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Ch. 4 Types of Chemical Reactions and Solution Stoichiometry

Ch. 4 Types of Chemical Reactions and Solution Stoichiometry

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Ch. 4 Types of Chemical Reactions and Solution Stoichiometry

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  1. Ch. 4Types of Chemical Reactionsand Solution Stoichiometry

  2. Solute A solute is the dissolved substance in a solution. Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water in soda

  3. Saturation of Solutions • A solution that contains the maximum amount of solute that may be dissolved under existing conditions is • saturated. • A solution that contains less solute than a saturated solution under existing conditions is • unsaturated. • A solution that contains more dissolved solute than a saturated solution under the same conditions is • supersaturated.

  4. Definition of Electrolytes and Nonelectrolytes An electrolyte is: • A substance whose aqueous solution conducts • an electric current. A nonelectrolyte is: • A substance whose aqueous solution does not • conduct an electric current. Try to classify the following substances as electrolytes or nonelectrolytes…

  5. Electrolytes? • Pure water • Tap water • Sugar solution • Sodium chloride solution • Hydrochloric acid solution • Ethyl alcohol solution • Pure, solid sodium chloride

  6. Answers… ELECTROLYTES: NONELECTROLYTES: • Tap water (weak) • NaCl solution • HCl solution • Pure water • Sugar solution • Ethanol solution • Pure, solid NaCl But why do some compounds conduct electricity in solution while others do not…?

  7. Ionic CompoundsDissociate NaCl(s)  Na+(aq) + Cl-(aq) AgNO3(s)  Ag+(aq) + NO3-(aq) MgCl2(s)  Mg2+(aq) + 2 Cl-(aq) Na2SO4(s)  2 Na+(aq) + SO42-(aq) AlCl3(s)  Al3+(aq) + 3 Cl-(aq)

  8. Ions tend to stay in solution where they canconduct a current rather than re-forming a solid. The reason for this is the polar nature of the water molecule… Positive ions associate with the negative end of the water dipole (oxygen). Negative ions associate with the positive end of the water dipole (hydrogen).

  9. Some covalent compounds IONIZE in solution Covalent acids form ions in solution, with the help of the water molecules. For instance, hydrogen chloride molecules, which are polar, give up their hydrogens to water, forming chloride ions (Cl-) and hydronium ions (H3O+).

  10. Strong acids such as HCl are completely100% ionized in solution. Other examples of strong acids include: • Sulfuric acid, H2SO4 • Nitric acid, HNO3 • Hydriodic acid, HI • Perchloric acid, HClO4

  11. Weak acids such as lactic acid usually ionize less than 5% of the time. Many of these weaker acids are “organic” acids that contain a “carboxyl” group. The carboxyl group does not easily give up its hydrogen.

  12. Because of the carboxyl group, organic acids aresometimes called “carboxylic acids”. Other organic acids and their sources include: • Citric acid – citrus fruit • Malic acid – apples • Butyric acid – rancid butter • Amino acids – protein • Nucleic acids – DNA and RNA • Ascorbic acid – Vitamin C This is an enormous group of compounds; these are only a few examples.

  13. However, most covalent compounds do not ionizeat all in solution. Sugar (sucrose – C12H22O11), and ethanol (ethyl alcohol – C2H5OH) do not ionize - That is why they are nonelectrolytes!

  14. Molarity The concentration of a solution measured in moles of solute per liter of solution. M = mol L

  15. Preparation of Molar Solutions Problem: How many grams of sodium chloride are needed to prepare 1.50 liters of 0.500 M NaCl solution? • Step #1: Ask “How Much?” (What volume to prepare?) • Step #2: Ask “How Strong?” (What molarity?) • Step #3: Ask “What does it weigh?” (Molar mass is?) 1.500 L 0.500 mol 58.44 g = 43.8 g 1 L 1 mol

  16. It is not practical to keep solutions of many different concentrations on hand, so chemists prepare more dilute solutions from a more concentrated “stock” solution. Serial Dilution Problem: What volume of stock (11.6 M) hydrochloric acid is needed to prepare 250. mL of 3.0 M HCl solution? MstockVstock = MdiluteVdilute (11.6 M)(x Liters) = (3.0 M)(0.250 Liters) x Liters = (3.0 M)(0.250 Liters) 11.6 M = 0.065 L

  17. A. Single Replacement Reactions A + BX  AX + B BX + Y  BY + X Replacement of: • Metals by another metal • Hydrogen in an acid by a metal • Hydrogen in water by a metal • Halogens by more active halogens • Ex. Mg(s) + HCl(aq) MgCl2(aq) + H2(g) • Ex. 2 Li(s) + 2 H2O(l) 2 LiOH(aq) + H2(g)

  18. The Activity Series of the Metals • Lithium • Potassium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • Hydrogen • Bismuth • Copper • Mercury • Silver • Platinum • Gold Metals can replace other metals provided that they are above the metal that they are trying to replace. Metals above hydrogen can replace hydrogen in acids. Metals from sodium upward can replace hydrogen in water

  19. The Activity Series of the Halogens • Fluorine • Chlorine • Bromine • Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl(s) + F2(g)  2NaF(s) + Cl2(g) ??? MgCl2(s) + Br2(g)  No Reaction ???

  20. Practice problems - Answers are unbalanced!  1. Mg + FeCl3 Fe + MgCl2 2. Sodium is added to water. Na + H2O  H2 + NaOH 3. Lithium is added to hydrochloric acid Li + HCl  H2 + LiCl 4. Zinc is added to a solution of sodium chloride Zn + NaCl  N.R. 5. Chlorine gas is bubbled into a solution of potassium iodide Cl2 + KI  I2 + KCl 6. Chlorine gas is bubbled into a solution of potassium fluoride Cl2 + KF  N. R.

  21. Double Replacement Reactions The ions of two compounds exchange placesin an aqueous solution to form two new compounds. AX + BY  AY + BX One of the compounds formed is usually a precipitate(an insoluble solid), an insoluble gasthat bubbles out of solution, or a molecular compound, usually water.

  22. Double replacement forming a precipitate… Double replacement (ionic) equation Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq) Complete ionic equation shows compounds as aqueous ions Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) +2 I-(aq)  PbI2(s) + 2K+(aq) + 2 NO3-(aq) Net ionic equation eliminates the spectator ions Pb2+(aq) + 2 I-(aq)  PbI2(s)

  23. Solubility Rules – Mostly Soluble

  24. Solubility Rules – Mostly Insoluble

  25. D.R. Practice problems 1. KBr(aq) + AgNO3(aq)  AgBr(s) + KNO3(aq) 2. Silver nitrate + potassium chromate  2AgNO3(aq) + K2CrO4(aq)  AgCrO4(s) + 2KNO3(aq) 3. Ammonium chloride + cobalt (II) sulfate  2NH4Cl(aq) + CoSO4(aq)  (NH4)2SO4(aq) + CoCl2(aq) N.R. 4. Lithium hydroxide + sodium chromate 2LiOH(aq) + Na2CrO4(aq)  2NaOH(aq) + Li2CrO4(s) 5. Zinc acetate + cesium hydroxide  Zn(C2H3O2)2(aq) + 2CsOH(aq)  Zn(OH)2(s) + 2CsC2H3O2(aq) 6. What is the net ionic equation for the rxn above? Zn+2(aq) + OH-(aq)  Zn(OH)2(s)

  26. Unstable Compounds!!! (own note paper) • Ammonium hydroxide • NH4OH • Carbonic Acid • H2CO3 • Sulfurous acid • H2SO3 • Sulfide salts (ex. Na2S) from acid (H+) • All break down to form other products! • NH4OH  NH3(g) + H2O(l) • H2CO3  CO2(g) + H2O(l) • H2SO3  SO2(g) + H2O(l) • S-2  H2S(g)

  27. 1. Sodium sulfite + hydrochloric acid Unstable Examples: Na2SO3(aq) + 2HCl(aq)  H2SO3(aq) + 2NaCl(aq) H2SO3(aq)  H2O(l) + SO2(g) Na2SO3(aq) + 2HCl(aq)  H2O(l) + SO2(g) + 2NaCl(aq) 2. Ammonium sulfate + sodium hydroxide (NH4)2SO4(aq) + NaOH(aq)  2 NH4OH(aq) + Na2SO4(aq) 2 NH4OH(aq)  2 NH3(g) + H2O(l) (NH4)2SO4(aq)+ 2NaOH(aq)  2NH3(g) + 2H2O(l) + Na2SO4(aq) What is the net ionic equation for the reaction above? NH4+(aq) + OH-(aq)  NH3(g) + H2O(l)