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Stoichiometry

Stoichiometry. College Chemistry Chapter 3. Law of Conservation of Mass. All chemical and physical reactions must follow the LCM. Lavoisier was the first to state this law and he was also given credit as the founder of quantitative chemistry.

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Stoichiometry

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  1. Stoichiometry College Chemistry Chapter 3

  2. Law of Conservation of Mass • All chemical and physical reactions must follow the LCM. • Lavoisier was the first to state this law and he was also given credit as the founder of quantitative chemistry. • All equations must be balanced in order to adhere to the LCM.

  3. Law of Conservation of Mass • Regular reactions: atoms and masses balance • Redox reactions: atoms, masses, and charges balance • Here there will be change in the oxidation state of ions during the reaction. • One element will be oxidized; that means that it will lose electrons and become more positive. • One element will be reduced; that means that it will gain electrons and become more negative. • A balanced equation may require that half reactions be considered in order to balance the charges.

  4. Redox examples • Mg0(s) + 2HCl(aq) MgCl2(aq) + H2(g) Half reactions: Oxidation: Mg0(s) Mg+2 + 2e- Reduction: 2H+(aq) + 2e-  H20 Spectator ions: Cl-(aq)

  5. Types of Reactions • Synthesis (direct combination) – A+BC • May use elements • May use compounds • Using metal oxides: get hydroxides • Using nonmetal oxides: get acids • Decomposition (analysis) – C  A+B • Ordinary binary compounds • Chlorates – chloride salt and oxygen • Carbonates – oxides and carbon dioxide • Hydroxides – metal oxide and water • Acids – nonmetallic oxide and water

  6. Examples • Ordinary binary compounds • 2NaCl(s) 2Na0(s) + Cl20(g) • Chlorates – chloride salt and oxygen • 2KClO3(s)  2KCl(s) + 3O2(g) • Carbonates – oxides and carbon dioxide • BaCO3(s) BaO(s) + CO2(g) • Hydroxides – metal oxide and water • NaOH(l) Na2O(s) + H2O(l) • Acids – nonmetallic oxide and water • H2SO4(s) SO3(g) + H2O(l)

  7. Types of Reactions • Single displacement – A + BC  AC + B • Must use activity series to determine if reaction happens • May be used with halogens with F2 > Cl2 > Br2 > I2 • Double replacement – AB + CD  AD + CB • Must consider the driving forces to determine if reaction happens • Formation of a precipitate • Formation of a gas • Formation of water • Formation of a small molecular compound

  8. Solubility Rules • Formation of a precipitate occurs if the product is insoluble in aqueous solution. • Soluble compounds are those containing • NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1 • NO3- • C2H3O2- except with Ag+ • Cl-, Br-, I- except with Ag+, Hg2+2, and Pb+2 • SO4-2 except with Sr+2, Ba+2, Hg2+2, Pb+2

  9. Solubility Rules • Insoluble compounds are those containing • S-2 except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1, Sr+2, Ba+2, Ca+2 • CO3-2 except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1 • PO4-3 except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1 • OH- except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1, Sr+2, Ba+2, Ca+2

  10. REMEMBER • Diatomics such as H2, O2, N2, F2, Cl2, Br2, or I2 • The charge on a compound must net zero. • The charges on the representative elements are predictable; those on the transition elements are not except for zinc, silver, and cadmium.

  11. Types of Reactions • Combustion – one reactant will be oxygen and the other one may be an organic compound, a metallic compound, or a nonmetallic compound • Complete – occurs when there is excess oxygen  CO2(g) + H2O(l) • Incomplete – occurs when there is a limited amount of oxygen  CO(g) + H2O(l)

  12. Types of Solution Reactions • Precipitation reactions – a precipitate forms as a result of two solutions reacting in solution • Acid-base reactions (neutralizations) – water and a salt form as a result of an acid and a base reacting in solution (may be a titration) • Oxidation-reduction reactions (redox) – products form such that the charges on reactant ions change as the reaction proceeds. • Two half reactions are sometimes required to determine the balanced equation. • May be in acidic or basic solution.

  13. Atomic and Molecular Weights • Atomic masses are based on the standard of carbon-12. • The masses on the periodic table are not the mass of any isotopes; they are the weighted average of all the isotopes.

  14. Empirical Formulas • Directly • Masses • Percent composition • Combustion analysis

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