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CH 4 Reactions in Aqueous Solutions

CH 4 Reactions in Aqueous Solutions. Types of Equations Used to Describe Reactions in Solution. Molecular : overall reaction stoichiometry - not actual forms MgSO 4 ( aq ) + Na 2 CO 3 ( aq ) --> MgCO 3 (s) + Na 2 SO 4 ( aq )

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CH 4 Reactions in Aqueous Solutions

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  1. CH 4 Reactions in Aqueous Solutions

  2. 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)

  3. Types of Chemical Reactions Combination (Synthesis) reaction A + B  AB Decomposition reactionsAB  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

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

  5. Combination Reactions (Synthesis): A + B → C Metals + nonmetals (Redox?) Cesium metal + iodine → Zinc + sulfur → Magnesium + nitrogen → Metal Oxides (most are solid) + Water: (Redox?) Magnesium oxide + water → Lithium oxide + water → Aluminum oxide + water → Iron(III) oxide + water →

  6. Combination Reactions (Synthesis): 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)

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

  8. Decomposition Reactions : C → A + B Reverse of 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?

  9. Decomposition Reactions (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

  10. Decomposition Reactions : (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)

  11. 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)

  12. 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

  13. Single Replacement or Displacement Reactions A0 + B+C- A+C- + B- ( metals) A0 + B+C- B+A- + C0 (halogens) All are Redox Active metal replaces less active metal Active metal replaces H in water or acids Nonmetal replaces less active nonmetal Activity series – used to predict Rx Standard Reduction Potential Chart and SHE

  14. Single Replacement Reactions If a < reactive element is combined with a > reactive element in compound form → no Rx 1. Zinc metal reacts with copper (II) sulfate in water solution Molecular equation: Net Ionic equation : Redox? 2. zinc metal reacts with hydrochloric acid 3. aluminum metal reacts with sulfuric acid

  15. Single Replacement Reactions Write formula and net ionic equations: sodium metal reacts with cold a 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? Activity series of metals: http://www.chem.vt.edu/RVGS/ACT/notes/activity_series.html

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

  17. Double Replacement Reactions orMetathetical 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 ?

  18. Double Replacement Reactions orMetathetical Reactions Write the molecular complete ionic net ionic forms Aqueous nickel (II) chloride reacts with aqueous sodium hydroxide Aqueous sodium sulfide reacts with lead (II) nitrate Aqueous potassium carbonate reacts with barium chloride

  19. Double Replacement Reactions orMetathetical Reactions • 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) 

  20. Double Replacement: Gas Formation Common gases formed in DR Rx S2- + acid → H2S (g)+ salt CO32- + acid → CO2 (g)+ H2O + salt SO3- + acid → SO2 (g)+ H2O + salt NH4+ + OH- + Δ → NH3 g) + H2O + salt 1. Sodium carbonate reacts with hydrochloric acid 2. Ammonium chloride reacts with sodium hydroxide 3. Magnesium nitride reacts with water 4. Calcium sulfite reacts with hydrobromic acid 5. Sodium chloride + sulfuric acid 6. Sodium sulfide reacts with hydrochloric acid

  21. 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?

  22. Combustion Write the products and balance the following combustion reaction: C6H12O6 (s) + O2 → C3H8O3 + O2 → CH3OH + O2

  23. Acids and Bases: Arrhenius Acid • Any substance that releases H+ ion in aqueous solution Base • Any substance that releases OH- ion in aqueous solution

  24. Brønsted-Lowery Acid-Base Definitions Anacid is a substance that donates a proton (H+) to a base A baseis a substance that accepts a proton (H+) from an acid conjugate base conjugate. acid

  25. Brønsted-Lowery Acid-Base Definitions An acid is a substance that donates a proton (H+) to a base A base is a substance that accepts a proton (H+) from an acid Acid-base reactions can be reversible: reactantsproducts or productsreactants Conjugate acid: ____________ Conjugate base: _________________

  26. Compounds that act as Brönsted Acid and Base Write equations for the following reactions. Identify the acid, base, conjugated acid and conjugated base: • HSO4-(aq) + H2O(l) → • HSO4-(aq) + H2O(l) → • H2O(l) + H2O(l) → • HCO3-1 (aq) + H2O(l)

  27. Important Acids and Bases Strong Acids: HCl hydrochloric HBr hydrobromic HI hydroiodic HNO3 nitric H2SO4 sulfuric HClO4 perchloric Weak Acid: CH3CO2H acetic Strong Bases: NaOH sodium hydroxide KOH potassium hydroxide Ca(OH)2 calcium hydroxide Weak Base: NH3 ammonia

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

  29. STRONG acids in water: 100% of acid molecules form ions: HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) H3O+ is hydronium ion

  30. WEAK acids in water: ~5% or less of acid molecules form ions (acetic, H3PO4, H2CO3)

  31. Polyprotic Acids: multiple acidic H atoms H2SO4 H+ + HSO4- HSO4-  H+ + SO42- Not all H’s are acidic: CH3CO2H

  32. If H3PO4 reacts as an acid, which of the following can it not make? 1. H4PO4+ 2. H2PO4- 3. HPO42- 4. PO43-

  33. Reactions Involving Weak Bases HCl(aq) + NH3(aq)  NH4+(aq) + Cl-(aq) Net-Ionic Equation: NH3(aq) + H+(aq)  NH4+(aq) Spectator Ion?

  34. Acid-Base Reactions: Neutralization The “driving force” is the formation of water. NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(liq) Net ionic equation OH-(aq) + H+(aq) → H2O(liq) “Spectator Ions”? __________________________ This applies to ALL reactions of STRONG acids and bases.

  35. Acid-Base Neutralization Rx Polyprotic acids H2SO4 (SA)or H3PO4 (WA) H2SO4 : First H+ is ionized completely H2SO4 → H+ + HSO4- • If base is excess: all H+ form H2O • If equimolar acid + base: only 1 H+ ionizes • Acidic anhydrides (NMO) + Basic anhydrides (MO) : react with H2O before acid or base

  36. CH3CO2H(aq) + NaOH(aq)  Choose the correct answer: 1. CH3CO2H2+(aq) + NaO(aq) 2. CH3CO2-(aq) + H2O(l) + Na+(aq) 3. CH4(g) + CO2(g) + H2O(l) Complete Ionic equation: CH3COOH(aq) + OH-(aq) → CH3COO-(aq) + H2O(l)

  37. HCN(aq) + NH3(aq)  Answer? 1. NH4+(aq) + CN-(aq) 2. H2CN+(aq) + NH2-(aq) 3. C2N2(s) + 3 H2(g)

  38. Hydrolysis Rx- Reverse Neutralization Salt + H2O → molecular species Formation of a weak acid and/or weak base NH4+ + Cl- + H2O → H+ + Cl- + NH4OH NH4Cl : salt from SA (HCl) + WB (NH3 ) Forms acidic solution due to: NH4+> OH- *Salts of SA + WB → Acidic Solution *Salts of SA + WB → Basic Solution *Salts of SA + SB → Neutral Solution *Salts of WA + SB → ?? Check Ka and Kb

  39. Aqueous potassium fluoride undergoes hydrolysis when placed in water. 2. Sodium chloride and water are mixed together. 3. Ammonium fluoride and water are mixed together.

  40. Oxidation-Reduction Reactions Redox reactions: involve a transfer of electrons. Assigning oxidation states to an element in a molecule: K2CrO4 LiSCN

  41. LEO GER: Leo the Lion Says GEROIL RIG Loss of Electrons is Oxidation Gain of Electrons is Reduction Oxidation Involves Loss Reduction Involves Gain

  42. Determination of Oxidation States Fe2O3 + 2Al → Al2O3 + 2Fe Iron (III) gains 3 electrons to become elemental iron. Elemental aluminum lost 3 electrons to become the aluminum ion. Write the half reactions:

  43. N2 H4 + N2O4→ N2 + H2O The combustion of hydrazine with dinitrogen tetroxide helps to keep the space shuttle in Earth Orbit. Is it a Redox reaction? Explain.

  44. Fe2O3 + 2Al → Al2O3 + 2Fe Iron (III) ion gained electrons. It has been reduced. The aluminum lost electrons. It has been oxidized. The oxidizing agent is the species that is reduced (Iron (III)). The reducing agent is the species that is oxidized (aluminum).

  45. Rules for Assigning Oxidation States (OS) 1. OS of an atom in an element is 0. Na (s), O2 (g) 2. OS of a monatomic ion is the same as its charge. Na+ OS = +1, Cl- OS = -1 3. In its covalent compounds with nonmetals, hydrogen is assigned an OS of +1. HCl, NH3, H2O. 4. Oxygen is assigned an OS of -2 in its covalent compounds. CO, CO2, SO2, SO3 The exception to this rules occurs in peroxides (compounds contains the O22- group), where each oxygen is assigned an OS of -1. H2O2

  46. 5. In binary compounds the element with the greater attraction for the electrons in the bond is assigned a negative OS equal to its charge in its ionic compounds. HF, NH3, H2S, HI 6. The sum of the oxidation states must be zero for an electrically neutral compound and must be equal to the overall charge for an ionic species. NH4+, CO32-

  47. Which Atoms Undergo Redox? 2H2 (g) + O2 (g) → 2H2O (g) Zn (s) + Cu2+(aq) → Zn2+ (aq) + Cu(s) 2AgCl (s) + H2 (g) → 2H+ (aq) + 2Ag(s) + 2Cl- (aq) 2MnO4- (aq) + 16H+ (aq) + 5C2O42- (aq) → 2Mn2+(aq) + 10 CO2 (g) + 8 H2O (l)

  48. Methods for Balancing Redox Reactions 1. Oxidation states method CdS + I2 + HCl → CdCl2 + HI + S (1, 1, 2, 1, 2, 1) Cl2 + Ca(OH)2 → CaCl2 + Ca(ClO3) + H2O ( 6, 6, 5, 1, 6) 2. Half reaction method

  49. Balancing a Redox Equation by the Oxidation States Method • Assign the oxidation states of all atoms. • Decide which element is oxidized and determine the increase in oxidation state. • Decide which element is reduced and determine the decrease in oxidation state. • Choose coefficients for the species containing the atom oxidized and the atom reduced such that the total increase in oxidation state equals the total decrease in oxidation state. • Balance the remainder of the equation by inspection.

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