Chapter 4: Reactions in Aqueous Solution (Sections )

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1 Chapter 4: Reactions in Aqueous Solution (Sections ) Chapter Goals Be able to: Classify substances as electrolytes or nonelectrolytes. Write molecular, ionic, and net ionic equations for precipitation, acidbase, and redox reactions. Use solubility rules to predict whether a precipitate might form when aqueous salt solutions are mixed. Identify common strong acids and strong bases. Assign oxidation numbers to each atom in a chemical species. In a redox reaction, identify the species oxidized, the species reduced, the oxidizing agent, and the reducing agent. Using an activity series, predict whether a redox reaction will occur when a metal is placed in contact with a solution containing an ion of a different metal. Balance redox reactions by the oxidation-number method or by the halfreaction method. Determine the concentration of a species using data from a redox titration. 4.1 Ways Chemical Reaction Occur: Important types of Aqueous Reactions 1) precipitation reactions process in which reactants an that drops out of solution. formation of this stable product removes material from the aqueous solution and is driving force for reaction (rxn) e.g. takes place when the of two compounds. 2) acid-base neutralization reactions process in which an acid reacts with a base to yield water plus an ionic compound called a salt HNO 3 (aq) + KOH(aq) KNO 3 (aq) + H 2 O(l) 1

2 definitions: acid - compounds that produce H + ions when dissolved in water base - compounds that produce OH ions when dissolved in water driving force - production of water molecule formed from reaction of H + and OH ions in solution 3.) oxidation-reduction reactions (Redox Reactions) - process in which one or more are transferred between reaction partners - driving force - decrease in electrical potential - oxidation numbers and/or charges on atoms in various reactants change 4.2 Electrolytes in Aqueous Solution electrolyte solution substances dissolved in H 2 O to produce Usually ionic compounds (metal and non-metal) exception: molecular compounds that dissociate (split apart) to ions when dissolved in water e.g. HCl (aq) water H + + Cl (aq) non-electrolyte solution substances dissolved in H 2 O which in aqueous solution 2

3 Electrical Conductivity of an Ionic Solution a) solution of an ionic compound electrolyte solution conducts electricity b) solution of a molecular compound nonelectrolyte doesn t conduct electricity Further Classification of Electrolyte Solutions i.) Strong Electrolyte compounds that into ions when dissolved in water ii.) Weak Electrolyte compounds that into ions when dissolved in water Note: to write a dissociation - use double arrow to show that it takes place simultaneously in both directions a dissociation is a dynamic process in which an equilibrium is established between the forward and reverse reactions 3

4 Classification of Common Substances 4.3 Aqueous Reactions and Net Ionic Equations spectator ions ions that of the reaction arrow and therefore during the reaction their role is to balance charge net ionic equations give in the actual reaction and exclude any species that do not take part in the reaction (i.e. spectator ions) focus upon ions undergoing change Ba(NO 3 ) 2 (aq) + Na 2 (SO 4 )(aq) BaSO 4 (s) +2 NaNO 3 (aq) Cancel out spectator ions: 4

5 Net ionic equations must include: i) mass or atom balance same of each element on both sides ii) charge balance must be on both sides Practice Problems: Write the net ionic equations for the following reactions. a) S 8 (s) + 8 O 2 (g) 8 SO 2 (g) b) NiCl 2 (aq) + Na 2 S(aq) NiS(s) + 2 NaCl(aq) c) 2 CH 3 CO 2 H(aq) + Ba(OH) 2 (aq) 2 CH 3 CO 2 Ba(aq) + 2 H 2 O(l) 4.4 Precipitation Reaction and Solubility Rules precipitation reaction (ppt n rxn) chemical reaction which produces an insoluble product a precipitate (ppt) solubility of each compound in a given amount of solvent at a given temperature i) if a substance has a low solubility in water, it is likely to precipitate from an aqueous solution ii) if a substance has a high solubility in water, no precipitate will form. Ba(NO 3 ) 2 (aq) + Na 2 (SO 4 )(aq) BaSO 4 (s) +2 NaNO 3 (aq) soluble soluble insoluble soluble 5

6 Solubility rules: (memorize!) 1. Salts which are always soluble Group 1A cations: Li +, Na +, K +, Rb +, Cs + Ammonium ion: NH 4 + All salts of the NO 3, ClO 3, ClO 4, C 2 H 3 O 2, and HCO 3 ions 2. Salts which are soluble with exceptions: Cl, Br, I ion salts except with Ag +, Pb 2+, & Hg 2 2+ SO 2 4 Ba 2+ ion salts except with Ag +, Pb 2+, Hg 2 2+, Ca 2+, Sr 2+, & 3. Salts which are insoluble with exceptions: O 2 & OH ion salts except with the alkali metal ions, and Ca 2+, Sr 2+, & Ba 2+ ions (group I and II except Mg 2+ ) CO 32, PO 43, S 2, CrO 42, & SO 32 ion salts except with the alkali metal ions and the ammonium ion Practice Problems Predict the solubility of: (a) CdCO 3 (b) MgO (c) Na 2 S (d) PbSO 4 (e) (NH 4 ) 3 PO 4 (f) HgCl 2 Predict whether a precipitate will form for: (a) NiCl 2 (aq) + (NH 4 ) 2 S(aq) (b) Na 2 CrO 4 (aq) + Pb(NO 3 ) 2 (aq) (c) AgClO 4 (aq) + CaBr 2 (aq) 6

7 4.5 Acids, Bases and Neutralization Reactions Acid Substance that dissociates in H 2 O to form hydrogen ions (H + ) In reality, H +, as a bare proton, is too reactive to exist by itself and therefore molecule to form H 3 O + Will sometimes use H + (especially for balancing equations) but correctly written as H 3 O + (aq) to represent aqueous solution e.g. HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl - (aq) Base - substance that dissociates in H 2 O to form hydroxide ions (OH ) Strong and Weak Acids & Bases differentiated by their in water Strong Acid - ionizes completely (i.e. acid solution where substance is completely converted to ions in aqueous solution) - strong electrolytes (e.g. HCl, HNO 3, H 2 SO 4 ) Weak Acid - partially ionized to H + ions in water - molecules containing ionizable hydrogen atom - weak electrolytes (e.g. CH 3 CO 2 H, HF) 7

8 Strong Base completely ionized to OH ions and corresponding cation strong electrolytes e.g. strong bases are hydroxides of Group 1 and Group 2 metals NaOH (s) Na + (aq) + OH (aq) Weak Base reacts with H 2 O molecules, acquiring H + ions and leaving OH ions behind e.g. reactions with ammonia NH 3 (aq) + H 2 O (l) NH 4+ (aq) + OH (aq) Neutralizations Reactions - when an acid and base are mixed in the right stoichiometric proportion, both and forming water and a salt 1) Strong Acid-Strong Base when solutions are mixed, H + and OH - ions combine to form H 2 O molecules Generally, HA(aq) + MOH(aq) H 2 O(l) + MA (aq) Note: A - of salt comes from acid M + of salt comes from base e.g. HCl(aq) + NaOH(aq) net ionic equation: NaCl(aq) + H 2 O(l) 8

9 2) Weak Acid - Strong Base involves a two-step reaction e.g. HB - weak acid NaOH - strong base i) ionization of HB molecule HB(aq) H + (aq) + B (aq) ii) neutralization of H + by OH - ions of NaOH H + (aq) + OH (aq) H 2 O(l) Overall reaction: (add the two steps) HB(aq) + OH (aq) B (aq) + H 2 O(l) Note: in these reaction, the spectator ions, such as Na + are not included in the net ionic equation 3) Strong Acid - Weak Base involves a two step reaction e.g. NH 3 - weak base HCl - strong acid i) reaction of NH 3 with H 2 O NH 3 (aq) + H 2 O(l) NH 4+ (aq) + OH (aq) ii) neutralization of OH ions by the H + of HCl H + (aq) + OH (aq) H 2 O(l) Overall reaction: H + (aq) + NH 3 (aq) NH 4+ (aq) Practice Problems Write ionic and net ionic equations for the following: (a) Ca(OH) 2 (aq) + 2 CH 3 CO 2 H(aq) (b) HBr(aq) + Ba(OH) 2 (aq) (c) HCl(aq) + NH 3 (aq) 9

10 4.6 Oxidation-Reduction Reaction (Redox Reaction) - reactions in aqueous solution involving an exchange of electrons between two species Oxidation loss of electrons oxidation no. more positive Reduction gain of electrons oxidation number more negative Example: Zn (s) + 2 H + (aq) Zn 2+ (aq) + H 2 (g) Reduction or Oxidation (half reactions) - breaking reaction into half reactions: oxidation (loss of e - ): reduction (gain of e - ): NOTES: 1) oxidation and reduction occur together can t have one without the other 2) there is in the number of electrons in a redox reaction the number of electrons lost in the oxidation of one species are gained by the species involved in the reduction 10

11 Oxidation Number - can identify a redox reaction by looking for change in oxidation number of an element in the course of a reaction oxidation number - provides a measure of whether the atom is neutral, electron-rich or electron-poor - comparing oxidation number of an atom before and after reaction allows us to determine whether the atom has gained or lost electrons Note: - oxidation numbers do not necessarily imply ionic charges - simply convenient way to keep track of electrons in a redox reaction Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) 2 Zn (s) + O 2 (g) 2 ZnO (s) Rules for determining oxidation number (apply in order) 1. The sum of oxidation numbers must equal the charge on a molecule, atom, or ion. oxidation no. of an atom in a free element = 0 (e.g. O 2, O 3, Cl 2, P 4, S 8 ) oxidation no. of a monoatomic ion = charge (e.g. Na + is +1, S 2- is -2 ) 2. The oxidation number of F is Group 1A (including H) metals +1 Group 2A metals Oxidation number of O is Cl, Br and I are always -1 in compounds except when combined with F or O (rules 2 and 4 apply) 11

12 A few other rules which may be helpful, but are usually not needed if you follow the first 5 rules in order: 6. When H forms binary compounds with a metal the oxidation number on H is -1 (rule 3 applies) +1-1 e.g. LiH 7. In the formation of a peroxide (H 2 O 2 ) compounds are based on O 2 2- ion in which O has oxidation number of -1 (follows rule 3) Practice Problems a) VOCl 3 f) NO 2 b) Mn 2 O 7 g) C 2 O 2 4 c) CuSO 4 h) S 2 O 2 3 d) CH 2 O i) HNO 3 e) ClO 3 k) KClO Identifying Redox Reactions determine oxidation numbers of each species in the reaction and then identify process as oxidation or reduction oxidation - increase in oxidation number (i.e. loss of e - ) reduction - decrease in oxidation number (i.e. gain in e - ) e.g. 4 Fe (s) + 3 O 2 (g) 2 Fe 2 O 3 (s) 12

13 Reducing and Oxidizing Agents Reducing agent the substance that causes a reduction by giving up e s characteristics: ; loses one or more electrons ; oxidation number at atom increases Oxidizing agent the substance that causes an oxidation by accepting e s characteristics: ; gains one or more electrons ; oxidation number at atom decreases 4.8 The Activity Series of the Elements used to determine whether a reaction occurs between a given ion and a given element depends on the relative ease with which the various species gain or lose electrons i.e. the relative ease with which the species are reduced or oxidized Activity series - ranks elements in order of their reducing ability in aqueous solution Important things to note in Table 4.3: i) elements at the e readily and are stronger reducing agents (easily oxidized) ii) elements at the bottoms give up e and are weak reducing agents (easily reduced) 13

14 ** any element higher in the activity series will reduce the ion of any element lower in the activity series 4.10 Balancing Redox Reactions: The Half-Reaction Method (no oxidation numbers needed) Half-reaction method realize that the overall reaction can be broken into two parts or half-reactions Half-reactions one half-reaction describes of process one half-reaction describes of process Example: Balance for an acidic solution: MnO 4 (aq) + Br (aq) Mn 2+ (aq) + Br 2 (aq) The steps involved follow the same basic procedure as described for the oxidation-number method. 1. Determine oxidation and reduction half-reactions: Oxidation half-reaction: Reduction half-reaction: 14

15 2. Balance for atoms other than H and O: Oxidation: Br (aq) Br 2 (aq) Reduction: MnO 4 (aq) Mn 2+ (aq) 3. Balance for oxygen by adding H 2 O: Oxidation: 2 Br (aq) Br 2 (aq) Reduction: MnO 4 (aq) Mn 2+ (aq) + 4. Balance for hydrogen by adding H + : Oxidation: 2 Br (aq) Br 2 (aq) Reduction: MnO 4 (aq) + Mn 2+ (aq) + 4 H 2 O(l) 5. Balance for charge by adding electrons (e ): Oxidation: 2 Br (aq) Br 2 (aq) + Reduction: MnO 4 (aq) + 8 H + (aq) + Mn 2+ (aq) + 4 H 2 O(l) Final Steps 6. Balance for numbers of electrons by multiplying: Oxidation: [2 Br (aq) Br 2 (aq) + 2 e ] Reduction: [MnO 4 (aq) + 8 H + (aq) + 5 e Mn 2+ (aq) + 4 H 2 O(l)] 7. Combine and cancel to form one equation: Oxidation: Br (aq) Br 2 (aq) + Reduction: MnO 4 (aq) + H + (aq) + 2 Mn 2+ (aq) + H 2 O(l) Final Equation: 15

16 Practice Problems Balance the following equations: 1. MnO 2 (s) + HNO 2 (aq) Mn 2+ (aq) + NO 3 (aq) (acidic) 2. Cr 2 O 72 (s) + C 2 O 42 (aq) Cr 3+ (aq) + CO 2 (g) (acidic) 3. MnO 4 (aq) + NO 2 (aq) MnO 2 (s) + NO 3 (aq) (basic) 4. Zn(s) + HNO 3 (aq) Zn(NO 3 ) 2 (aq) + NH 4 NO 3 (aq) 5. Cr 2 O 72 (aq) + H 2 O 2 (aq) Cr 3+ (aq) + O 2 (aq) (acidic) 4.9 Balancing Redox Reactions: The Oxidation-Number Method Uses the change in oxidation numbers and necessary coefficients so that the number of electrons gained in a reaction equals the number lost. example: Equation balance by usual methods : MnO 4 (aq) + H 2 C 2 O 4 (aq) + 6H + (aq) MB CB Mn 2+ (aq) + 2CO 2 (g) + 4H 2 O This equation is NOT balanced 16

17 MnO 4 (aq) + H 2 C 2 O 4 (aq) + H + (aq) Mn 2+ (aq) + 2CO 2 (g) + H 2 O How to balance using oxidation numbers: 1. Determine oxidation numbers 2. Find a coefficient for each pair that will result in an equal number of electrons transferred. 3. Apply the coefficients to the equation. 2MnO 4 (aq) + 5H 2 C 2 O 4 (aq) + H + (aq) 2 Mn 2+ (aq) + 10CO 2 (g) + H 2 O (notice that we had to use double the coefficients predicted because H 2 C 2 O 4 contains 2 carbons and we don t want to use 5/2 in front of it!) 4. Add H + to balance charge if in acidic media or add OH - to balance charge if in basic media. 2MnO 4 (aq) + 5H 2 C 2 O 4 (aq) + 6 H + (aq) (2 x-1) + (6 x +1) = +4 2 x +2= Balance with H 2 O. 2MnO 4 (aq) + 5H 2 C 2 O 4 (aq) + 6 H + (aq) 2 Mn 2+ (aq) + 10CO 2 (g) + H 2 O 2 Mn 2+ (aq) + 10CO 2 (g) + 8H 2 O 6. Check charge balance and mass balance. 17

18 4.11 Redox Titrations Redox titration - titration used to determine the concentration of many oxidizing or reducing agents Note: i) the substance whose concentration you want to determine must undergo an oxidation or reduction in 100% yield ii) there will usually be a colour change to signal when the reaction is complete There are examples in the text and some will be provided in class Some Applications of Redox Reactions Combustion - the burning of a fuel by oxidation with oxygen in air Bleaching - makes use of redox reaction to decolorize or lighten colored materials Batteries - all are based on redox reactions Metallurgy - the science of extracting and purifying metals from their ores Corrosion - the deterioration of a metal by oxidation Respiration - the process of breathing and using oxygen for many biological redox reactions that provide the energy needed by living organisms 18

19 This document was created with Win2PDF available at The unregistered version of Win2PDF is for evaluation or non-commercial use only.

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