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Chapter 3

Chapter 3. Chemical Reactions. Chemical and Physical Properties. Chemical Changes rusting or oxidation chemical reactions Physical Changes changes of state density, color, solubility, melting, boiling Extensive Properties: depend on quantity

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Chapter 3

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  1. Chapter 3 Chemical Reactions

  2. Chemical and Physical Properties • Chemical Changes • rusting or oxidation • chemical reactions • Physical Changes • changes of state • density, color, solubility, melting, boiling • Extensive Properties: depend on quantity • Intensive Properties: do not depend on quantity

  3. States of Matter • Changes from one state to another: Physical Change • heating • cooling

  4. Physical Change vs. Chemical Change

  5. Physical Change vs. Chemical Change

  6. Chemical Equations for a reaction to occur molecules, atoms, ions must interact with one another in the appropriate orientation under the right conditions Symbolic representation of a chemical reaction (chemical change) that shows: • -reactants on left side of reaction • -products on right side of equation • -relative amounts of each using coefficients H2 + O2 H2O

  7. Chemical Equations • Are an attempt to show on paper what is happening at the molecular level

  8. Chemical Equations • Look at the information an equation provides: • reactants products 1 formula unit 3 molecules 2 atoms 3 moles (molecule/mole) (moles/f.u.) (moles/f.u.) (molecules.f.u.) the states of matter also listed

  9. Chemical Equations

  10. Chemical Equations • Law of Conservation of Matter • Matter is neither created nor destroyed in a chemical reaction • -There is no detectable change in quantity of matter in an ordinary chemical reaction • -Balanced chemical equations must always include the same number of each kind of atom on both sides of the equation Balancing equations is a skill acquired only with a lot of practice!!! • By working many problems

  11. Balancing Composition Reactions Na(s) + Cl2(g)  NaCl(s) Mg(s) + O2(g)  MgO(s) Al(s) + Br2(l)  AlBr3(s)

  12. Balancing Reactions On Your Own P4(s) + O2(g)  P4O10(s) CO(g) + O2(g)  CO2(g) P4(s) + Cl2(g)  PCl3(l) SO2(g) + O2(g)  SO3(g) P4O6(g) + O2(g)  P4O10(s)

  13. Balancing Decomposition Reactions N2O(g)  N2(g) + O2(g) H2O2(aq)  H2O(l) + O2(g) AgBr(s)  Ag(s) + Br2(l) NH4HCO3(s)  NH3(g) + H2O(g) + CO2(g)

  14. Balancing Displacement Reactions on Your Own AgNO3(aq) + Cu(s)  CuNO3(aq) + Ag(s) Al(s) + H2SO4(aq)  Al2(SO4)3(aq) + H2(g) Cl2(g) + NaI(aq)  I2(s) + NaCl(aq) CaCl2(aq) + Na3PO4(aq)  NaCl(aq) + Ca3(PO4)2(s) Ca(OH)2(aq) + HNO3(aq)  Ca(NO3)2(aq) + H2O(l) Ca(NO3)2(aq) + K2CO3(aq)  KNO3(aq) + CaCO3(s)

  15. Law of Conservation of Matter Combustion reaction: the burning of a fuel in oxygen producing oxides or oxygen containing compounds • -NH3 burns in oxygen to form nitrogen monoxide and water

  16. Law of Conservation of Matter • C7H16 burns in oxygen to form carbon dioxide and water.

  17. Solutions a mixture of two or more substances dissolved in another Solute: substance present in the smaller amount that is dissolved by the solvent Solvent: substance present in the larger amount that dissolves the solute

  18. Properties of Aqueous Solutions • Electrolytes • produce ions in solution and conduct electricity • Strong electrolytes • ionize or dissociate 100% in water • NaCl(s)Na+(aq) + Cl-(aq) • Weak electrolytes • ionize or dissociate much less than 100% in water • HF(l) H+(aq) + F-(aq)

  19. Strong Electrolytes conduct electricity extremely well in dilute aqueous solutions • -ionize in water 100% Examples: • HCl, HNO3, etc • strong soluble acids • NaOH, KOH, etc • strong soluble bases • NaCl, KBr, etc • soluble ionic salts

  20. Strong Ionic Salts

  21. Weak Electrolytes conduct electricity poorly in aqueous solutions -ionize much less than 100% in water Examples: • CH3COOH, (COOH)2 • weak acids • NH3, Fe(OH)3 • weak bases

  22. Properties of Aqueous Solutions Nonelectrolytes solutes that do not conduct electricity in water – do not “ionize” • Examples: • C2H5OH – ethanol • Sugars – glucose, sucrose, etc.

  23. Aqueous Solution Conductivity

  24. Solubility • maximum amount of solute that can dissolve in a given amount of solvent • -defined as the amount of solute that dissolves in 100 g solvent • Unsaturated Solution: • contains less than the maximum amount that dissolves • Saturated solution: • contains the maximum amount that dissolves • Supersaturated solution: • contains more than the maximum amount that normally dissolves

  25. Solubility Rules for determining solubility: • soluble (dissolves) vs. insoluble (does not dissolve) OH- and O2-, except Ba2+ Figure 5.3 on page 179

  26. Solubility

  27. Metathesis Reactions two ionic aqueous solutions are mixed and the ions switch partners AX + BY  AY + BX Metathesis reactions remove ions from solution in 3 ways: • form H2O – neutralization (acid-base reactions) • form an insoluble solid (precipitation reactions) • form a gas • -Ion removal is the driving force of metathesis reactions

  28. Three representation: 1. Molecular equation 2. Total ionic equation Precipitation Reactions Ag+(aq) + NO3-(aq) + Na+ (aq) + Cl-(aq)  AgCl(s) + Na+ (aq) + NO3-(aq) 3. Net ionic equation Ag+(aq) + Cl-(aq)  AgCl(s)

  29. Precipitation Reactions • 1. Molecular equation • 2.Total ionic reaction • 3. Net ionic reaction

  30. Arrhenius Acids substances that generate H3O+ (H+) in aqueous solutions -Strong acids ionize 100% in water (l)

  31. Bronsted-Lowry Acids Substances that donate protons (H+) • Strong Acids • FormulaName • HCl hydrochloric acid • HBr hydrobromic acid • HI hydroiodic acid • HNO3 nitric acid • H2SO4 sulfuric acid • HClO3 chloric acid • HClO4 perchloric acid

  32. Acids • -Weak acids ionize <100% in water

  33. Acids • Common Weak Acids • FormulaName • HF hydrofluoric acid • CH3COOH acetic acid (vinegar) • HCN hydrocyanic acid • HNO2 nitrous acid • H2CO3 carbonic acid (soda water) • H3PO4 phosphoric acid

  34. Arrhenius Bases • Substance that produce OH- ions in aqueous solution (water) • Strong bases ionize 100% in water C C • Weak bases are covalent compounds that ionize <100% in water (l)

  35. Bronsted-Lowry Bases Substances that accept protons (H+) • Strong bases: • LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2,Sr(OH)2 • Notice that they are all hydroxides of IA and IIA metals

  36. Acid-Base (neutralization) Reactions form water and salt (ionic compound) • acid + base  salt + water • 1. Molecular equation l • 2. Total ionic equation (l) • 3. Net ionic equation

  37. Acid-Base (neutralization) Reactions 1. Molecular equation (l) 2. Total ionic equation (l) 3. Net ionic equation

  38. Acids and Bases There are four acid-base reaction combinations that are possible: • strong acids – strong bases • weak acids – strong bases • strong acids – weak bases • weak acids – weak bases

  39. Acids and Bases • Polyprotic acids: • Have more than 1 hydrogen ion that it can donate to a base 1 mol sulfuric acid reacts with 1 mol sodium hydroxide H2SO4(aq) + NaOH(aq) NaHSO4(aq) + H2O(l) 1 mol sulfuric acid reacts with 2 mols sodium hydroxide H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H2O(l)

  40. Gas Forming Reactions H2CO3 H2O(l) + CO2 (g) H2SO3  H2O(l) + SO2 (g) NH4OH  NH3(g) + H2O(l)

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