1. Chapter 3 Chemical Reactions

2. 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. 3 States of Matter Changes from one state to another: Physical Change heating cooling

4. 4 Physical Change vs. Chemical Change

5. 5

6. 6 Chemical Equations Symbolic representation of a chemical reaction (chemical change) that shows: 1.-reactants on left side of reaction 2.-products on right side of equation 3.-relative amounts of each using coefficients H 2 + O 2 H 2 O for a reaction to occur molecules, atoms, ions must interact with one another in the appropriate orientation under the right conditions

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

8. 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. 9 Chemical Equations

10. 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. 11 Balancing Composition Reactions Na(s) + Cl 2 (g)  NaCl(s) Mg(s) + O 2 (g)  MgO(s) Al(s) + Br 2 (l)  AlBr 3 (s)

12. 12 Balancing Reactions On Your Own P 4 (s) + O 2 (g)  P 4 O 10 (s) CO(g) + O 2 (g)  CO 2 (g) P 4 (s) + Cl 2 (g)  PCl 3 (l) SO 2 (g) + O 2 (g)  SO 3 (g) P 4 O 6 (g) + O 2 (g)  P 4 O 10 (s)

13. 13 Balancing Decomposition Reactions N 2 O(g)  N 2 (g) + O 2 (g) H 2 O 2 (aq)  H 2 O(l) + O 2 (g) AgBr(s)  Ag(s) + Br 2 (l) NH 4 HCO 3 (s)  NH 3 (g) + H 2 O(g) + CO 2 (g)

14. 14 Balancing Displacement Reactions on Your Own AgNO 3 (aq) + Cu(s)  CuNO 3 (aq) + Ag(s) Al(s) + H 2 SO 4 (aq)  Al 2 (SO 4 ) 3 (aq) + H 2 (g) Cl 2 (g) + NaI(aq)  I 2 (s) + NaCl(aq) CaCl 2 (aq) + Na 3 PO 4 (aq)  NaCl(aq) + Ca 3 (PO 4 ) 2 (s) Ca(OH) 2 (aq) + HNO 3 (aq)  Ca(NO 3 ) 2 (aq) + H 2 O(l) Ca(NO 3 ) 2 (aq) + K 2 CO 3 (aq)  KNO 3 (aq) + CaCO 3 (s)

15. 15 Law of Conservation of Matter Combustion reaction: the burning of a fuel in oxygen producing oxides or oxygen containing compounds –-NH 3 burns in oxygen to form nitrogen monoxide and water

16. 16 Law of Conservation of Matter C 7 H 16 burns in oxygen to form carbon dioxide and water.

17. 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. 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. 19 Strong Electrolytes conduct electricity extremely well in dilute aqueous solutions –-ionize in water 100% Examples: 1.HCl, HNO 3, etc strong soluble acids 2.NaOH, KOH, etc strong soluble bases 3.NaCl, KBr, etc soluble ionic salts

20. 20 Strong Ionic Salts

21. 21 Weak Electrolytes conduct electricity poorly in aqueous solutions -ionize much less than 100% in water Examples: 1.CH 3 COOH, (COOH) 2 weak acids 2.NH 3, Fe(OH) 3 weak bases

22. 22 Properties of Aqueous Solutions Nonelectrolytes solutes that do not conduct electricity in water – do not “ionize” Examples: C 2 H 5 OH – ethanol Sugars – glucose, sucrose, etc.

23. 23 Aqueous Solution Conductivity

24. 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. 25 Solubility Rules for determining solubility: soluble (dissolves) vs. insoluble (does not dissolve) Figure 5.3 on page 179 OH - and O 2-, except Ba 2+

26. 26 SolubleInsolubleExceptions 1. Group IA and ammonium salts (Li +, Na +, K +, NH 4 + ) ___________ 2. Acetates, nitrates, chlorates, perchlorates (CH 3 COO -, NO 3 -, ClO 3 -, ClO 4 - ) ___________ 3. most chlorides, bromides, and iodides (Cl -, Br -, I - ) Salts formed with Ag +, Hg 2+, Pb 2+ 4. most fluorides (F - )Salts formed with Group IIA 5. most sulfates (SO 4 2- )Salts formed with Group IIA (Ca 2+, Sr 2+, Ba 2+ ), Ag +, Hg 2+, Pb 2+ 6. most carbonates, phosphates, sulfides (CO 3 2-, PO 4 3-, S 2- ) Salts formed with Group IA and NH 4 + (rule #1) 7. most oxides (O 2- )_______________________ 8. most hydroxides (OH - )Salts formed with Group IA and Ca 2+, Sr 2+

27. 27 Solubility

28. 28 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: 1.form H 2 O – neutralization (acid-base reactions) 2.form an insoluble solid (precipitation reactions) 3.form a gas -Ion removal is the driving force of metathesis reactions

29. 29 Precipitation Reactions Three representation: 1.1. Molecular equation 2.2. Total ionic equation Ag + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq)  AgCl(s) + Na + (aq) + NO 3 - (aq) 3. Net ionic equation Ag + (aq) + Cl - (aq)  AgCl(s)

30. 30 Precipitation Reactions 1. Molecular equation 2.Total ionic reaction 3. Net ionic reaction

31. 31 Arrhenius Acids substances that generate H 3 O + (H + ) in aqueous solutions -Strong acids ionize 100% in water (l)(l)

32. 32 Substances that donate protons (H + ) Strong Acids FormulaName 1.HClhydrochloric acid 2.HBrhydrobromic acid 3.HIhydroiodic acid 4.HNO 3 nitric acid 5.H 2 SO 4 sulfuric acid 6.HClO 3 chloric acid 7.HClO 4 perchloric acid Bronsted-Lowry Acids

33. 33 Acids -Weak acids ionize <100% in water

34. 34 Common Weak Acids FormulaName 1.HF hydrofluoric acid 2.CH 3 COOHacetic acid (vinegar) 3.HCNhydrocyanic acid 4.HNO 2 nitrous acid 5.H 2 CO 3 carbonic acid (soda water) 6.H 3 PO 4 phosphoric acid Acids

35. 35 Substance that produce OH- ions in aqueous solution (water) –Strong bases ionize 100% in water Arrhenius Bases Weak bases are covalent compounds that ionize <100% in water (l)(l) CC

36. 36 Substances that accept protons (H + ) Strong bases: 1.LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) 2, Sr(OH) 2 2.Notice that they are all hydroxides of IA and IIA metals Bronsted-Lowry Bases

37. 37 Acid-Base (neutralization) Reactions form water and salt (ionic compound) –acid + base  salt + water 1. Molecular equation 2. Total ionic equation 3. Net ionic equation l (l)(l)

38. 38 Acid-Base (neutralization) Reactions 1. Molecular equation 2. Total ionic equation 3. Net ionic equation (l)(l) (l)(l)

39. 39 There are four acid-base reaction combinations that are possible: 1.strong acids – strong bases 2.weak acids – strong bases 3.strong acids – weak bases 4.weak acids – weak bases Acids and Bases

40. 40 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 H 2 SO 4(aq) + NaOH (aq)  NaHSO 4(aq) + H 2 O ( l ) 1 mol sulfuric acid reacts with 2 mols sodium hydroxide H 2 SO 4(aq) + 2NaOH (aq)  Na 2 SO 4(aq) + 2H 2 O ( l ) Acids and Bases

41. 41 Gas Forming Reactions H 2 CO 3  H 2 O(l) + CO 2 (g) H 2 SO 3  H 2 O(l) + SO 2 (g) NH 4 OH  NH 3 (g) + H 2 O(l)