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Chemical Reactions

Chemical Reactions. Potassium iodide (aq) reacts with lead nitrate (aq) producing a yellow precipitate of lead iodide. Empirical formula Molecular formula Structural formula. Chemical Formulas. Empirical. Formulas?. Molecular. Metal and non-metal

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Chemical Reactions

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  1. Chemical Reactions Potassium iodide (aq) reacts with lead nitrate (aq) producing a yellow precipitate of lead iodide

  2. Empirical formula Molecular formula Structural formula Chemical Formulas

  3. Empirical Formulas? Molecular • Metal and non-metal • Ionic - lacking discrete unit, or molecule • Simplest whole number ratio • Covalent compounds • Molecular and empirical formulas can be different • Glucose: molecular C6H12O6 versus empirical CH2O.

  4. Formula Weight • General term ; Molecular Weight used more often • “Sum” of the atomic weights of all the atoms in a chemical formula

  5. Hydrocarbons and Carbohydrates(Organic Chemistry) • Hydrocarbons • Composed of H and C • Some simple ; some complex • Examples: C3H8 (propane) C4H10 (butane) • Complete “combustion” yields: • CO2, H2O + energy • Carbohydrates • Composed of H, C, and O • Sugars, starches, cellulose • Examples: C12H22O11 (sugar) • Complete “oxidation” yields: • CO2, H2O + energy.

  6. Chemical Reactions • Occur through formation and breaking of chemical bonds between atoms • Involve changes in matter, creation of new materials, and energy exchange • Chemical equations • Concise representation of chemical reactions.

  7. Chemical Equations • Reactants - substances existing before reaction • Products - substances existing after reaction • Chemical symbols and formulas needed for quantitative purposes.

  8. Balancing Equations • Law of conservation of mass: atoms are neither created nor destroyed in chemical reactions • Mass of reactants = mass of products (i.e. balanced) • To balance a chemical equation • Change coefficients in front of chemical formulas • Do not change the subscripts (numbers within formulas).

  9. Fig 10.5 Subscripts vs Coefficients

  10. Example: Fig. 10.6

  11. Stepwise balancing procedurepage 279 Like an Inventory or “Bean Counting” • Law of conservation of mass (atoms are conserved) • Don’t change subscripts of formulas (compounds) • Multiply everything within a compound by the Coefficient • Look for the most complex reactants and products • Try to balance atoms within them first • Treat “Polyatomic” ions that appear on both sides as independent units with a charge • Cross-over technique and use of fractional coefficients top find least common multiple to balance the equation • See the next few examples: 10.5, 10.6, and 10.7

  12. Chemical Reactions Potassium iodide (aq) reacts with lead nitrate (aq) producing a yellow precipitate of lead iodide

  13. Bal Eq Classifications of Chemical Reactions • Combination reactions • Decomposition reactions • Replacement reactions • (1-3 = redox reaction subclasses) • Ion exchange reactions

  14. 4 Fe(s) + 3 O2(g) 2 Fe2O3(s) Combination Reactions Rust: Fig 10.10 • Two or more substances combine to form a single compound

  15. Decomposition Reactions Δ 2 HgO (s) 2 Hg (s) + O2 • Breakdown into simpler compounds or elements • Usually require some form of energy for Rx to occur

  16. 2 Al (s) + 3 CuCl2 (aq) 2 AlCl3 (aq) + 3 Cu (s) Example: Replacement Reaction Fig 10.13

  17. 2 Al (s) + 3 CuCl2 (aq) 2 AlCl3 (aq) + 3 Cu (s) Ag (s) + CuCl2 (aq) No Rx Replacement ReactionFig 10.12 • Occur because some elements have a stronger electron-holding ability • More active metals (Li, K, Ca, Na) give up electrons to elements lower on the list

  18. Ion Exchange Reaction AX + BY AY + BX 3 Ca(OH)2 (aq) + Al2(SO4)3 (aq) 3 CaSO4 (aq) + 2 Al(OH)3 • Ion Exchange: • ions of one compound interact with ions of another compound • Possible results: • Solid precipitates: ↓ • Gas forms: ↑ • Water formed: H2O (l) • No ion exchange reaction occurred if both products are soluble (See appendix B) • “ S ” versus “ i ”

  19. Information from Chemical Equations • Atoms are conserved • Mass is conserved • Law of combining volumes (gases) • Gases at the same temperature and pressure contain equal numbers of molecules

  20. Atomic mass unit (u) = 1/12th mass of carbon-12 One mole of a substance contains Avogadro’s number (6.02x1023) of the basic chemical unit of that substance (atoms, molecules, ions, …) Example: A mole of carbon-12 atoms is defined as having 6.02 x 1023 atoms totaling a mass of 12.00g Units of Measurement used with Equations

  21. Molar Weights • Gram-atomic weight: mass in grams equal to atomic weight • Gram-formula weight: mass in grams equal to formula weight • Gram-molecular weight: mass in grams equal to molecular weight

  22. Quantitative use of Equations

  23. Next Time: Water and Solutions

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