Chemical Equations and Reactions in Aqueous Solutions

1. Chemical Equations and Reactions in Aqueous Solutions Chapter 3

2. Chemical Reaction A chemical reaction: is a process in which one set of substances called reactants is converted to a new set of substances called products. reactantsproducts 3.1

3. Types of Chemical Equations • Word : hydrogen + oxygen → water • Skeleton: H2 (g) + O2(g) → H2O(l) • Balanced: 2H2(g) + O2(g) → 2H2O(l) • Molecular Equation • Ionic Equation • Net ionic Equation 3.1

4. Physical State Representation • (g) = ________________ • (l) = ________________ • (s) = ________________ • (aq) = ________________ These representations will commonly be observed in chemical equations.

5. Chemical Equations • There is no loss in the quantity of matter from a chemical reaction. This is the _____ __ __________ __ _____. • A balanced chemical equation must have the _____ _______ of each kind of atom on both sides of the equation. • Smallest possible whole-number coefficients.

6. Indicators for a Chemical Reaction • A color change • Formation of a solid (precipitate) within a clear solution • Evidence of a gas formation • Evolution or absorption of heat and/or light 3.1

7. Balancing Chemical Equations • By inspection. • NH3 + O2 → NO + H2O • C2H6 + O2 → CO2 + H2O • Fe2O3 + CO → Fe + CO2 • Using oxidation numbers (later) 3.1

8. Balancing Chemical Equations • __Al(s) + __Br2(l)  ___Al2Br6(s) (covalent) • __C3H8(g) + __O2(g)  __CO2(g) + __H2O(g) • __C4H10(l) + __O2(g)  __CO2(g) + __H2O(g) • __P4(s) + __Cl2(g)  __PCl3(s) Be sure to balance equations before proceeding with calculations.

9. Meaning of Chemical Equation • Interpret the following equation: 2NO(g) + O2(g) → 2NO2(g) 3.1

10. Balancing a Chemical Equation Balancing by Inspection: adjust stoichiometric coefficients by trial and error until a balanced condition is found. Useful strategies: 1. If an element occurs in only one compound on each side of the equation, try balancing this element first. 2. When one of the reactants or products exists as the free element, balance this element last. 3. In some reactions, certain groups of atoms (e.g. polyatomic ions) remain unchanged. In such case balance these groups as a unit.

11. 4.It is permissible to use fractional as well as integral numbers as coefficients. At times, an equation can be balanced most easily by using one or more fractional coefficients and then, if desired, clearing the fractions by multiplying all coefficients by a common multiplier. C2H6 + O2 CO2 + H2O  E.g. Balance the following chemical equations (a) H2 + O2 H2O (b) ZnS+ O2 ZnO +SO2 (c) Pb(NO3)2 + KI  PbI2 + KNO3 (d)Cu + AgNO3 Cu(NO3)2 + Ag

12. Practice Problems • Pages 86, 89, Examples 3.1, 3.2 • Page 119; questions 3.7, 3.8 • In the future: check the textbook (Chapter 3) for relevant problems and reading 3.1

13. Properties of Aqueous Solutions • Aqueous reactions account for virtually all chemistry that takes place in living systems • Solutions: a homogenous mixture • Solute ______________ • Solvent _____________ • Aqueous solution: (aq) water is the solvent • Nonaqueous solution: _______________ Many reactions that involve ionic compounds occur in water — aqueous solutions. 3.2

14. An Ionic Compound, CuCl2, in Water 3.2

15. Formation of ions in Aqueous Solution Dissociation: process by which an ionic substance breaks into ions when dissolved in water NaCl(s) → Na+ (aq) + Cl- (aq) Ionization: molecular substance reacts with water and forms ions. What is that? HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq) 3.2

16. Aqueous Solutions How do we know ions are present in aqueous solutions? The solutions conduct electricity! They are called ELECTROLYTES Strong Electrolytes: 100% separation into ions Some acids Some bases All the salts, when soluble NaCl(s) Na+ (aq) + Cl- (aq) 3.2

17. Aqueous Solutions Weak electrolytes: all other ionic compounds Acetic acid ionizes only to a small extent, so it is a weak electrolyte. CH3CO2H(aq) CH3CO2-(aq) + H+(aq) State of EQUILIBIUM is established 3.2

18. Aqueous Solutions Some compounds dissolve in water but do not conduct electricity. They are called nonelectrolytes. C12H22O11(s) → C12H22O11(aq) Examples include: sugar ethanol ethylene glycol 3.2

19. Ca(OH)2, Ba(OH)2 and Sr(OH)2 Group IIA, heavy metals) Know the strong acids & bases! 3.2

20. Classify each of the following as a strong, weak, or nonelectrolyte a. HClO4 _________________ b. LiOH ______________ c. C6H12 _____________________ • NH3 ____________________ CaCl2 __________________________ • HC2H3O2 ____________________ 3.2

21. Additional Definitions • PHYSICAL EQUILIBRIUM: a system in which two opposite physical changes occur at the same rate • Solution equilibrium • Vapor pressure equilibrium • CHEMICAL EQUILIBRIUM: • A system in which the forward and reverse reaction are equal • REVERSIBLE REACTION: (↔), the reaction proceeds in both directions. Usually for weak electrolytes. • CH3COOH(aq) + H2O(l) CH3COO-(aq) + H30+(aq) • REACTION THAT GO TO COMPLETION: (→) • all reactants form products • not reversible

22. Properties of Aqueous Solutions Assignment: Page 119, question 3.13 3.2

23. Types of Chemical Reactions: 3.3- 3.6 Assign oxidation numbers to all atoms in the following compounds: • Ni(ClO3)2 • K2Cr2O7 • Cu[NH3]4+2

24. Types of Chemical Reactions • Reaction involving no changes in oxidation states dinitrogen trioxide gas was bubbled into flask full of water • Reactions involving changesin oxidation states (Redox Reactions) magnesium metal was immersed into a solution of hydrochloric acid

25. Types of Chemical Reactions • Combination (Synthesis) reaction A + B  AB • Decomposition reactions AB  A + B • Displacement reactions AB + C  AC + B • Metathetical (change of position) reactions (double-replacement reactions) AB + CD  AD + CB • Combustion reactions reactions with oxygen CxHy + nO2  xCO2 + (y/2) H2O

26. Combination Reactions (Synthesis): A + B → C • Redox or non Redox? • Metals + Oxygen: • Lithium + oxygen → • Magnesium + oxygen → • Gold + oxygen → • Platinum + oxygen → Remember the diatomics Metals with multiple charges: choose the one with higher charge; Cu+2 and not Cu+1 Rule:

27. Combination Reactions: A + B → C • Nonmetals + Oxygen (Redox?) • Excess carbon with oxygen → • Limited amount of carbon with excess of oxygen → • Phosphorus + excess oxygen → • Phosphorus with limited amount of oxygen → • Rule:

28. Combination Reactions: A + B → C • Metals + nonmetals (Redox?) • Cesium metal + iodine → • Zinc + sulfur → • Magnesium + nitrogen → • Rule:

29. Synthesis Reactions Metal Oxides (most are solid) + Water: (Redox?) Magnesium oxide + water → Lithium oxide + water → Aluminum oxide + water → Iron(III) oxide + water → Rule:

30. Combination Reactions: A + B → C Nonmetal Oxides + Water : (Redox?) solid calcium oxide + water → solid lithium oxide + water → Can be Redox: 2NO2(g) + H2O (l) → HNO3 (aq) + HNO2(aq) • Rule:

31. Combination Reactions: A + B → C Metal Oxides + Nonmetal Oxides (Redox?) calcium oxide + silicon dioxide → lithium oxide + tetra phosphorus deca oxide → Rule: Notes: the more electropositive (most metallic) element is always written first P4O10; CaO; H2O, CO2 Check Periodic table

32. Decomposition Reactions • C  A + B reverse from combination (synthesis) • Metallic oxides  metal + oxygen • Nonmetallic oxides  nonmetal + oxygen • Hydroxide  metal oxide + water • Acid  nonmetallic oxide + water Which are Redox and which are not?

33. Decomposition (Special Cases) • Metal carbonates metallic oxide + CO2 • Metal bicarbonates: metal oxide + CO2(g) + H2O (l) • Metal sulfite  metallic oxide + SO2 • Metal chlorate metal chloride + oxygen (O2) • Binary compounds  elements • Electrolysis of molten salts (ionic compounds)  elements

34. Decomposition: Special Cases • Decomposition of peroxides: peroxide  water + oxygen (O2) • Ammonium compounds acid + ammonia; the acid may decompose (NH4)2CO3 (s)  2NH3(g) + CO2(g) + H2O(l) NH4NO2 (s)  N2(g) + 2H2O (l) NH4NO3(s)  N2O (g) + 2H2O(l)

35. Types of Equations Used to Describe Reactions in Solution • Molecular: overall reaction stoichiometry- not actual forms MgSO4 (aq) + Na2CO3 (aq) --> MgCO3 (s) + Na2SO4 (aq) 2. Complete Ionic: reactants and products that are strong electrolytes are represented as ions. Mg2+ + SO42- + 2Na+ + CO32- --> MgCO3 (s) + 2Na+ + CO32- 3. Net Ionic: includes only those solution components undergoing a change. Spectator ions not included. Mg2+ + CO32- --> MgCO3 (s)

36. Writing Equations Write a balanced molecular, ionic and net ionic equations for the following reactions: • Solution of silver nitrate was added to a solution of sodium chromate • A piece of solid zinc was placed in a solution of Copper(II) chloride 3.1

37. Single Replacement Reactions • A0 + B+C- A+C- + B-( metals) • A0 + B+C- B+A- + C0(halogens) • All are Redox / NonRedox? • Metal replaces less active metal • Nonmetal replaces less active nonmetal • Not all occur • Activity series (more about later)

38. Single Replacement (SR) : Metal Replaces Less Active Metal • Zinc metal reacts with copper (II) sulfate in water solution Molecular equation: Net Ionic equation : Redox?

39. Single Replacement Reactions Industrial processes to produce metals: Vanadium (V) oxide (solid) reacts with molten calcium metal Titanium (IV) metal solid reacts with molten magnesium • Can we write net ionic equations for the above?

40. Single R: Hydrogen Displacement Write molecular and net ionic equations: • From water sodium metal reacts with water aluminum reacts with steam magnesium reacts with hot water • Which metals will replace hydrogen from cold water? • Which metals will replace hydrogen from hot water? • Which metals will replace hydrogen from steam?

41. Single R: Hydrogen Displacement Write molecular and net ionic equations: • From acids zinc metal reacts with hydrochloric acid aluminum metal reacts with sulfuric acid

42. Single R: Halogen Displacement Write molecular and net ionic equations: • Chlorine gas reacts with aqueous solution with sodium bromide • Activity series: F2 > Cl2 > Br2 > I2

43. Double Replacement (Metathetical Reactions) • A+B- + C+D- A+D- + C+B- • Reactions occur to completion when: Precipitate is produced Gas is produced Molecular substance such as H2O, CO2, NH3, SO2 are produced • Redox or NonRedox ?

44. Methatetical Reactions (Double Replacement) • A+B- + C+D-→ A+D- + C+B- • Start with two reactants (acid, base, salt, water) and produce two products • The products can be predicted by exchanging the positive parts of the two reactants

45. Double Replacement : Precipitate Formations • Solubility Rules: MEMORIZE • All reactions occur in aqueous solution

46. Double Replacement: Precipitate Write the molecular, complete ionic, and net ionic forms • Aqueous nickel (II) chloride reacts with aqueous sodium hydroxide

47. Double Replacement: Precipitate • Aqueous sodium sulfide reacts with lead (II) nitrate molecular net ionic • Aqueous potassium carbonate reacts with barium chloride molecular net ionic

48. Double Replacement: Precipitate • Predict whether a reaction will occur in each of the following case. If so, write a net ionic equation for the reaction. If no reaction occurs, write NR after arrow. • Al2(SO4)3 + NaOH  • K2SO4(aq)+FeBr3(aq)  • CdCl2(aq) + (NH4)2S(aq) 

49. Double Replacement: Precipitate Complete and balance the following reactions: • KCl (aq) + Pb(NO3)2 (aq)  • AgNO3 (aq) + MgBr2 (aq) 

50. Selective Precipitation Precipitation reactions allow us to target specific substances, and separate and recover them from a solution. Example: A solution contains Ca2+, Cu2+, and Pb2+. What anions can we add, and in what order , to separate and recover each cation?