1 / 50

Chapter 4: Solution Chemistry and the Hydrosphere

Chapter 4: Solution Chemistry and the Hydrosphere. Problems: 4.1-4.80, 4.85-4.96, 4.99-4.100, 4.111-4.113, 4.119-4.120, 4.129, 4.131-4.132. Solutions on Earth and Other Places. aqueous solution: a solution where water is the dissolving medium (the solvent )

donagh
Télécharger la présentation

Chapter 4: Solution Chemistry and the Hydrosphere

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 4:Solution Chemistry and the Hydrosphere Problems: 4.1-4.80, 4.85-4.96, 4.99-4.100, 4.111-4.113, 4.119-4.120, 4.129, 4.131-4.132

  2. Solutions on Earth and Other Places aqueous solution: asolution where water is the dissolving medium (the solvent) • For example, when table salt (NaCl) is dissolved in water, it results in an aqueous solution of sodium chloride, NaCl(aq), with Na+ and Cl- ions dissolved in water. • Note: The physical state aqueous,(aq), indicates an element or compound dissolved in waterwhile the physical state liquid,(l), means a pure substance in the liquid state. • Thus, NaCl(aq)NaCl(l), which is molten NaCl requiring very high temperatures. A solution consists of a solute dissolved in a solvent.

  3. Solutions solute: component present in smaller amount solvent: component present in greater amount The formation of a solution: As a solute crystal is dropped into a solution, the water molecules begin to pull apart the ionic compound ion by ion • Solvent molecules surround the solute particles, forming a solvent cage • the ions are now hydrated (surrounded by polar water molecules) • solute is now dissolved in the solvent and cannot be seen because the ions are far apart, like the particles in a gas

  4. Unsaturated, Saturated and Supersaturated Solutions • In general, if a solid is soluble in a solvent, more solid dissolves in the solvent at higher temperatures. unsaturated: contains less than the maximum amount of solute that a solvent can hold at a specific temperature saturated: contains the maximum amount of solute that a solvent can hold at a specific temperature supersaturated:contains more than the maximum amount of solute that a solvent should be able to hold at specific temperature

  5. Unsaturated, Saturated and Supersaturated Solutions • A supersaturated NaC2H3O2 solution recrystallizing after addition of more solute:

  6. Unsaturated, Saturated and Supersaturated Solutions • How can a solution hold more solute than it should be able to hold? • If a given amount of solute is dissolved in a solvent at a higher temperature, and the solution is allowed to cool without being disturbed, the solute will remain in solution. • But the solution is unstable, and the solute will come out of solution (i.e. recrystallize) if the solution is disturbed (e.g. by adding more solute, scratching the glass, etc.)

  7. Unsaturated, Saturated and Supersaturated Solutions For some substances, recrystallization is exothermic, releasing heat to the surroundings. • Hot packs used to warm hands and feet in winter (though some of these are oxidation reactions, which we will discuss later) For other substances, recrystallization is endothermic, absorbing heat and making the surroundings colder. • Cold packs used for sports injuries

  8. How do we measure concentration? solution: homogeneous mixture of substances present as atoms, ions, and/or molecules solute: component present in smaller amount solvent: component present in greater amount Note: Unless otherwise stated, the solvent for most solutions considered in this class will almost always be water! Aqueous solutions are solutions in which water is the solvent.

  9. How do we measure concentration? • A concentrated solution has a large quantity of solute present for a given amount of solution. • A dilute solution has a small quantity of solute present for a given amount of solution. SOLUTION CONCENTRATION= The more solute in a given amount of solution  the more concentrated the solution Example: Explain the difference between the density of pure ethanol and the concentration of an ethanol solution.

  10. How do we measure concentration? Concentration can be measured a number of ways: • ppm (parts per million) – one part in a million parts • ppb (parts per billion) – one part in a billion parts • g/kg (grams per kilogram) – one gram solute per one kilogram of solvent The chemical standard most used is Molarity Molarity = units: M (molar = mol/L) We’ll come back to concentration later in the chapter…

  11. Evidence of a Chemical Reaction a) A gas is produced. b) A precipitate forms. c) Heat is released or absorbed

  12. Types of Chemical Reactions • Precipitation Reactions • Acid-Base Neutralization Reaction • Oxidation-Reduction (Redox) Reactions • Further classified as: • Combination • Decomposition • Combustion • Single-replacement reactions

  13. Precipitation Reactions • Solubility Rules: Indicate if an ionic compound is soluble or insoluble in water. • Keep in mind that these are just general guidelines, and in reality, some ionic compounds are slightly soluble, and solubility may depend on the temperature.

  14. Precipitation Reactions soluble = compound dissolves in water exists as individual ions in solution physical state is aqueous, (aq) Insoluble = compound does not dissolve in water but remains a solid physical state is shown as solid, (s)

  15. Precipitation Reactions • Example 1: Use the Solubility Rules and identify the ionic compounds are soluble or insoluble by indicating the physical state of each compound.

  16. Precipitation Reactions • Example 1: Use the Solubility Rules and identify the ionic compounds are soluble or insoluble by indicating the physical state of each compound. Soluble (aq) Insoluble (s)

  17. Precipitation Reactions • In a precipitation reaction, two solutions react to form a precipitate (an insoluble solid): AX(aq) + BZ(aq)AZ(s) + BX(aq) precipitate For example: KI (aq)+ Pb(NO3)2(aq)  PbI2 (s) + KNO3 (aq)

  18. Precipitation Reactions To balance and complete the precipitation reactions: 1. Exchange the anions, writing the formulas for the products based on the charges of the ions! 2. Use the Solubility Rules to determine if each product is soluble or insoluble. • If at least one product is insoluble, a precipitation reaction has occurred. Write the formulas for both products, indicating the precipitate as (s), then balance the equation. • If both products are soluble, write NR (=no reaction). 3. Keep in mind that the charges on ions do NOT change in precipitation reactions. For metals that can form more than one charge, use the charge on the metal ion from the reactant side of the equation.

  19. Precipitation Reactions Ex 1. MgSO4(aq)+ NaOH(aq) Ex 2.K2CO3(aq)+ AlCl3(aq) Ex 3. SrBr2(aq)+ Zn(NO3)2(aq) Ex 4. CuSO4(aq)+ NaOH(aq) Ex 5. KI(aq) + Pb(NO3)2(aq)

  20. Acid-Base Neutralization Reactions: Proton Transfer Properties of Acids and Bases Acids Bases –produce hydrogen ions, H+–produce hydroxide ions, OH– –taste sour –taste bitter; feel soapy, slippery –turn blue litmus paper red –turn red litmus paper blue

  21. Arrhenius Definitions acid: A substance that releases H+ when dissolved in water • Some acids are monoprotic (release only H+ per molecule) • e.g. HCl, HBr, HI, HNO3, HClO4 • Some acids are polyprotic (release more than on H+ per molecule) • e.g. H2SO4 and H2CO3 are both diprotic; H3PO4 is triprotic. base: A substance that releases OH– when dissolved in water

  22. Acid-Base Reactions In an acid-base reaction, • H+ from acid reacts with the OH– from base to formwater, H2O • The cation (M+) from base combines with anion from acid (X–) to form asalt. A general equation for an acid-base neutralization reaction is shown below: HX(aq) + MOH(aq) H2O(l) + MX acid base water salt Because water is always produced, an acid always reacts with a base!

  23. Examples Complete and balance each of the equations below: a. HCl(aq) + NaOH(aq) b. H2SO4(aq)+ KOH(aq) c. H3PO4(aq)+ Ca(OH)2(aq)

  24. Acid-Base Reactions with Gas Formation Some acid-base reactions produce carbon dioxide gas, CO2(g), along with water and salt. When the base contains carbonate ion (CO32–) or hydrogen carbonate ion (HCO3–), then the products of the acid-base reaction are water, carbon dioxide gas, and a salt. The general equations for the unbalanced acid-base reactions are: HX(aq)+ MCO3(s) H2O(l) + CO2(g) + MX acid base water carbon dioxide salt HX(aq)+ MHCO3(s)H2O(l)+ CO2(g) + MX acid base water carbon dioxide salt Because water is always produced, an acid always reacts with a base!

  25. Acid-Base Reactions with Gas Formation Complete and balance each of the equations below: a.HCl(aq) + Na2CO3(s) b. HNO3(aq) + CaCO3(s) c. H2SO4(aq) + KHCO3(s) d. HClO4(aq) + Sr(HCO3)2(s)

  26. A double-replacement reaction that produces NH4OH(aq)actually produces ammonia, NH3(g). NH4OH(aq) NH3(g) + H2O(l) Example: Complete and balance the equation below: (NH4)2SO4(aq) + KOH(aq)

  27. Brønsted-Lowry Definition of Acids and Bases • Brønsted-Lowry acid: A substance that donates a proton (H+)—i.e., a proton donor • Brønsted-Lowry base: A substance that accepts a proton (H+)—i.e., a proton acceptor • Unlike an Arrhenius base, a Brønsted-Lowrybase does not need to contain OH–. Why is H+ called a proton?

  28. Brønsted-Lowry Acids and Bases A Brønsted-Lowry acid-base reaction simplyinvolves a proton (H+) transfer. NH3(aq)+ H2O(l)⇄NH4+(aq)+ OH–(aq) • Note: This reaction simply involves H2O donating a H+ ion to NH3 to produce NH4+ and OH–. • In this reaction, H2O is the Brønsted-Lowry acid, and NH3 is the Brønsted-Lowry base. • The conjugate acid-base pairs differ only by a H+. • In this reaction, the conjugate acid-base pairs are NH3 and NH4+and H2O and OH–.

  29. Conjugate Acid-Base Pairs Conjugate acid-base pairs: a Brønsted-Lowry acid/base and its conjugate differ by a H+ HA(aq) + H2O(l)⇄H3O+(aq) + A–(aq) • For the reaction above, when HA donates H+ to H2O, it leaves behind A–, which can act as a base for the reverse reaction. • An acid and base that differ only by the presence of H+ are conjugate acid-base pairs. • The general reaction for the dissociation (or ionization) of an acid can be represented as above, where the double-arrow indicates both the forward and reverse reactions can occur. • Note: The double arrow (⇄) indicates the reaction is reversible (goes in both directions).

  30. Conjugate Acid-Base Pairs Determine the Brønsted-Lowry acid and base in each of the following reactions: a. CH3COOH(aq)+ NH3(aq)⇄NH4+(aq)+ CH3COO–(aq) b.NH3(aq)+ H2O(l)⇄NH4+(aq)+ OH–(aq) c. H2O(l)+ H2SO4(aq)⇄ H3O+(aq)+ HSO4–(aq)

  31. Oxidation-Reduction (Redox) Reactions Types of Redox Reactions • Combination Reaction • Decomposition Reaction • Single-Replacement (or Displacement) Reaction • Combustion Reaction

  32. Combination Reactions A + Z  AZ Usually a meal and a non-metal react to form a solid ionic compound: metal + nonmetal ionic compound(s) Δ

  33. Combination Reactions:A + Z  AZ Complete and balance each of the equations below: a. Na(s) + Cl2(g) b. Al(s) + O2(g) c. Zn(s) + S8(s) d. Mg(s) + N2(g) Δ Δ Δ Δ

  34. Decomposition Reactions:AZ  A + Z Be able to classify and balance decomposition reactions. (You won’t need to predict products.) a. ___ KHCO3(s) ___ K2CO3(s)+ ___ H2O(l)+ ___ CO2(g) b. ___ Al2(CO3)3(s) _____ Al2O3(s)+ _____ CO2(g) c. ___ KClO3(s)_____ KCl(s)+ _____ O2(g) Δ, MnO2 Δ Δ

  35. Single-Replacement Reactions and the Activity Series Activity Series: Relative order of elements arranged by their ability to replace cations in aqueous solution Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Au Note: The Activity Series will be given to you on quizzes and exams.

  36. Single Replacement Reactions:A + BZ  AZ + B metal A + aqueous solution B aqueous solution A + metal B Zn(s) + Sn2+(aq)Sn(s) + Zn2+(aq) Cu(s) + 2 Ag+(aq) 2 Ag(s) + Cu2+(aq)

  37. Activity Series: Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Au To balance and complete the following rxns: • Check the Activity Series to see which metal is more active. • The more active metal replaces the less active by going into solution as an ion and the less active metal ion comes out as a solid. 1. Mg(s) + AlCl3(aq) 2.Al(s) + CdSO4(aq) 3. Cd(s) + AgNO3(aq) 4.Ag(s) + Mg(NO3)2(aq)

  38. Activity Series: Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Au To balance and complete the following reactions: • Check the Activity Series to see which metal is more active, the metal or H. • The more active metal replaces the less active by going into solution as an ion and the H comes out as hydrogen gas, H2(g). metal A + acid solution  aqueous solution A + H2(g) 1. Zn(s)+ HCl(aq) • Al(s)+ HNO3(aq) 3.Cu(s) + HI(aq)

  39. Active Metals: Li > K> Ba > Sr > Ca > Na Active metals react directly with water: The activemetal replaces the less active by going into solution as an ion with hydroxide, OH–, and the H comes outas hydrogen gas, H2(g). active metal + H2O(l)metal hydroxide + H2(g) • Ca(s)+H2O(l) • Na(s)+ H2O(l) • Fe(s)+ H2O(l)

  40. Combustion Reactions CxHy + O2(g)H2O(g) + CO2(g) CxHyOz+ O2(g)H2O(g) + CO2(g) • C3H8(g) + O2(g) • C6H6O(l) + O2(g) • C2H2(g) + O2(g) Δ Δ Δ

  41. Identify the Reactions For each of the following, 1. Identifythe type of reaction using the letters designated below: Combination (C) Precipitation (P) Decomposition (D) Acid-Base Neutralization (N) Combustion (B) Single Replacement/Displacement (SR) 2. Balancethe equations: ____a. ___ Mg(NO3)2(aq) + ___ K3PO4(aq)___ Mg3(PO4)2(s) + ___ KNO3(aq) ____b. ___ Ni(OH)3(s) + ___ HCl(aq)___ H2O(l) + ___ NiCl3(aq) ____c. ___ Al(HCO3)3(aq) ___ CO2(g) + ___ H2O(g) + ___ Al2(CO3)3(s) ____d. ___ Fe(s) + ___ Pb(NO3)2(aq)___ Pb(s) + ___ Fe(NO3)3(aq) Δ

  42. Electrolytes and Non-Electrolytes Electrolytes and electrical conductivity • If a solution conducts electricity, it contains ions • A solution that contains many ions is a strong electrolyte. • A solution that contains only a few ions is a weak electrolyte. • A solution that contains only a no ions is a nonelectrolyte.

  43. Electrolytes and Non-Electrolytes • A solution that contains many ions is a strong electrolyte. • Lightbulb burns brightly in a light bulb conductivity apparatus. • A solution that contains only a few ions is a weak electrolyte. • Light bulb burns dimly in a light bulb conductivity apparatus. • A solution that contains only a no ions is a nonelectrolyte. • Light bulb does not light in a light bulb conductivity apparatus.

  44. Strong Electrolytes strong electrolytes: substances that are good conductors of electricity • These substances break up to produce many ions in water • many ions present to move electrons/conduct electricity strong electrolyte For example, NaCl(s) Na+(aq) + Cl–(aq) KOH(s) K+(aq) + OH–(aq) HBr(aq)H+(aq) + Br–(aq) Examples: strong acids, strong bases, all soluble ionic compounds H2O H2O H2O

  45. Weak Electrolytes weak electrolytes: substances that are weak/poor conductors of electricity • These substances mostly remain intact as compounds, producing very few ions in water • only a few ions present to move electrons/conduct electricity weak electrolyte For example, Mg(OH)2(s) Mg(OH)2(s) HNO2(aq) HNO2(aq) Examples: weak acids, weak bases, insoluble ionic compounds H2O

  46. Non-electrolytes nonelectrolytes: substances that cannot conduct electricity • These molecules never break down into ions. • They always remain intact as neutral molecules that have no charge no ions to move electrons/conduct electricity For example, C12H22O11(s) C12H22O11(aq) Examples: sugar (e.g. sucrose), ethanol (C2H5OH), and all other molecules that are not acids H2O

  47. Acids and Bases Know the following acids and bases. All other acids and bases are weak! Strong acids and bases dissolve in water to form ions (or species) in solution. HNO3(aq) H+(aq) + NO3–(aq) Ca(OH)2(aq) Ca2+(aq) + 2 OH–(aq) Note: H2SO4(aq) is a strong acid and diprotic (able to release 2 H+ ions), but it generally ionizes to release only one H+ ion in water: H2SO4(aq) H+(aq) + HSO4–(aq) Recognize that both protons are not released in water!

  48. Molecular, Ionic and Net Ionic Equations • molecular equation: equation showing reactants and products as compounds • total/complete ionic equation:shows strong electrolytes as individual ions while all solids, liquids, gases, and weak electrolytes remain intact as compounds • spectator ions:ionsthat do not form solids, liquids, gases, weak electrolytes – appear on both sides of total ionic equation as ions • net ionic equations:show only solids, liquids, gases, weak electrolytes (weak acids and weak bases), and ions undergoing a chemical change/reaction – excludes spectator ions

  49. Net Ionic Equations Guidelines for Writing Net Ionic Equations • Balance the chemical/molecular equation. • Convert the molecular equation to total ionic equation • Leave solids, liquids, gases, and weak acids and bases as compounds • Show strong acids and all aqueous ionic compounds as ions in solution. • Cancel spectator ions to get net ionic equation • If canceling spectator ions eliminates all ions  NO REACTION (NR) • If coefficients can be simplified, do so to get the lowest ratio. • Make sure total charges on both sides of the equation are equal.

More Related