Reactions in Aqueous Solution

Reactions in Aqueous Solution
Chapter 4 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

A solution is a homogenous mixture of 2 or more substances
The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount Aqueous Solution: A solution in which the solvent = water Solutes in an aqueous solution may be ionic compounds, acids, or molecular compounds. Solution Solvent Solute Soft drink (l) H2O Sugar, CO2 Air (g) N2 O2, Ar, CH4 Soft Solder (s) Pb Sn 4.1

Behavior of Solutes in Aqueous Solution & Their Classification
Water causes some solutes to dissociate or ionize. Dissociation or Ionization is the separation of a compound into separated cation and anion. The solutes that undergo dissociation or ionization include water soluble ionic compounds and acids. Molecular compounds dissolve intact in aqueous solutions. They do not dissociate or ionize. When a solute dissociates (ionizes), ions are generated in solution and the solution is therefore able to conduct electricity. If a lot of ions are generated, the solution conducts electricity strongly. If only a small amount of ions are generated, the ability to conduct is weak. If no ions are generated, the solution cannot conduct electricity.

An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. That is all solutes that ionize or dissociate are electrolytes. A strong electrolyte ionizes nearly completely generating a lot of ions. A weak electrolyte only ionizes partially and hence fewer ions are generated. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. nonelectrolyte weak electrolyte strong electrolyte 4.1

Strong Electrolytes: all water soluble ionic compounds (see solubility rules) & the 7 strong acids (memorize list) 7 Strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3 Weak Electrolytes: All weak acids (not on list of 7: HC2H3O2, HF, HCN, H3PO4, H2CO3), NH3 Nonelectrolytes: All molecular compounds that are not weak acids or weak bases. Examples: alcohol, sugar

Conduct electricity in solution?
Cations (+) and Anions (-) Strong Electrolyte – 100% dissociation NaCl (s) Na+ (aq) + Cl- (aq) H2O Weak Electrolyte – not completely dissociated HC2H3O C2H3O2- (aq) + H+ (aq) 4.1

Nonelectrolyte does not conduct electricity
No cations (+) and anions (-) in solution C6H12O6 (s) C6H12O6 (aq) H2O 4.1

Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature. 4.2

Metathesis Reactions = Double Displacement Reactions
Metathesis Reaction: A reaction in which two reactants exchange ions (cations or anions). These reactions therefore involve ionic compounds, acids as reactants. A metathesis reaction will result in a net change in the solution if one of the following happens when ions are exchanged: At least 1 new pairing of ions results in the formation of a gas which escapes from the solution. (bubbling may be observed). Gases commonly occurring as products: CO2, NH3, H2S, SO2 At least 1 new pairing of ions results in the formation of a water insoluble compound or precipitate. Such a metathesis reaction is called a precipitation reaction. (settling of a solid, cloudiness or turbidity will be observed) At least 1 new pairing of ions results in the formation of a weak electrolyte or a nonelectrolyte. A visual change may not be observed for these reactions. The reaction between an acid & a base falls under this category

Equations for Metathesis Reactions
1. Molecular Equation = Balanced Formula Equation: Balanced chemical equation with physical states included. Physical State Substances l : liquid Only H2O s: solid Only for water-insoluble compounds (solubility rules g: gas Only CO2, NH3, H2S, SO2 aq: aqueous All water-soluble compounds (solubility rules), acids, ions 2. Complete Ionic Equation: Such an equation is obtained when all strong electrolytes in molecular equation are dissociated. 3. Spectator Ions & Net Ionic Equation: Ions that occur in the same physical state on both sides of a complete ionic equation are spectator ions. Omitting them yields the net ionic equation. The net ionic equations describes the net change in solution.

Compounds that decompose in Metathesis Reactions & Their Decomposition Equations
3 compounds, namely, NH4OH, H2CO3, H2SO3, DECOMPOSE (different from dissociate) in aqueous solutions. These compounds are replaced by the products of their decomposition in all equations (molecular, complete ionic & net ionic) in which they occur.

Writing Net Ionic Equations
Write the balanced molecular equation. Write the ionic equation showing the strong electrolytes completely dissociated into cations and anions. Cancel the spectator ions on both sides of the ionic equation Check that charges and number of atoms are balanced in the net ionic equation Write the net ionic equation for the reaction of silver nitrate with sodium chloride. AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Ag+ + NO3- + Na+ + Cl AgCl (s) + Na+ + NO3- Ag+ + Cl AgCl (s) 4.2

Precipitation Reactions
Precipitate – insoluble solid that separates from solution PbI2 precipitate Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2NaNO3 (aq) molecular equation Pb2+ + 2NO3- + 2Na+ + 2I PbI2 (s) + 2Na+ + 2NO3- ionic equation Pb2+ + 2I PbI2 (s) net ionic equation Na+ and NO3- are spectator ions 4.2

Precipitation of Lead Iodide
Pb2+ + 2I PbI2 (s) PbI2 4.2

Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus
fruits contain citric acid. React with certain metals to produce hydrogen gas. 2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g) React with carbonates and bicarbonates to produce carbon dioxide gas 2HCl (aq) + CaCO3 (s) CaCl2 (aq) + CO2 (g) + H2O (l) Aqueous acid solutions conduct electricity. 4.3

Diprotic acid: H2SO4: ionize in steps losing 1 H+ per step
Monoprotic acids HCl H+ + Cl- Strong electrolyte, strong acid HNO H+ + NO3- Strong electrolyte, strong acid HC2H3O H+ + C2H3O2- Weak electrolyte, weak acid Diprotic acid: H2SO4: ionize in steps losing 1 H+ per step H2SO H+ + HSO4- Strong electrolyte, strong acid HSO H+ + SO42- Weak electrolyte, weak acid Triprotic acids : H3PO4 H3PO H+ + H2PO4- Weak electrolyte, weak acid H2PO H+ + HPO42- Weak electrolyte, weak acid HPO H+ + PO43- Weak electrolyte, weak acid 4.3

Bases (Arrhenius Definition): These are compounds that increase the amount of OH- in aqueous solution. Examples: all water soluble metal hydroxides; NH3 Strong bases: These generate a lot of OH- in aqueous solution. Examples: all water soluble metal hydroxides; NaOH(aq)  Na+(aq) + OH-(aq) (dissociation) Ba(OH)2(aq)  Ba+2(aq) + 2OH-(aq) (dissociation) Weak Bases: These generate a small amount of OH- in aqueous solution. Example: NH3 NH3(aq) + H2O  NH4+(aq) + OH-(aq) (hydrolysis)

Neutralization Reaction
acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H2O H+ + Cl- + Na+ + OH Na+ + Cl- + H2O H+ + OH H2O HF (aq) + NaOH (aq)  NaF (aq) + H2O HF(aq) + Na+ + OH-  Na+ + F- + H2O HF(aq) + OH-  F- + H2O 4.3

Metal hydroxide + acid  water + salt; salt: cation = metal cation; anion = anion of acid
Balancing the above equation: number of H+ from acid = number of OH- from base = number of H2O formed 3Ca(OH)2 + 2H3PO4 6H2O + Ca3(PO4)2; salt = Ca3(PO4)2; Ca(OH)2 + H2SO4 2H2O + CaSO4; salt = CaSO4; 3NaOH + H3PO4 3H2O + Na3PO4; salt = Na3PO4; Ammonia + acid  ammonium salt; cation = NH4+; anion = anion of acid; To balance: number of H+ from acid= number of NH3 = number of NH4+ 3NH3 + H3PO4  (NH4)3PO4; ammonium salt = (NH4)3PO4; 2NH3 + H2SO4  (NH4)2SO4; ammonium salt = (NH4)2SO4 NH3 + HNO3  NH4NO3; ammonium salt = NH4NO3

Oxidation-Reduction Reactions
(electron transfer reactions) 2Mg Mg2+ + 4e- Oxidation half-reaction (lose e-) O2 + 4e O2- Reduction half-reaction (gain e-) 2Mg + O2 + 4e Mg2+ + 2O2- + 4e- 2Mg + O MgO 4.4

Redox reactions or oxidation-reduction reactions involve electron transfer from one reactant to another. Oxidation is the loss of electrons: OIL Reduction is the gain of electrons: RIG The two are dependant and you cannot have one without other The reactant that loses electrons is said to undergo oxidation. The reactant that gains electrons is said to undergo reduction The reactant that undergoes oxidation causes the other reactant to undergo reduction and is hence called the REDUCING AGENT. The reactant that undergoes reduction takes away electrons from the reactant causing its oxidation. It is therefore the OXIDIZING AGENT. Reducing agent is oxidized, the oxidizing agent is reduced. The electrons flow from the reducing agent to the oxidizing agent When a metal & nonmetal react, metal =reducing agent, nonmetal = oxidizing agent.

The transfer of electrons in a redox reaction is followed by looking at changes in OXIDATION NUMBER.
Oxidation numbers are imaginary numbers obtained from a set of rules (MEMORIZE) Every atom in a molecule or polyatomic ion is assigned an oxidation number OXIDATION NUMBERS SHOULD NOT BE CONFUSED WITH CHARGES FOR POLYATOMIC SPECIES The oxidation number of monoatomic ions (NOT POLYATOMIC IONS) = the charge of ion The oxidation number of elements (NOT COMPOUNDS) = 0 When a reactant gives up electrons, the oxidation number of at least one of its constituent elements becomes less negative or more positive (larger). When a reactant gains electrons, the oxidation number of at least one of its constituent elements becomes more negative or less positive (smaller).

Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
Oxidation number Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H2, O2, P4 = 0 In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 The oxidation number of oxygen is usually –2 in most of its compounds & polyatomic ions. Exceptions: In OF2, the oxidation number of O is +2. In H2O2 and O22- the oxidation number of O is –1. The oxidation number of O in O2- (superoxide) is -1/2. The oxidation number of hydrogen is +1 in its compounds with other nonmetals and it -1 in its compounds with metals. The oxidation number of Group IA metals in all their compounds are +1, IIA metals are +2 and fluorine is always –1 in all its compounds. Other atoms whose oxidation numbers do not change in their various compounds are Zn, Ag, Al. Oxidation of Zn is +2 in all its compounds, oxidation number of Ag is +1 in all its compounds. The oxidation number of Al is +3 in all its compounds. The sum of the oxidation numbers of all the atoms in a molecule or polyatomic ion is equal to the charge on the molecule or ion. 4.4

Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
Zn Zn2+ + 2e- Zn is oxidized Zn is the reducing agent Cu2+ + 2e Cu Cu is reduced Cu2+ is the oxidizing agent Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s) Cu Cu2+ + 2e- Ag+ + 1e Ag Ag+ is reduced Ag+ is the oxidizing agent 4.4

Cl in Cl2, ClO-, ClO2-, ClO3-, ClO4-
Homework Cr in CrO4-2 Cl in Cl2, ClO-, ClO2-, ClO3-, ClO4- Br in HBr, HBrO, HBrO2, HBrO3, HBrO4, Br2 S in S8, CaS, H2S, H2SO3, H2SO4 N in N2O, NO, NO2, HNO3, HNO2 P in H3PO3, H3PO4

HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4 7.,
Oxidation numbers of all the elements in HCO3- ? O = -2 H = +1 3x(-2) ? = -1 C = +4 4.4

The oxidation numbers of elements in their compounds
4.4

IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2 O = -2
Oxidation numbers of all the elements in the following ? F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2 O = -2 K = +1 3x(-2) ? = 0 7x(-2) + 2x(+1) + 2x(?) = 0 I = +5 Cr = +6 4.4

Types of Oxidation-Reduction Reactions
Combination Reaction A + B C +3 -1 2Al + 3Br AlBr3 Decomposition Reaction C A + B +1 +5 -2 +1 -1 2KClO KCl + 3O2 4.4

Combustion Reaction A + O B +4 -2 S + O SO2 +2 -2 2Mg + O MgO 4.4

Displacement Reaction A + BC AC + B +1 +2 Sr + 2H2O Sr(OH)2 + H2 Hydrogen Displacement +4 +2 TiCl4 + 2Mg Ti + 2MgCl2 Metal Displacement -1 -1 Cl2 + 2KBr KCl + Br2 Halogen Displacement 4.4

The Activity Series for Metals
Hydrogen Displacement Reaction M + BC MC + B M is metal; MC is an ionic compound whose cation is the cation of M and the anion in MC & BC are the same. BC is an acid or H2O B is H2 Ca + 2H2O Ca(OH)2 + H2 Pb + 2H2O Pb(OH)2 + H2 4.4

All metal above H2 in the activity series will react with acids such as HCl to produce H2(g)
Metal(s) + HCl(aq)  metal chloride + H2(g); for all metals more active than H2(g). The oxidizing agent = H+ of the acid The reducing agent = metal Similarly, any metal will react with salts of less active metals. The cation of the less active metal will be reduced to the element. The more active active metal will be oxidized to its corresponding cation. Mmore active + Mless activeC  Mmore activeC + Mless active Mmore active: more active metal; Mless active: less active metal Mless activeC, Mmore activeC: ionic compounds with the same anion C but the cations differ. More active metal(s) + cation of less active metal(aq)  cation of less more active metal (aq) + less active metal(s) Metals will not react with salts of more active metals. Less active metal + cation of more active metal  no reaction.

The Activity Series for Halogens
F2 > Cl2 > Br2 > I2 Halogen Displacement Reaction -1 -1 Cl2 + 2KBr KCl + Br2 I2 + 2KBr KI + Br2 4.4

Disproportionation Reaction Element is simultaneously oxidized and reduced. +1 -1 Cl2 + 2OH ClO- + Cl- + H2O Chlorine Chemistry 4.4

Classify the following reactions.
Ca2+ + CO CaCO3 Precipitation NH3 + H NH4+ Acid-Base Zn + 2HCl ZnCl2 + H2 Redox (H2 Displacement) Ca + F CaF2 Redox (Combination) 4.4

Chemistry in Action: Breath Analyzer
+6 3CH3CH2OH + 2K2Cr2O7 + 8H2SO4 +3 3CH3COOH + 2Cr2(SO4)3 + 2K2SO4 + 11H2O 4.4

Recall: Stoichiometry of a reaction is the relationship between the moles of reactants used and moles of products of formed. This relationship is defined by the ratios of the coefficients of the reactants and products in the balanced chemical equation. aA + bB  cC + dD Moles of A used : moles of B used : moles of C formed : moles of D formed = a : b: c: d To express the ratio in grams of reactants and products, convert moles to grams using the respective molar masses (moles x molar mass = grams)

Solution Stoichiometry
The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. M = molarity = moles of solute liters of solution What mass of KI is required to make 500. mL of a 2.80 M KI solution? M KI M KI volume of KI solution moles KI grams KI 1 L 1000 mL x 2.80 mol KI 1 L soln x 166 g KI 1 mol KI x 500. mL = 232 g KI 4.5

What is the molarity of an aqueous solution containing 50
What is the molarity of an aqueous solution containing 50.0 g of glucose, C6H12O6 in 250 mL of solution. How many grams of Na3PO4 are present in mL of a M aqueous solution of Na3PO4 What is the volume of a M aqueous solution containing 5.00g of urea NH2CONH2. Molarity = moles of solute/L of solution Moles of solute = molarity x L of solution Volume in L of solution = moles of solute/Molarity of solution Moles of solute = (50.0g) (1mol/180.18g) = mol L of solution = 250 x 10-3 L = L Molarity = 0.288mol/0.250 L = 1.11 mol/L = 1.11 M 2. L of solution = L; Molarity = M = moles/L Moles of solute, Na3PO4 = ( mol/L) x L = moles = ( mol) x (163.94g/mol) = g

Solute: 5.00g of NH2CONH2 = (5.00g) x (1mol/60.06g) = 0.0833 mol
Molarity of solution = M = moles/L Volume of solution = mol/molarity = mol/(0.500 mol/L) = L = 167 mL

Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. Dilution Add Solvent Moles of solute before dilution (i) after dilution (f) = MiVi MfVf = 4.5

If 5. 00 mL of a 2. 00 M aqueous solution of KCl is diluted to 0
If 5.00 mL of a 2.00 M aqueous solution of KCl is diluted to 0.500M, what is the final volume of the solution and how much water was added. What is the volume of a 2.50 M solution required to prepare 500 mL of a 0.100M solution by dilution. If mL of water was added to mL of a 4.00M aqueous solution of a solute, what is the resulting molarity of the solution. MiVi = MfVf 5.00 mL x 2.00 M = X mL x M X mL = (5.00 mL x 2.00 M)/0.500 M = 20.0 mL 5.00 mL  20.0 mL; Therefore 15.0 mL of water was added. 3. Mi = 4.00 M; Vi = mL; Vf = mL = mL Mf = ?; Mf = (10.00 mL x 4.00 M)/(110.0 mL) = M

How would you prepare 60.0 mL of 0.200 M
HNO3 from a stock solution of 4.00 M HNO3? MiVi = MfVf Mi = 4.00 Mf = 0.200 Vf = 0.06 L Vi = ? L Vi = MfVf Mi = 0.200 x 0.06 4.00 = L = 3 mL 3 mL of acid + 57 mL of water = 60 mL of solution 4.5

Gravimetric Analysis Dissolve unknown substance in water
React unknown with known substance to form a precipitate Filter and dry precipitate Weigh precipitate Use chemical formula and mass of precipitate to determine amount of unknown ion 4.6

Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete. At this point neither reactant is in excess. They have been mixed in the proportion specified by the balanced equation. Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color 4.7

What volume of a 1.420 M NaOH solution is
Required to titrate mL of a 4.50 M H2SO4 solution? WRITE THE CHEMICAL EQUATION! H2SO4 + 2NaOH H2O + Na2SO4 M acid rx coef. M base volume acid moles acid moles base volume base 4.50 mol H2SO4 1000 mL soln x 2 mol NaOH 1 mol H2SO4 x 1000 ml soln 1.420 mol NaOH x 25.00 mL = 158 mL 4.7

What is the volume of 0. 500 M Ba(OH)2 needed to neutralize 60
What is the volume of M Ba(OH)2 needed to neutralize 60.0 mL of M H3PO4? What is the mass of Ca(OH)2 needed to neutralize 25.0 mL of M H2SO4 What is the molarity of NaOH is 25.0 mL of NaOH neutralized 50.0 mL of M HNO3. 1. moles of Ba(OH)2 : moles of H3PO4 = 3: 2 Moles of H3PO4 = (0.500 mol/L) x ( L) = moles Moles of Ba(OH)2 = .0300 mol H3PO4 x (3 mol Ba(OH)2/2 mol H3PO4) = mol Volume in L of Ba(OH)2 = moles of Ba(OH)2/M of Ba(OH)2 = mol/0.500(mol/L) = L = 90.0 mL 2. moles of Ca(OH)2 : moles of H2SO4 = 1:1 Moles of H2SO4 = (0.400 mol/L) x L = moles Moles of Ca(OH)2 = also moles Mass of Ca(OH)2 = moles x molar mass = mol x (74.10 g/mol) = g

Chemistry in Action: Metals from the Sea
CaCO3 (s) CaO (s) + CO2 (g) CaO (s) + H2O (l) Ca2+ (aq) + 2OH (aq) - Mg2+ (aq) + 2OH (aq) Mg(OH)2 (s) - Mg(OH)2 (s) + 2HCl (aq) MgCl2 (aq) + 2H2O (l) Mg2+ + 2e Mg 2Cl Cl2 + 2e- MgCl2 (aq) Mg (s) + Cl2 (g)

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