2NH3 (g) • Particles • Molecules • Atoms • Moles • Mass

  16. N2(g)+ 3H2 (g) ---> 2NH3 (g) • How many moles of hydrogen will react with 0.00326 mol N2? • How many molecules of ammonia are produced when 4.55 x 1018 molecules of hydrogen react? • How many atoms of hydrogen are in the ammonia produced? • How many grams of nitrogen are required to react with 75.8 g hydrogen?

  17. Stoichiometry • Use thestoichiometricmole ratio to convert from moles of one substance to moles of another substance within the reaction • C4H10 + O2 ---> CO2 + H2O • How many moles of water are produced when 0.48 moles of butane, C4H10, react? • How many molecules of butane are needed to produce 12.00 grams of H2O?

  18. Stoichiometry, cont. • In the lab, we determine the masses of different substances, rather than moles, and therefore must be able to convert from grams of one substance in a reaction to grams of another. • NaOCl + KI + HC2H3O2 ---> I2 + NaCl + KC2H3O2 + H2O • How many grams of iodine are produced when 0.35 g of potassium iodide react?

  19. I2 (aq) + Na2S2O3(aq) --> Na2S4O6(aq) + NaI • How many grams of sodium iodide are produced when 0.203 g of iodine react with excess Na2S2O3? • Which reactant limits the amount of product that can be made? Why?

  20. Limiting Reactants • The product(s) of a reaction is/are limited by how much of each reactant is present (available) in the reaction. • Two types of reactants • Limiting - this is the reactant you run out of first! • Excess - at the end of the reaction there will be some of this reactant left over (excess:-)).

  21. Combustion of Magnesium • Mg(s) + O2(g in air)--> MgO(s) • How much magnesium oxide can be produced when 1.085 g Mg burns in air? • Determine the limiting and excess reactants • Calculate the theoretical yield from the limiting reactant • Mg(s) + O2(g in air)--> MgO(s)

  22. Molar Mass 208.3g/mol 76.1 g/molBaCl2 + NH4SCN --> Ba(SCN) 2 + NH4Cl • 34.5 g BaCl2 react with 44.3 g NH4SCN. • How much NH4Cl can be produced (theoretical yield)? • STEP 1: Determine the LR and ER! (use mole ratio) • STEP 2: Determine the Theoretical Yield from LR

  23. Experimental Reaction Yield • Balanced equations can be used to calculate the amount of product that will form during a reaction - called the THEORETICAL YIELD • The amount of product that actually forms during a chemical reaction is called the ACTUAL YIELD • The actual yield is often less than the theoretical yield.

  24. Percent Yield • The percent yield is the ratio of the actual yield compared to the theoretical yield, converted to a percent. • actual yield • % YIELD = --------------------- x 100 • theoretical yield

  25. N2(g)+ H2(g)--->NH3 (g) • 0.075 g N2 react with 0.0095 g H2 to produce 0.051 g NH3. • Which reactant is the limiting reactant? Which reactant is in excess? • How much ammonia should be produced? • What is the percent yield for this reaction?

  26. Mole Ratios in Chemical Formulas • The Empirical (Simplest) Formula is a ratio of atoms in the compound (this is equivalent to the mole ratio of atoms). • Ex. If 3.10 g Fe reacts with chlorine to make 9.01 g of a compound, what is the simplest formula of the compound?

  27. Mole Ratios in Chemical Formulas • For hydrated compounds, the mole ratio of water to the compound is expressed in the formula. • Ex. If 2.00 g of a copper (II) sulfate hydrate is heated and the mass of the anhydrate is 1.28 g, what is the formula of the hydrate?

  28. Oxidation & ReductionRedOx • Reactions that involve the transfer of electrons between “particles” are known as RedOx rxns. • Oxidation is the loss of electron(s) from an atom. • Reduction is the gain of electron(s) by an atom. • OIL RIG • Examples: • Rusting; Batteries; Antiseptics; Combustion of Hydrocarbons; Reactions in Biochemical Pathways

  29. Chemistry 104Quiz #6 1. How many grams do 8.5 x 1025 molecules of water weigh?2. Balance the equation: C6H14 + O2 CO2 + H2O3. For the reaction: N2H4 + 2H2O2  N2 + 4H2O How many grams of dinitrogen tetrahydride are needed to form 20.0 g dihydrogen monoxide?4. If 18.0 g hydrogen peroxide react with the amount of dinitrogen tetrahydride determined in Q.#3 and produces 15.6 g water, what is the percent yield?

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