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1 Chapter 4 Aqueous solutions Types of reactions.

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1 1 Chapter 4 Aqueous solutions Types of reactions

2 2 Parts of Solutions l Solution- homogeneous mixture. l Solute- what gets dissolved. l Solvent- what does the dissolving. l Soluble- Can be dissolved. l Miscible- liquids dissolve in each other.

3 3 Aqueous solutions l Dissolved in water. l Water is a good solvent because the molecules are polar. l The oxygen atoms have a partial negative charge. l The hydrogen atoms have a partial positive charge. l The angle is 105ºC.

4 4 Hydration l The process of breaking the ions of salts apart. l Ions have charges and attract the opposite charges on the water molecules.

5 5 Hydration H H O H H O H H O H H O H H O H H O H H O H H O H H O

6 6 Solubility l How much of a substance will dissolve in a given amount of water. l Usually g/100 mL l Varies greatly, but if they do dissolve the ions are separated, l and they can move around. l Water can also dissolve non-ionic compounds if they have polar bonds.

7 7 Types of Solutions l Electricity is moving charges. l The ions that are dissolved can move. l Solutions of ionic compounds can conduct electricity. l Strong electrolytes- completely dissociate (fall apart into ions). l Many ions- Conduct well. l Weak electrolytes- Partially fall apart into ions. l Few ions -Conduct electricity slightly. l Non-electrolytes- Dont fall apart. l No ions- Dont conduct.

8 8 Solubility Rules All nitrates are soluble Alkali metals ions and NH 4 + ions are soluble Halides are soluble except Ag +, Pb +2, and Hg 2 +2 Most sulfates are soluble, except Pb +2, Ba +2, Hg +2,and Ca +2

9 9 Solubility Rules Most hydroxides are slightly soluble (insoluble) except NaOH and KOH Sulfides, carbonates, chromates, and phosphates are insoluble Lower number rules supersede so Na 2 S is soluble

10 10 Measuring Solutions l Concentration- how much is dissolved. l Molarity (M)= Moles of solute Liters of solution l Calculate the molarity of a solution with 34.6 g of NaCl dissolved in 125 mL of solution. l 34.6g/58.44g = mol l mol / L = 4.74 M

11 11 Molarity l How many grams of HCl would be required to make 50.0 mL of a 2.7 M solution? l 2.7 mol = X mol 1000 mL 50 mL X = mol (0.135 mol HCl) x ( g) = 4.9 g HCl 1 mol HCl l What would the concentration be if you used 27g of CaCl 2 to make 500. mL of solution? l What is the concentration of each ion?

12 12 Molarity l Calculate the concentration of a solution made by dissolving 45.6 g of Fe 2 (SO 4 ) 3 to 475 mL. l What is the concentration of each ion?

13 13 Making solutions l Describe how to make mL of a 1.0 M K 2 Cr 2 O 4 solution. l Describe how to make 250. mL of an 2.0 M copper (II) sulfate dihydrate solution.

14 14 Dilution l Adding more solvent to a known solution. l The moles of solute stay the same. l Moles = M x L = M 1 V 1 l Moles = M 2 V 2 l Moles = Moles l M 1 V 1 = M 2 V 2 l Stock solution is a solution of known concentration used to make more dilute solutions

15 15 Dilution l What volume of a 1.7 M solutions is needed to make 250 mL of a 0.50 M solution? l 18.5 mL of 2.3 M HCl is added to 250 mL of water. What is the concentration of the solution? l 18.5 mL of 2.3 M HCl is diluted to 250 mL with water. What is the concentration of the solution?

16 16 Dilution l You have a 4.0 M stock solution. Describe how to make 1.0L of a 0.75 M solution. l Determine the volume needed l This volume is then added to a 1.0 L graduated flask and filled with water to 1.0 liters

17 17 Types of Reactions Precipitation reactions l When aqueous solutions of ionic compounds are poured together a solid forms. l A solid that forms from mixed solutions is a precipitate l If youre not a part of the solution, your part of the precipitate

18 18 Precipitation reactions NaOH(aq) + FeCl 3 (aq) NaCl(aq) + Fe(OH) 3 (s) l is really Na + (aq)+OH - (aq) + Fe +3 + Cl - (aq) Na + (aq) + Cl - (aq) + Fe(OH) 3 (s) l So all that really happens is OH - (aq) + Fe +3 Fe(OH) 3 (s) l Double replacement reaction

19 19 Precipitation reaction l We can predict the products l Can only be certain by experimenting l The anion and cation switch partners AgNO 3 (aq) + KCl(aq) Zn(NO 3 ) 2 (aq) + BaCr 2 O 7 (aq) CdCl 2 (aq) + Na 2 S(aq)

20 20 Precipitations Reactions l Only happen if one of the products is insoluble l Otherwise all the ions stay in solution- nothing has happened. l Need to memorize the rules for solubility (pg 152)

21 21 Three Types of Equations l Molecular Equation- written as whole formulas, not the ions. K 2 CrO 4 (aq) + Ba(NO 3 ) 2 (aq) l Complete Ionic equation show dissolved electrolytes as the ions. 2K + + CrO Ba NO 3 - BaCrO 4 (s) + 2K NO 3 - l Spectator ions are those that dont react.

22 22 Three Types of Equations l Net Ionic equations show only those ions that react, not the spectator ion. Ba +2 + CrO 4 -2 BaCrO 4 (s) l Write the three types of equations for the reactions when these solutions are mixed. l Iron (III) nitrate and potassium hydroxide l Silver nitrate and Potassium Chloride.

23 23 Stoichiometry of Precipitation l What mass of solid is formed when 1.25L of M Pb(NO 3 ) 2 is mixed with 2.00 L of M Na 2 SO 4 ? –Determine what reaction occurs –Write the balanced net ionic equation –Calculate the moles of reactants –Determine which reactant is limiting –Calculate the moles of product –Convert to grams of product

24 24 l 25 mL 0.67 M of H 2 SO 4 is added to 35 mL of 0.40 M CaCl 2. What mass CaSO 4 Is formed? l What volume of M HCl is needed to precipitate the silver from of M silver nitrate solution ?

25 25 Types of Reactions Acid-Base Reactions l Acid: proton donor. l Base: proton acceptor usually OH - l What is the net ionic equation for the reaction of HCl(aq) and KOH(aq)? Acid + Base salt + water H + + OH - H 2 O

26 26 Acid - Base Reactions l Often called a neutralization reaction Because the acid neutralizes the base. l Often titrate to determine concentrations. l Solution of known concentration (titrant), l is added to the unknown (analyte), l until the equivalence point is reached where enough titrant has been added to neutralize it.

27 27 Titration l Endpoint: Where the indicator changes color. l Not always at the equivalence point. l A mL sample of aqueous Ca(OH) 2 requires mL of M Nitric acid for neutralization. What is [Ca(OH) 2 ]? l # of H + x M A x V A = # of OH - x M B x V B

28 28 Types of solutions l Acids- form H + ions when dissolved. l Strong acids fall apart completely (many ions) H 2 SO 4,HNO 3,HCl, HBr, HI, HClO 4 l Weak acids- dont dissociate completely. l Bases - form OH - ions when dissolved. l Strong bases: KOH, NaOH l Strong bases react completely with weak acids l Strong acids react completely with weak bases

29 29 Acid-Base Reaction l 75 mL of 0.25M HCl is mixed with 225 mL of M Ba(OH) 2. What is the concentration of the excess H + or OH - ?

30 30 Types of Reaction Oxidation-Reduction called Redox l Ionic compounds are formed through the transfer of electrons. l An Oxidation-reduction reaction involves the transfer of electrons. l We need a way of keeping track.

31 31 Oxidation States l A way of keeping track of the electrons. l Not necessarily true of what is in nature, but it works. l Rules for assigning (memorize). The oxidation state of elements in their standard states is zero. Oxidation state for monoatomic ions are the same as their charge.

32 32 Oxidation States Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide. In compounds with nonmetals hydrogen is assigned the oxidation state +1. In its compounds fluorine is always –1. The sum of the oxidation states must be zero in compounds or equal the charge of the ion.

33 33 Oxidation States l Assign the oxidation states to each element in the following. l CO 2 l NO 3 - l H 2 SO 4 l Fe 2 O 3 l FeO

34 34 Oxidation-Reduction l Transfer electrons, so the oxidation states change. Na + 2Cl 2 2NaCl CH 4 + 2O 2 CO 2 + 2H 2 O l Oxidation is the loss of electrons. l Reduction is the gain of electrons. l OIL RIG l LEO GER

35 35 Oxidation-Reduction l Oxidation means an increase in oxidation state - lose electrons (becomes more +). l Reduction means a decrease in oxidation state - gain electrons (becomes more -). l The substance that is oxidized is called the reducing agent (more +). l The substance that is reduced is called the oxidizing agent (more -).

36 36

37 37 Identify the l Oxidizing agent l Reducing agent l Substance oxidized l Substance reduced l in the following reactions Fe (s) + O 2 (g) Fe 2 O 3 (s) Fe 2 O 3 (s)+ 3 CO(g) 2 Fe(l) + 3 CO 2 (g) SO H + + MnO 4 - SO H 2 O + Mn +2

38 38 Half-Reactions l All redox reactions can be thought of as happening in two halves. l One produces electrons - Oxidation half. l The other requires electrons - Reduction half. l Write the half reactions for the following and identify as reduction or oxidation. –2Na(aq) + Cl 2 (aq) 2NaCl(s) –Ce 4+ (aq)+ Sn 2+ (aq) Ce 3+ (aq)+ Sn 4+ (aq)

39 39 Balancing Redox Equations l In aqueous solutions the key is the number of electrons produced must be the same as those required. Ce 4+ (aq)+ Sn 2+ (aq) Ce 3+ (aq)+ Sn 4+ (aq) Ce 4+ (aq)+ e - Ce 3+ (aq) Sn 2+ (aq) Sn 4+ (aq) + 2e -

40 40 Acidic Solution – Write separate half reactions – For each half reaction balance all reactants except H and O – Balance O using H 2 O – Balance H using H + – Balance charge using e - – Multiply equations to make electrons equal – Add equations and cancel identical species – Check that charges and elements are balanced.

41 41 Practice l The following reactions occur in aqueous solution. Balance them Cr(OH) 3 + OCl - + OH - CrO Cl - + H 2 O MnO Fe +2 Mn +2 + Fe +3 Cu + NO 3 - Cu +2 + NO(g) Pb + PbO 2 + SO 4 -2 PbSO 4 Mn +2 + NaBiO 3 Bi +3 + MnO 4 -

42 42 Now for a tough one Fe(CN) MnO 4 - Mn +2 + Fe +3 + CO 2 + NO 3 -

43 43 Basic Solution l Do everything you would with acid, but add one more step. l Add enough OH - to both sides to neutralize the H + CrI 3 + Cl 2 CrO IO Cl - Fe(OH) 2 + H 2 O 2 Fe(OH) -

44 44 Redox Titrations l Same as any other titration. l the permanganate ion is used often because it is its own indicator. MnO 4 - is purple, Mn +2 is colorless. When reaction solution remains clear, MnO 4 - is gone. l Chromate ion is also useful, but color change, orangish yellow to green, is harder to detect.

45 45 Example l The iron content of iron ore can be determined by titration with standard KMnO 4 solution. The iron ore is dissolved in excess HCl, and the iron reduced to Fe +2 ions. This solution is then titrated with KMnO 4 solution, producing Fe +3 and Mn +2 ions in acidic solution. If it requires mL of M KMnO 4 to titrate a solution made with g of iron ore, what percent of the ore was iron?

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