Chemistry 40S Aqueous Solutions (This unit has been adapted from

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1 Chemistry 40S Aqueous Solutions (This unit has been adapted from Name: 1

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3 Lesson 1: Aqueous Reactions Goals: Operationally define precipitate. Identify ionic substances that precipitate from aqueous solution using experimentation. Write balanced equations for double displacement and net-ionic precipitation reactions. Water is essential for life on Earth. More than two-thirds of the Earth s surface is water. The cells in our bodies contain between 65% and 90% water by weight. Many of the reactions of life occur in water, we say they are aqueous. Three important types of aqueous reactions are: Aqueous Solutions of Ionic Compounds 3

4 For sodium chloride, we represent the dissolving in the equation NaCl( s) Na + (aq) + Cl (aq) Precipitates A precipitation reaction For example, when a solution of aqueous silver nitrate and aqueous sodium chloride are mixed, the two clear colourless solutions form a cloudy mixture. The cloudy mixture clears as a white solid settles to the bottom of the container. 4

5 Solubility Rules Several ionic substances have a low solubility in water because the force of attraction between the ions is greater than the force water molecules use to dissociate the ions. These compounds do not dissociate in water, rather they will usually form suspensions and settle to the bottom. Solubility Rules for Selected Ionic Compounds in Water at 25 C 1. Compounds of Li +, Na +, K +, Cs +, Fr +, all metals in Main Group 1 of the periodic table, H +, and NH4 + ions are water-soluble. Potassium perchlorate (KClO4) is an exception and is somewhat insoluble. 2. Most compounds of chlorides (Cl - ), bromides (Br - ), and iodides (I - ) are water-soluble. Exceptions are compounds formed between these ions and copper I (Cu + ), silver (Ag + ), lead (Pb 2+ ), and mercury I (Hg2 2+ ), which are insoluble at room temperature. 3. Compounds containing the nitrate (NO3 - ), acetate (C2H3O2 - ), perchlorate (ClO4 - ), and chlorate (ClO3 - ) ions are water-soluble. Potassium perchlorate (KClO4) is an exception and is somewhat insoluble and silver acetate (AgC2H3O2) is considered sparingly soluble. 4. Most compounds of sulfates (SO4 2- ) are water-soluble. Exceptions are compounds formed between sulfates and calcium (Ca 2+ ), strontium (Sr 2+ ), barium (Ba 2+ ), and lead (Pb 2+ ) ions, which are water-insoluble. Silver sulfate (Ag2SO4) is considered sparingly soluble. 5

6 5. Most compounds of sulfides (S 2- ) are water-insoluble. Exceptions are the sulfides of Main Group 1 and Main Group 2 elements and of the ammonium ion (NH4 + ). 6. Most compounds containing the hydroxide ion (OH - ) are waterinsoluble. Exceptions are the hydroxides of metals from Main Group 1 of the periodic table, and of compounds formed between hydroxide ion and H +, Sr 2+, Ba 2+, and NH Most compounds containing the carbonate (CO3 2- ), phosphate (PO4 3- ), and sulfite (SO3 2- ) ions are water-insoluble. Exceptions are the compounds formed with H +, NH4 +, and with metals from Main Group 1 of the periodic table. Example 1: Use the solubility rules to predict if the following compounds are soluble. a) Sodium sulphate (Na2SO4) b) Silver nitrate (AgNO3) c) Aluminum hydroxide (Al(OH)3) d) Lead (II) chloride (PbCl2) 6

7 Precipitation Reactions AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) 7

8 Example 2: Write the complete set of equations (molecular, ionic and net ionic) for the reaction between aqueous lead (II) nitrate and aqueous potassium iodide. Step 1: Write the formulas of the initial compounds. Step 2: Identify the ions present. Step 3: Write the formulas of the 2 possible compounds formed from the ions. Step 4: Determine which, if any, of these compounds would form a precipitate. 8

9 Example 3: Write balanced precipitation, complete ionic and net ionic equations for the mixing of the following solutions. If no reaction occurs, write no reaction. Show each step. a) Ba(NO3)2 (aq) and Na2CO3 (aq) b) Pb(NO3)2 (aq) and NH4I (aq) c) CuSO4 (aq) and Na3PO4 (aq) d) CoCl2(aq) and KOH(aq) 9

10 Practice: Aqueous Reactions 1. Using the Ion Solubility Chart, identify each of the following compounds as soluble or insoluble, low solubility. a) Tin (III) sulphate b) Nickel (II) iodide c) Lithium carbonate d) Barium phosphate e) Silver iodide 2. Write balanced precipitation, complete ionic and net ionic equations for the mixing of the following solutions. If no reaction occurs, write no reaction. Show each step. a) Ammonium sulfate and rubidium carbonate 10

11 b) Sodium hydroxide and nickel (II) chloride c) Strontium hydroxide and calcium iodide d) Ammonium phosphate and barium chloride e) Aluminum nitrate and magnesium sulfate 11

12 Lesson 2: Neutralization Reactions Goals: Describe the uses of common acids and bases in daily experience. Operationally define acids and bases in terms of observable and measurable properties. Define acids and bases in terms of Arrhenius' theory. Write formulas from memory for common acids and bases. Write balanced equations for neutralization reactions. Compare the neutralization of polyprotic acids and polyhydroxic bases. Defining Acids and Bases Example 1: Hydrochloric acid: HCl (aq) H + (aq) + Cl - (aq) Example 2: Acetic acid (vinegar): HC2H3O2 (aq) H + (aq) + C2H3O2 - (aq) Example 3: Sodium hydroxide: NaOH(s) Na + (aq) + OH - (aq) 12

13 Example 4: Aqueous ammonia: NH3 (g) + H2O (l) NH4OH (aq) NH4 + (aq) + OH - (aq) Ammonia is a little troublesome. It must first be shown to react with water then dissociate. Aqueous ammonia is often called ammonium hydroxide. Properties of Acids Acidic solutions in water have the following properties: Here is a group of common acids. To help you remember the names and formulas, here are a few general naming rules for acids: Compounds ending in ide become hydro...ic acids. For example, hydrogen chloride (HCl) becomes hydrochloric acid. Compounds ending in ate become ic acids. For example, hydrogen sulfate (H2SO4) becomes sulfuric acid. Compounds ending in ite become ous acids. For example, hydrogen sulfite (H2SO3) becomes sulfurous acid. 13

14 Acid Formula Common Uses or Commonly Found Hydrochloric HCl Commonly known as muriatic acid or stomach acid. Produced in and secreted by the stomach, and used commonly in laboratories. It is also used in food processing, as a cleaning agent, in the activation of oil wells, in the production of other chemicals, and in recovering magnesium from seawater. Sulfuric H2SO4 Battery acid. The sulfur compounds, released by onions when cut, form sulphuric acid when dissolved in water causing teary eyes. Used in fertilizer manufacturing, petroleum refining, the production of metals, paper, paint, dyes, detergents, and many chemical raw materials. Pure sulfuric acid is a dense, oily liquid that has a high boiling point. When you combine water with concentrated sulfuric acid, a large amount of heat is released. Concentrated sulfuric acid is a good dehydration agent. Nitric HNO3 Pure nitric acid is an unstable liquid. Concentrated nitric acid is more stable and is 70% by mass of the acid dissolved in water. It is used in making rubber, chemicals, dyes, plastics, drugs, and explosives. Acetic HC2H3O2 Also known as vinegar. Carbonic H2CO3 CO2 dissolved in water, as in carbonated beverages, forms carbonic acid. Nitrous HNO2 One component of acid rain. Used in the manufacture of fertilizers and in organic reactions. Sulfurous H2SO3 A component of acid rain. Used to manufacture sulfuric acid and in the pulp and paper industry. Phosphoric H3PO4 In carbonated beverages. Dilute phosphoric acid is not toxic and has a sour taste. It is used as a flavouring agent in beverages and is used as a cleaning agent for dairy equipment. It is also used in the manufacture of detergents, ceramics, and phosphorous-containing chemicals. Hypochlorous HClO When chlorine gas is dissolved in water hypochlorous acid is a product. Hypochlorous acid is used as a disinfectant in water supplies and swimming pools. Also used in the manufacture of bleaches. 14

15 Acids such as hydrochloric, nitric, and acetic are called monoprotic acid. Acids that release more than one hydrogen ion in water are called polyprotic acid. Properties of Bases Basic solutions in water have the following properties: 15

16 Here is a list of common bases: Base Formula Common Uses Sodium hydroxide NaOH Also called lye. Used in oven cleaners and drain cleaners. It is also used in the manufacturing of soap. Potassium hydroxide KOH Also known as potash. Used for soap making, drain cleaners and making alkaline batteries. Magnesium hydroxide Mg(OH)2 Milk of magnesia, antacids, laxatives. Calcium hydroxide Ca(OH)2 Also called slaked lime or hydrated lime. Used in antacids, making plaster and concrete, removing hair from leather hides, and steel making. Ammonia NH3 Used in window and general household cleaners, manufacturing of fertilizers and explosives. Aluminum hydroxide Al(OH)3 Used in antacids, antiperspirants, and ceramics. Bases such as sodium hydroxide and potassium hydroxide are called monohydroxic. Bases that release more than one hydroxide ion in water are called polyhydroxic. 16

17 Neutralization Reactions In general, the reaction for a neutralization reaction is given by acid + base salt + water For example, when hydrochloric acid and sodium hydroxide are mixed the following reaction occurs: HCl (aq) + NaOH (aq) NaCl (aq) + HOH (l) The hydrogen from the acid combines with the hydroxide from the base to form HOH, or water. The remaining metal and non-metal ions combine to form the aqueous salt. Neutralization Equations HCl (aq) + NaOH (aq) NaCl (aq) + HOH (l) The reaction between the hydrogen and hydroxide ions is seen more easily in the ionic equation. The aqueous sodium ions and chloride ions are spectators in this reaction, since they start as separate aqueous ions and end as the same. The net ionic equation shows the reaction that occurs, omitting the spectator ions. The net ionic equation for this reaction omits the sodium and chloride ions from the ionic equation. H + (aq) + OH - (aq) H2O (l) The net result of a neutralization reaction is the reacting of the hydrogen ion from an acid with the hydroxide of a base, to form water. 17

18 Example 5: Write the complete set of equations, molecular, ionic and net ionic, showing the reaction between potassium hydroxide and sulphuric acid. 18

19 Example 6: Write the complete set of reactions that occur when the following acid and bases are reacted. a) Acetic acid and sodium hydroxide b) Sulphuric acid and potassium hydroxide 19

20 Practice: Neutralization Reactions Write the complete set of reactions that occur when the following acid and bases are reacted. 1. Nitric acid and calcium hydroxide 2. Phosphoric acid and lithium hydroxide 20

21 3. Sulphuric acid and aluminum hydroxide 4. Sulfurous acid and magnesium hydroxide 21

22 5. Nitrous acid and barium hydroxide 6. Hydrochloric acid and magnesium hydroxide 22

23 Lesson 3: The Stoichiometry of Neutralization Goals: Calculate the concentration or volume of an unknown acid or base from the concentration and volume of a known acid and base required for neutralization. Compare the neutralization of polyprotic acids and polyhydroxic bases. Defining Neutralization Steps for solving neutralization problems: Step 1: Write the balance neutralization equation. Step 2: Substitute given values into the equation above and solve for the required amount, or find moles of given amount, then use stoichiometry to find the moles, then required amount of unknown. 23

24 Example 1: Calculate the concentration of hydrochloric acid, if 25.0 ml is just neutralized by 40.0 ml of a mol/l sodium hydroxide solution. 24

25 Example 2: What volume of 1.00 mol/l potassium hydroxide is needed to neutralize 750. ml of a mol/l nitric acid solution? 25

26 Neutralization With Polyhydroxic Bases and Polyprotic Acids Polyhydroxic Bases Polyprotic Acids Example 3: What is the concentration of a sample of magnesium hydroxide, if 225 ml of the base is neutralized by 125 ml of a mol/l hydrochloric acid solution? 26

27 Example 4: Calculate the volume of mol/l carbonic acid needed to neutralize 25.0 ml of mol/l sodium hydroxide. 27

28 Practice: The Stoichiometry of Neutralization 1. What volume of mol/l sulfuric acid can be neutralized by 50.0 ml of mol/l sodium hydroxide? 2. Concentrated hydrochloric acid (11.7 mol/l) is added to a spill of 5.00 L of sodium hydroxide solution with a concentration of 2.00 mol/l. What volume of acid is required to neutralize the spill? 3. A clumsy chemistry teacher (not this one, of course) spills 75.0 ml of concentrated sulfuric acid (18.0 mol/l). He has a stock solution of 1.00 mol/l sodium hydroxide on hand to neutralize the spill. How much base does he need to neutralize the spill? 28

29 4. If ml of an antacid containing magnesium hydroxide is completely neutralizes L of mol/l hydrochloric acid solution, what is the concentration of the antacid? 5. What is the concentration of a solution of phosphoric acid if 325 ml is required to neutralize 250 ml of a mol/l solution of potassium hydroxide? 6. Calculate the concentration of nitrous acid if 25.0 ml of the acid is needed to neutralize 19.0 ml of mol/l potassium hydroxide. 29

30 Lesson 4: Oxidation and Reduction Reactions Goals: Relate the role of oxygen to the process of rusting and burning. Define oxidation and reduction. Determine oxidation numbers for atoms in simple compounds and ions. Identify the oxidizing agent, the reducing agent, the substance reduced and the substance oxidized, given a balanced chemical equation. Defining Oxidation and Reduction Oxidation-reduction reactions For example, a reaction of the burning of magnesium: 2Mg (s) + O2 (g) 2MgO (s) In this reaction magnesium begins as a neutral atom and loses two electrons to become a Mg 2+ ion in MgO. Magnesium is oxidized. Oxygen begins as a neutral atom and gains the two electrons from magnesium to become an O 2 ion in MgO. Oxygen is reduced. 30

31 Oxidation Numbers Not all redox reactions are as simple as the ones on the previous page, so chemists have designed a system to keep track of electrons in a chemical reaction. They have assigned oxidation numbers to all atoms and ions. The oxidation number represents the charge the atom would have if every bond were ionic. Not every bond is ionic, but chemists assume they are for this system. The oxidation number is not always the actual charge, but it is very helpful to follow electrons in redox reactions. The rules for assigning oxidation numbers are as follows: Oxidation numbers are always assigned PER ATOM. The oxidation numbers of all uncombined elements is zero. For example: O2, H2, etc The oxidation number of monatomic ions equals the charge of that ion. In compounds, the oxidation number for alkali metals. For example, Li, Na, K, etc is always +1. In compounds, the oxidation number for the alkaline earth metals. For example, Be, Mg, Ca, etc is always +2. In compounds, the oxidation number of aluminum is always +3. In compounds, the oxidation number of fluorine is -1. In compounds, the oxidation number of hydrogen is +1. An exception is in metal hydrides, such as NaH or MgH2, when hydrogen is -1. In compounds, the oxidation number of oxygen is -2. An exception is in peroxides, such as H2O2 or Na2O2, when its oxidation number is -1. For any neutral compound, the sum of the oxidation numbers for each atom must be zero. For a polyatomic ion, the sum of the oxidation numbers for each atom must be the charge of that ion. There is a difference between ion charge and oxidation numbers. For example, the magnesium ion, Mg 2+, has an ion charge of 2 + and an oxidation number of

32 Example 1: Assign oxidation numbers to each atom in SO2. Example 2: Assign oxidation numbers for each atom in K2Cr2O7. Example 3: Assign oxidation numbers for each atom in Fe(NO3)3. 32

33 Example 4: Is the reaction SO2 + H2O H2SO3 a redox reaction? Example 5: Is the reaction Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s) a redox reaction? 33

34 Practice: Oxidation and Reduction Reactions 1. Assign oxidation numbers for each element in the following compounds. a) Na2Cr2O7 b) KNO3 c) FeCl2 d) H2C2O4 e) HClO3 f) KMnO4 g) MnSO4 h) H2SO4 i) H2 k) NH4NO3 l) NO2 l) C2H5OH 2. Identify which of the following are redox reactions. a) 2NO2 N2O4 b) 2 Mg + O2 2MgO c) Mg + + 2Ag + + 2NO3 - Mg NO Ag 34

35 Lesson 5: Oxidizing and Reducing Agents Goals: Identify the substance oxidized and the substance reduced in a redox reaction. Identify the oxidizing and reducing agents in a redox reaction. Determine the number of electrons transferred in a redox reaction. Defining Oxidizing and Reducing Agents 35

36 Example 1: Nitric acid reacts with hydrogen sulfide according to the balanced equation below. Identify the substance oxidized, the substance reduced, the oxidizing agent and the reducing agent for the burning of propane. 2HNO3 (aq) + 3H2S (g) 2NO (g) + 3S (s) + 4H2O (l) 36

37 Electrons Transferred Redox reaction is defined as a chemical reaction in which electrons are transferred from one atom to another. We can use a balanced equation to determine how many electrons have been transferred. To determine the number of electrons transferred in a redox reaction Step 1: Assign oxidation numbers Step 2: Determine atoms gaining and losing reactions Step 3: Multiply the number electrons lost or gained by the number of atoms gaining or losing the electrons Example 2: For the redox reaction Cu + 2AgNO3 Cu(NO3)2 + 2Ag. How many electrons are transferred? 37

38 Example 3: How many electrons are transferred in the reaction below? 2HNO3 (aq) + 3H2S (g) 2NO (g) + 3S (s) + 4H2O (l) 38

39 Practice: Oxidizing and Reducing Agents Which of the following equations represents a redox reaction? For the redox reactions, identify: element oxidized element reduced oxidizing agent reducing agent number of electrons transferred 1. H2 + I2 2HI 2. 2Ag + Mg(NO3)2 2AgNO3 + Mg 39

40 3. I2 + 2Na2S2O3 Na2S4O6 + 2NaI 4. MgO + SO3 MgSO4 5. 2KMnO4 + 8H2SO4 + 10KBr 5Br2 + 6K2SO4 + 2MnSO4 + 8H2O 40

41 6. 3Ag2S + 8HNO3 6AgNO3 + 2NO + 3S + 4H2O 7. 2C2H6 + 3O2 4CO2 + 6H2O 41

42 Lesson 6: Balancing Equations Using Redox Methods Goals: Balance redox equations using the oxidation number method. Balance redox equations in acidic and basic solutions using the half reaction method. The Oxidation Method Example 1: Use the oxidation method to balance the equation below. K2Cr2O7 + H2O + S SO2 + KOH + Cr2O3 42

43 Example 2: Balance the following equation using the oxidation method. P + HNO3 + H2O NO + H3PO4 43

44 Why We Balance Electrons Lost and Gained First? Example 3: Balance the reaction of Cu + Cl2 Cu + + Cl -. Half-Reaction 44

45 Balancing Redox Reactions by Half Reaction Method Steps for Balancing Redox Reactions in Acidic Solutions: There are about 8 steps to this method. Eventually, you may wish to combine some of these steps to speed up the balancing process, but never skip steps. The order of the steps is also important. Step 1: Assign oxidation numbers. Step 2: Identify the substances oxidized and reduce then write the oxidation and reduction half-reactions. Step 3: Balance all elements except hydrogen and oxygen. Step 4: Add electrons lost and gained. Step 5: Balance oxygen atoms by using H2O. Step 6: Balance hydrogen atoms using H + ions. Step 7: Balance the number of electrons lost and gained. Step 8: Add the two half-reactions. 45

46 Example 4: Balance the following reaction in an acidic solution. Cr2O7 2- (aq) + SO3 2- (aq) Cr 3+ (aq) + SO4 2- (aq) 46

47 Example 5: Balance the following equation in an acidic solution. MnO4 - + I - MnO2 + I2 47

48 Steps for Balancing Redox Reactions in Acidic Solutions: The first part of balancing redox reactions in basic solutions follows the same steps as that for acidic solutions. Step 1: Assign oxidation numbers. Step 2: Identify the substances oxidized and reduce then write the oxidation and reduction half-reactions. Step 3: Balance all elements except hydrogen and oxygen. Step 4: Add electrons lost and gained. Step 5: Balance oxygen atoms by using H2O. Step 6: Balance hydrogen atoms using H + ions. Since the reaction occurs in a base, we must add hydroxide ions. Step 7: Before adding the two half-reactions, you add the same number of hydroxide ions as hydrogen ions to BOTH sides of the equation. Step 8: We eliminate the hydrogen ions by forming water (H + + OH - H2O). Then continue as we did with acid. Step 9: Balance the number of electrons lost and gained. Step 10: Add the two half-reactions. 48

49 Example 6: Balance the following reaction in a basic solution. MnO4 - + C2O4 2- CO2 + MnO2 49

50 Example 7: Balance the following reaction in a basic solution. N2O + ClO - NO2 - + Cl - 50

51 Practice: Balancing Equations Using Redox Methods Complete the questions on loose-leaf. 1. Balance the following using the Oxidation Number Method. a) HNO3 + C2H6O + K2Cr2O7 KNO3 + C2H4O + H2O + Cr(NO3)3 b) Al + CuSO4 Al2(SO4)3 + Cu c) Cu + HNO3 Cu(NO3)2 + NO2 + H2O d) KMnO4 + KCl + H2SO4 MnSO4 + K2SO4 + Cl2 + H2O 2. Use the Half-Reaction Method to balance the following redox reactions in acidic solutions. a) S 2- + Cr2O7 2- S + Cr 3+ b) Br - + SO4 2- Br2 + SO2 c) Zn + NO3 - Zn 2+ + N2O d) C2H5OH + NO3 - HC2H3O2 + NO2 3. Use the Half-Reaction Method to balance the following redox reactions in a basic solution. a) AsO NO2 - AsO2 - + NO3 - b) ClO3 - + N2H4 NO + Cl - c) Zn + BrO4 - Zn(OH) Br - d) ClO3 - + BH4 - Cl - + H2BO3-51

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