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Chemical Reaction and Equations

Chemical Reaction and Equations. Evidence of a Reaction Chemical Equations Balancing Chemical Equations. 1. List four observations that suggest that a chemical reaction has taken place. 2. List three requirements for a correctly written chemical equation.

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Chemical Reaction and Equations

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  1. Chemical Reaction and Equations Evidence of a Reaction Chemical Equations Balancing Chemical Equations

  2. 1. List four observations that suggest that a chemical reaction has taken place. 2. List three requirements for a correctly written chemical equation. 3. Write a word equation and a formula equation for a given chemical reaction. 4. Write a chemical equation by balancing a formula equation by inspection. 8.1 Writing and Balancing Chemical Reactions and Equations

  3. Evidence of a Chemical Reaction • The release of light or heat .

  4. More Light and Heat Evidence

  5. Evidence of a Chemical Reaction • The production of a gas (bubbles) .

  6. More Gas Production Evidence O2 O2 O2 2 H2 2H2 HCl H2 Zn H2O

  7. Evidence of a Chemical Reaction • The formation of a precipitate (solid) when two aqueous solutions are mixed together. Pb(NO3)2 KI PbI2

  8. When gases form a Precipitate NH4Cl(s)

  9. Evidence of a Chemical Reaction PbI2 • A color change is often evidence of a chemical reaction.

  10. Characteristics of Chemical Equations • The equation represents known facts. • The equation must contain correct formulas! Reactant + Reactant  Product + Product • The equation must obey The Law of Conservation of Mass. Mass of Reactants = Mass of Products

  11. Word Equations • A chemical equation in which the reactants and products are represented by words (names). Methane + Oxygen  Carbon dioxide + Water The  means yields or produces The + means “and”

  12. Formula Equation • Reactants and products are described by using their symbols and/or formulas. CH4 + O2 CO2 + H2O • A formula equation does not obey “The Law of Conservation of Mass.”

  13. Chemical Equations 1. Use symbols and formulas to represent the reactants and products. 2. Adjust the coefficients in front of each reactant or product to identify the number of each reactant or product. 1 CH4 + 2 O2 1 CO2 + 2 H2O

  14. Chemical Equations 3. Use special symbols in parenthesis to indicate the state of the reactants or products. (s) = solid or (cr) Crystal (l) = liquid (g) = gas (aq) = aqueous (dissolved in water)

  15. Chemical Equations 4. Use symbols above the yield arrow to represent catalysts or special conditions. heat Δ atm pressure MnO2

  16. Reversible Reactions and Equations • Some chemical reactions occur in both directions. N2 (g) + 3 H2 (g)2 NH3 (g)

  17. Balanced Chemical Equations The coefficients represent the relative amounts of reactants and products. 1 CH4 + 2 O2 1 CO2 + 2 H2O 1 molecule 2 molecules 1 molecule 2 molecules 1 mole 2 moles 1 mole 2 moles

  18. The Importance of a Balanced Chemical Equation cont. • Coefficients can be used to determine the relative masses of reactants and products 1 CH4 + 2 O2 1 CO2 + 2 H2O 1 mole 2 moles 1 mole 2 moles 1(16.05g) + 2(32.00g) = 1(44.01g) + 2(18.02g) 80.05 g = 80.05 g Law of Conservation of Mass 

  19. A Chemical ReactionEvidence = ? Reactants CH4 2 O2

  20. Light from a Big Bang! Reactants Products CH4 CO2 2 H2O 2 O2 ENERGY

  21. Balancing Chemical Equations 1.Write the Word (Name) Equation. Methane + Oxygen  Carbon dioxide + Water 2. Write the Formula Equation: CH4 + O2 CO2 + H2O

  22. Balancing Chemical Equations 3. Change the Coefficients as needed. 4. NEVER change a subscript ! 5. Balance different types of atoms one at a time.

  23. Balancing Chemical Equations 6. Balance elements that only appear once on each side of the equation first. 7. Balance polyatomic ions that appear on both sides of the equation as one unit. 8. Balance “H” and “O” last.

  24. 9. Elements • Use the symbol from the periodic table for an element. Na Au S C Sn P • The following elements always appear as diatomic molecules in chemical equations. H2 N2 O2 F2 Cl2 Br2 I2

  25. Checking Your Work • Count atoms to be sure that the equation is balanced. • Counting = Coefficient x subscript

  26. Example One • Sodium metal (solid) combines with chlorine gas to produce solid sodium chloride.

  27. Example One • Sodium metal (solid) combines with chlorine gas to produce solid sodium chloride. Na (s) + Cl2 (g) NaCl (s) • Remember that chlorine is diatomic!

  28. Example One • Sodium metal (Solid) combines with chlorine gas to produce solid sodium chloride. Na (s) + Cl2 (g) NaCl (s) 2 Na (s) + Cl2 (g) 2 NaCl (s)

  29. Example Two • When copper metal (solid) reacts with aqueous silver nitrate, the products formed are aqueous copper (II) nitrate and silver metal (solid).

  30. Example Two • When copper metal (solid) reacts with aqueous silver nitrate, the products formed are aqueous copper (II) nitrate and silver metal (solid). Cu (s) + AgNO3 (aq) Cu(NO3)2 (aq) + Ag(s)

  31. Example Two • When copper metal (solid) reacts with aqueous silver nitrate, the products formed are aqueous copper (II) nitrate and silver metal (solid). Cu (s) + AgNO3 (aq) Cu(NO3)2 (aq) + Ag(s) Cu (s) + 2 AgNO3 (aq) Cu(NO3)2 (aq) + 2 Ag(s)

  32. Example Three • Solid Iron (III) oxide and carbon monoxide gas react to form solid iron and carbon dioxide gas.

  33. Example Three • Solid Iron (III) oxide and carbon monoxide gas react to form solid iron and carbon dioxide gas. Fe2O3 (s) + CO(g) Fe(s) + CO2(g)

  34. Example Three • Solid Iron (III) oxide and carbon monoxide gas react to form solid iron and carbon dioxide gas. Fe2O3 (s) + CO(g) Fe(s) + CO2(g) Fe2O3 (s) + 3 CO(g) 2 Fe(s) + 3 CO2(g)

  35. THE END

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