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Chapter 5

Lecture and Textbook Collaborated Notes - Chapter 5 - CHEM 1A03

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Aadil Merali Juma

Chem 1A03 Chapter 5: Introduction to Reactions in Aqueous Solutions Water: A Vital Natural Resource  Significance of Water – Energy, Health and Environment  “No single measure would od more to reduce disease and save lives in the developing world than bringing safe water and adequate sanitation to all” – UN Secretary-General Kofi Annan, 2003 international Year of Freshwater  “Water is probably the only natural resource to touch all aspects of human civilization – from agricultural and industrial development to the cultural and religious values embedded in society” – Koichiro Matsuura, Director-General, UNESCO Water: An Enigmatic Medium  Unique Macroscopic Properties Essential to Life 1. Boiling Point – unusually high b.p. of liquid water at STP: o Bp = 100°C at 1atm (Sea Level)  TP dependent o Due to H-bonding 2. Density – of ice (0.92g/ml) < liquid water (1g/mL) at 0°C o Ice expands with freezing: Lower packing density (ice floats) 3. Specific Heat Capacity – high (74 J/molK at STP) o Absorbs thermal heat for storage/distribution/release 4. Ability to dissolve a variety of solutes as a solvent o Strong solvation properties for most polar molecules/ions Solute Solvation by Water  Solubilization of polar/ionic solutes: Hydration  Energetically favorable2H O-solute interactions o H-bonding (dipole-dipole) with solutes (O-H, N-H, C=O) o +/-Ion-dipole forces Properties of Water Vital for Molecule Behavior  Ion Hydration, Sodium Pumps and Ion Transport  Protein Folding, Activity and Drug Binding  Solubility, Toxicity and Bioaccumulation Autoionization of Water  Water auto-ionization: Water is amphoteric and conductive Chem 1A03 5.1 The Nature of Aqueous  Aqueous solution containing ions conduct electricity o Ions conduct electricity because the ions move essentially independently of each other, each one carrying a certain quantity of charge  Nonelectrolyes – solutes that do not provide ions when dissolved in water o Eg/ Lamp fails to light up  Electrolytes – soutes that provide ions when dssolved in water  Strong Electrolyte – a substance that is essentially completely ionized in aqueous solution o Has a strong tendency for providing ions o Eg/ Light lights up brightly  Weak Electrolyte – a substance that is only partially ionized in aqeous solution o Has a weak (or small) tendency for providing ions o Eg/ Light lights up dimly Solute Properties and Electrolytes  Degree of ionization of an electrolyte in aqueous solution  Essentially all soluble ionic compounds and only a few relatively few molecular compounds are strong electrolytes  Most molecular compounds are either nonelectrolytes or weak electrolytes Solutes as Electrolytes 1. Strong Electrolyte (Ionic) o Sodium Chloride: NaCl (aq) Na +(aq) Cl-(aq) o Intravenous supply of electrolytes o HCl (aq) H +(aq) Cl-(aq) o H+, bare proton, interacts with water molecules surrounding it - hydration 2. Weak Electrolyte (Weakly Ionic) o Acetic Acid: Vinegar (5-8%w): CH CO3H (aq) H 2  CH C3O (aq)H3O +(aq) pKa = 4.8Conductive Ions o Preservative/Sour 3. Non-electrolyte (Neutral) o Ethanol: Vodka (>40%v): CH CH 3H 2 pK a16  no ionization 5.2 Precipitation Reactions Major Types of Aqueous Chemical Reactions 1. Solubility or Precipitation Reactions (ion transfer) Cu(NO )3 2(aq) 2 NaOH (aq) Cu(OH) 2(s) 2 NaNO 3(aq) 2. Reduction-Oxidation (Redox) Reactions (electron transfer) 2+ 2+ Zn (s)u (aq) Zn (aq)+ Cu(s) 3. Acid-Base Reactions (proton transfer) - + CH 3OOH (aq) H 2 (l)H COO3 (aq) H 3 (aq) Chem 1A03  K is a measure of extent of chemical reaction and transformation Net Ionic Reaction – an equation that includes only the actually participants in a reaction  Precipitation of AgI (s) Chemical Equation AgNO 3(aq)+ KI(aq) AgI +(s)O 3(aq) Total Ionic Equation Ag (aq)+NO 3 (aq) K +(aq) I-(aq) AgI +(s) (aq)+ NO 3 (aq) Spectator Ions Derived from Strong Electrolytes Net Ionic Equation Ag (aq)+ I(aq) AgI (s) Predicting Precipitation Reactions - Solubility: Qualitative Guidelines  Solubility Guidelines for Common Ionic Solids (Table 5.1, page 158) o Follow the lower-numbered guidelines when two guidelines are in conflict + + 1. Salts of group 1 cations (with some exceptions for Li ) and the NH cation a4e soluble 2. Nitrates, acetates and perchlorates are soluble 3. Salts of silver, lead and mercury (I) are insoluble 4. Chlorides, bromides and iodides are soluble 5. Carbonates, phosphates, sulfides, oxides and hydroxides are insoluble (sulfides of group 2 cations and hydroxides of Ca , Sr and2+ Ba2+are slightly soluble 6. Sulfates are soluble except for those of calcium, strontium and barium  Soluble: Strong Electrolytes (K >> sp o KI (aq) K +(aq)+ I(aq)  Insoluble: Weak Electrolytes (K <<1) sp o CdCO 3(s)D Cd 2+(aq) CO 32(aq) Solubility and Precipitation Demo #1  Solubility Produce (defined by K ) sp o AgI D Ag + + I- K = 8.3 x 10 -17 (s) (aq) (aq) sp  Precipitation + - 16 o Ag (aq) I (aq) AgI (s) K=1/K = sp2 x 10 Demo#2 + - -9  PbI 2(s) Pb 2 (aq) 2I (aq) Ksp7.1 x 10 Solubility Product for Salts: K sp  Measure of the
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