Chapter 4. Reactions in Aqueous Solution. Continued

Observing and Predicting Reactions Chapter 4 z Reactions in Aqueous Solution z How do we know whether a reaction occurs? What observations indicate a reaction has occurred? In your groups groups, make a list of changes that indicate a chemical reaction has occurred. 1 Properties of Aqueous Solutions z z z z 2 Electrolytes in Aqueous Solutions Substances behave differently when they are placed in water, specifically ionic versus covalent compounds. One breaks apart in water, the other does not. Which one is more likely to be pulled apart by water molecules? Electrolytes are ionic and strong acid solutions (e.g., GatoradeTM); Nonelectrolytes are covalent compounds (e.g., sugar); weak electrolytes are in between. Electrolytic Properties z z z Strong electrolyte: substance that, when dissolved in water, results in a solution that can conduct electricity (NaCl) Weak electrolyte: substance that is a poor conductor of electricity when dissolved in water (CH3COOH – vinegar) Nonelectrolyte: substance that doesn’t conduct electricity when dissolved in water (CH3OH – methanol) Strong/Weak Electrolytes Strong, weak, or nonelectrolyte z z z z z z z Place the following substances into groups. NaCl C6H12O6 CH3COOH NH3 HCl NaOH 1 Electrolytes in Aqueous Solutions Electrolytes 3 substances: A2X, A2Y, and A2Z. Which one is the strongest electrolyte? The weakest? (Water has been omitted for clarity.) z Strong/Weak Electrolytes Properties of Aqueous Solutions z z Most reactions in general chemistry take place in an aqueous environment. What does that mean? Terms: Ways Reactions Occur Three general categories: z  Precipitation: insoluble (solid) product is formed from aqueous solutions  Acid Acid-base base neutralization: acid and base react to form water and a salt (ionic compound)  Solution: homogeneous mixture of two or more  Oxidation-Reduction: electrons are transferred substances  Solute: substance present in smaller amount  Solvent: substance present in greater amount  Aqueous solution: solvent is water Products of Precipitation Reactions z z z Precipitation reactions always begin with two ionic compounds. Example: NaCl (aq) + AgNO3 (aq) Æ ? Draw these compounds in two separate aqueous environments. What are the possible products when they are combined? between atoms in reaction z z z Combination Decomposition Single-replacement (metal or hydrogen) Predict Products and Types of Reactions + NO3Ag+ Na+ Ag+ NO3- Æ - NO3 Ag+ Na+ Cl- Cl- Na+ Cl- Write formulas of products (based on charges), predict phases (Solubility Rules), and balance the equation. 2 Solubility Rules Solubility Rules z z z z z z z z If not covered by the rules, it is probably insoluble insoluble.. 13 Molecular, Ionic, and Net Ionic Equations z Formation of Silver Chloride There are 3 ways to represent ppt reactions: z Molecular equation: z Ionic equation (write separate ions for soluble (aq) compounds):  NaCl(aq) + AgNO3(aq) Æ AgCl(s) + NaNO3(aq)  As whole compounds (molecular equation)  As ionic species (ionic equation) – more accurate  As participants in reaction (net ionic equation) z z Determine if the following ionic compounds will be soluble (aq) or insoluble (s) in water: K2CO3 B SO4 BaSO PbI2 NaClO4 Ag2S (NH4)3PO4 Cu(OH)2  Na+(aq) + Cl-(aq) + Ag+(aq) +NO3-(aq) Æ AgCl(s) + Any aqueous ionic substance is written as a compound (e.g., AgNO3), but this isn’t accurate. What does this look like in water? Solids, liquids, and gases remain as compounds. Na+(aq) + NO3-(aq) z Net ionic equation (cancel any identical ion on both sides of the equation, called spectator ions): Ag+(aq) + Cl-(aq) Æ AgCl(s) z Note: s, l, and g stay together!!!!! Ag+/Cl- ions 15 z z z Precipitation Reactions Acid-Base Reactions Reaction of lead (II) nitrate and potassium iodide. What is the precipitate? Write the molecular, ionic, and net ionic equations. Worked Ex. 4.2 – 4.5; Problems 4.4 – 4.8 z Acid: substance that breaks apart in water to form H+ (e.g., HCl, HNO3, CH3COOH, lemon, lime, vitamin C). z Base: substance that breaks apart in water to form OH- (e.g., NH3, DranoTM, Milk of MagnesiaTM)  HA(aq) Æ H+(aq) + A-(aq)  MOH(aq) Æ M+(aq) + OH-(aq) acid strong acid base weak acid 18 3 Common Acids and Bases z Acid-Base Neutralization You need to KNOW these!!! z z What does “neutralization” mean? How might these two solutions react when combined? Na+ OH- Cl- + Na+ z z Strong acids: HCl, HBr, HI, HClO4, HClO3, H2SO4, HNO3 Strong bases: LiOH, KOH, NaOH, RbOH, CsOH, Ca(OH)2, H+ H+ Cl- OH- Sr(OH)2, Ba(OH)2 19 Acid-Base Neutralization z z z z Acid-Base Reaction Neutralization reaction: reaction between acid and base; products are salt (ionic compound) and water HCl(aq) + NaOH(aq) Æ NaCl(aq) + H2O(l) Acid + base Æ salt + water What are the ionic and net ionic equations for these reactions? HCl Acid + Base → H2O + Salt Water is created from H+ and OH- → z HF(aq) + NaOH(aq) z KOH (aq) + H2SO4 (aq) → z Worked Ex. 4.6; Problems 4.9, 4.10 2HCl(aq) + Mg(OH)2(s) → 2H2O(l) + MgCl2(aq) The salt is created from spectator ions Complete/Balance These Equations z z z z z z z Mg(OH)2 Solubility Rules Complete and balance these equations. Write ionic and net ionic equations, if applicable. Na2S(aq) + CuCl2(aq) Æ M SO4(aq) MgSO ( ) + BaCl B Cl2(aq) ( )Æ KNO3(aq) + CaCl2(aq) Æ CuSO4(aq) + NaOH(aq) Æ HF(aq) + NaOH(aq) Æ KOH(aq) + H2SO4(aq) Æ 23 If not covered by the rules, it is probably insoluble insoluble.. 24 4 Oxidation-Reduction Reactions Assigning Oxidation Numbers z Oxidation-Reduction (redox) reactions: electron-transfer reactions z z When iron rusts, it loses electrons to form a cation oxygen gain electrons to form an anion: cation, 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) Use oxidation number rules to determine gain and loss of electrons. Oxidation numbers are assigned as if elements in compounds completely transferred electrons (like in ionic compounds). z z z z z 1) An atom (or molecule) in its elemental state has an oxidation number of 0. 2) An atom in a monatomic ion (Na+, Cl-) has an oxidation number identical to its charge charge. 3a) Hydrogen has an oxidation number of +1, unless it is combined with a metal, in which case it has an oxidation number of –1. 3b) Oxygen usually has an oxidation number of -2. Oxygen in peroxides (O22-) has an oxidation number of -1. 26 Rules continued z z Example 3c) Halogens usually have an oxidation number of -1 (except when bonded to oxygen or in polyatomic ions). 4) The sum of oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion. (Example: NaCl = 0, SO42- = -2) z  4a) For binary ionic compounds, the position of the z H2SO4 z z H = +1; O = –2 S is unknown, so leave this for last.  The overall charge g on this compound p is 0.  Use algebra to solve for S:  2(+1) + 1(x) + 4(-2) = 0 z Solve for each element: MgCr2O7 Worked Ex. 4.7; Problem 4.11 element in the periodic table may be useful:  Group IA: +1; Group IIA: +2; Group VIIA: –1; Group VIA: –2; Group VA: –3 27 Assigning Oxidation Numbers z 28 Oxidation-Reduction Reactions Determine values of the oxidation number of each element in these compounds or ions: H2O SO2 CCl4 H2O2 Fe3(PO4)2 MnO4NaNO3 KClO4 z z z Oxidized: atom, molecule, or ion becomes more positively charged  Loss of electrons is oxidation (LEO) R d Reduced: d atom, t molecule, l l or iion b becomes lless positively charged (reduced charge)  Gain of electrons is reduction (GER) Or: OIL RIG (oxidation is loss; reduction is gain) Redox 1 Redox 2 29 5 Assigning Oxidation Numbers Oxidation-Reduction Reactions z z z z Redox Reactions Oxidation-Reduction Reactions z z The substance oxidized causes the other substance to be reduced and is called the reducing agent. The substance reduced causes the other substance to be oxidized and is called the oxidizing agent. 4 Fe(s) + 3O2(g) Æ 2Fe2O3(s) Worked Ex. 4.8; Problems 4.12, 4.13 Identify the element or ion oxidized/reduced. Also identify the oxidizing agent and the reducing agent. Zn(s) + 2HCl(aq) Æ ZnCl2(aq) + H2(g) z Combination (1 product) z Decomposition (1 reactant) – usually give off gases z Single Replacement (or Displacement) (start and end with an element and a compound)  Na(s) + Cl2(g) Æ  CuCO3(s) Æ Reduced  Zn(s) + HCl(aq) Æ Oxidized 34 Metals versus Nonmetals z Metals tend to react with water to form bases:  2Na (s) + 2H2O Æ 2NaOH + H2  MgO (s) + H2O Æ Mg(OH)2 z CO2 (s) + H2O (l) Æ H2CO3 (aq) Alkali metals + water Nonmetals tend to react with water to form acids:  2F2 (g) + 2H2O Æ 4HF + O2  CO2 (g) + H2O Æ H2CO3 6 Combination Reactions z Combination Reactions element + element → compound z Mg + O2 K + Cl2  H2(g) + O2(g) Æ z metal + nonmetal → ionic compound p  Na(s) + Cl2(g) z nonmetal + nonmetal → covalent compound  z → C(s) + O2(g) → Why are these redox reactions? 37 Decomposition Rxns Produce Gases z z Compound → 2 elements; element + compound; or 2 compounds Oxides, peroxides z z z  Give off O2 z Nitrates z Carbonates z Ammonium salts Decomposition Reactions Δ NH4Cl (s) ⎯⎯→ Δ NiCO3(s) ⎯⎯→ Δ CuO(s) ⎯⎯→ Δ ⎯ ⎯→ →  Give off NO2 , NO2 Give off CO2  Give off NH3 39 Single-Displacement Reactions z element + compound → compound + element (The more metallic/active element in the compound is displaced.) z Metal Displacement Single Displacement: Cu + AgNO3  Zn(s) + Cu(NO3)2(aq) Æ Cu(s) + Zn(NO3)2(aq) z Hydrogen Displacement  Mg(s) + HCl(aq) → MgCl2(aq) + H2(g) Alkali Metals Zn+SnCl2 41 42 7 Single Displacement: Fe + Cu(NO3)2 Activity Series z z The higher the metal on the activity series, the more active ti that th t metal. Translation: higher metals on the chart will form ions as products. 44 Predict whether a reaction occurs z z z z z z z Na + H2O Fe + H2O Fe + Cr(NO3)2 Ni + Pb2+ Ag + Mg2+ Zn + Co2+ Worked Ex. 4.9; Problems 4.14, 4.16 Determining Order of Activity z React w/ cold water z React w/ steam z z z React w/ acid Four metals (A, B, C, and D) are tested and the following observations are recorded. A + B+ Æ no reaction C + B+ Æ reaction ti occurs D + A+ Æ reaction occurs D + B+ Æ no reaction 45 Reaction Types Combustion Reactions z z z z Burning hydrocarbons CxHyOz + O2 (g) Æ CO2 (g) + H2O (g) Methanol, CH3OH _CH3OH (l) + _O2 (g) Æ _CO2 (g) + _H2O (g) Types of Reactions Summary z z z z z z Reaction Types Animated Precipitation: use Solubility Rules (AB + CD Æ AD + CB) Acid-Base Neutralization: acid + base Æ salt + water ((AB + CD Æ AD + H2O)) Combination: start with elements (A + B Æ AB) Decomposition: often produces gas (AB Æ A + B) Single Displacement: use Activity Series to predict if a reaction occurs (A + BC Æ B + AC) Combustion:  Hydrocarbon + O2(g) Æ CO2(g) + H2O(g) 8 Predict Products and Types of Reactions Predict Products and Types of Reactions Cl Cl + H+ H+ Predict Products and Types of Reactions ClClH+ + OHK+ OH- K+ ClH+ Cl- Na Na Na Na Na Na H+ Cl- + H+ Cl- Ca Ca Ca Ca Ca Ca Ca Classify (and balance) these rxns 1. Co(s) + AgNO3(aq) Æ 2. Fe(s) + HCl(aq) Æ Δ 3. Na2CO3(s) ⎯⎯→ 4. Ca(s) + H2O(l) Æ Δ 5. CaCO3(s) ⎯⎯→ 6. HClO4(aq) + KOH (aq) Æ 7. BaCl2(aq) + Na2SO4(aq) Æ 52 Classify (and balance) these rxns Δ 8. HgO(s) ⎯⎯→ 9. LiOH(aq) + H2SO4(aq) Æ 10. Na2CrO4(aq) + Ni(NO3)2(aq) Æ 11 Li(s) + O2(g) Æ 11. 12. Mg(OH)2(aq) + 2HCl(aq) Æ 13. NH3(g) + HCl(g) Æ Δ 14. NiCO3(s) ⎯⎯→ 15. Ca(s) + F2(g) Æ Classify these reactions by type Single disp. 1. Co(s) + 2AgNO3(aq) Æ Co(NO3)2(aq) + 2Ag(s) Single disp. 2. Fe(s) + 2HCl(aq) Æ FeCl2(aq) + H2(g) Decomp. 3. Na2CO3(s) Æ Na2O(s) + CO2(g) Single disp. 4. Ca(s) + 2H2O(l) Æ Ca(OH)2(aq) + H2(g) Decomp. 5. CaCO3(s) + heat Æ CaO (s) + CO2(g) Acid base neut Acid-base neut. 6. 6 HClO4(aq) + KOH(aq) Æ KClO4(aq) + H2O(l) Precip. 7. BaCl2(aq) + Na2SO4(aq) Æ BaSO4(s) + 2NaCl(aq) Decomp. 8. 2HgO(s) Æ 2Hg(l) + O2(g) Acid-base neut. 9. 2LiOH(aq) + H2SO4(aq) Æ Li2SO4(aq) + 2H2O(l) Precip. 10. Na2CrO4(aq) + Ni(NO3)2(aq) Æ 2NaNO3(aq) + NiCrO4(s) Combo. 11. 4Li(s) + O2(g) Æ 2Li2O(s) Acid-base neut. 12. Mg(OH)2(aq) + 2HCl(aq) Æ 2MgCl2(aq) + H2O(l) Combo. 13. NH3(g) + HCl(g) Æ NH4Cl(s) Decomp. 14. NiCO3(s) Æ NiO(s) + CO2(g) Combo. 15. Ca(s) + F2(g) Æ CaF2(g) 9