Solutions

1. Solutions – Solution, A homogeneous mixture in which all of the material is in the same state. – Substances present in lesser amounts, called solutes, are dispersed uniformly throughout the substance in the greater amount, the solvent – Aqueous solution — a solution in which the solvent is water – Nonaqueous solution — any substance other than water is the solvent – Many of the chemical reactions that are essential for life depend on the interaction of water molecules with dissolved compounds.

2. Aqueous Solutions Polar Substances • An individual water molecule consists of two hydrogen atoms bonded to an oxygen atom in a bent (V-shaped) structure. • The oxygen atom in each O–H covalent bond attracts the electrons more strongly than the hydrogen atom. • O and H nuclei do not share the electrons equally. • – Hydrogen atoms are electron-poor compared with a neutral hydrogen atom and have a partial positive charge, indicated by the symbol δ+. • – The oxygen atom is more electron-rich than a neutral oxygen atom and has a partial negative charge, indicated by the symbol 2δ-. • Unequal distribution of charge creates a polar bond, which makes them good solvents for ionic compounds. • Individual cations and anions are called hydrated ions.

3. Aqueous Solutions • Electrolyte — any compound that can form ions when it dissolves in water – When strong electrolytes dissolve, constituent ions dissociate completely, producing aqueous solutions that conduct electricity very well. – When weak electrolytes dissolve, they produce relatively few ions in solution. Aqueous solutions, of weak electrolytes do not conduct electricity as well as solutions of strong electrolytes. – Nonelectrolytes dissolve in water as neutral molecules and have no effect on conductivity. CH3CO2H(aq) → CH3CO2-(aq) +H+(aq)

4. Aqueous Solutions Molarity • Most common unit of concentration • Most useful for calculations involving the stoichiometry of reactions in solution • Molarity of a solution is the number of moles of solute present in exactly 1 L of solution: • Units of molarity — moles per liter of solution (mol/L), abbreviated as M • Relationship among volume, molarity, and moles is expressed as:

5. Preparation of Solutions

6. Calculating Moles from Volume Calculating Volume from Mass

7. Preparation of Solutions Problem: What mass of oxalic acid, H2C2O4, is required to make 250. mL of a 0.0500 M solution?

8. Preparation of Solutions 0.250 L of water was used to make 0.250 L of solution. Notice the water left over.

9. Preparation of Solutions (Vs) (Ms) = moles of solute = (Vd) (Md).

10. Preparation of Solutions PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do?

11. Ion Concentrations in Solution • Classify each compound as either a strong electrolyte or a nonelectrolyte. • If a compound is a nonelectrolyte, the concentration is the same as the molarity of the solution. • If a compound is a strong electrolyte, determine the number of each ion contained in one formula unit and find the concentration of each species by multiplying the number of each ion by the molarity of the solution.

12. Preparation of Solutions PROBLEM: Dissolve 5.00 g of NiCl2•6 H2O in enough water to make 250 mL of solution. Calculate molarity of the solution and the concentration of each of the ions.

13. Zinc reacts with acids to produce H2 gas. Have 10.0 g of Zn What volume of 2.50 M HCl is needed to convert the Zn completely? SOLUTION STOICHIOMETRY

14. Zinc reacts with acids to produce H2 gas. If you have 10.0 g of Zn, what volume of 2.50 M HCl is needed to convert the Zn completely? SOLUTION STOICHIOMETRY

15. Limiting Reactants in Solutions • The concept of limiting reactants applies to reactions that are carried out in solution and reactions that involve pure substances. • If all the reactants but one are present in excess, then the amount of the limiting reactant can be calculated. • When the limiting reactant is not known, one can determine which reactant is limiting by comparing the molar amounts of the reactants with their coefficients in the balanced chemical equation. • Use volumes and concentrations of solutions of reactants to calculate the number of moles of reactants.

16. Ionic Equations • Chemical equation for a reaction in solution can be written in three ways: 1. Overall equation — shows all of the substances present in their undissociated form 2. Complete ionic equation — shows all of the substances present in the form in which they actually exist in solution 3. Net ionic equation – Derived from the complete ionic equation by omitting all spectator ions, ions that occur on both sides of the equation with the same coefficients – Demonstrate that many different combinations of reactants can give the same net chemical reaction

17. Ionic Equations • There are three ways to write reactions in aqueous solutions. • Molecular equation: • Show all reactants & products in molecular or ionic form • Total ionic equation: • Show the ions and molecules as they exist in solution • Net ionic equation: • Shows ions that participate in reaction and removes spectator ions. Spectator ions do not participate in the reaction.

18. Mg(s)+ 2HCl(aq)→ H2(g) + MgCl2(aq) The molecular formula above can be written as the total ionic formula Mg(s)+ 2H+(aq)+ 2Cl-(aq)→ H2(g)+ Mg2+(aq)+ 2Cl-(aq) The two Cl- ions are SPECTATOR IONS — they do not participate. Could have used NO3- for the spectator ion as salts of nitrates are all soluble. By leaving out the spectator ions out you get the net ionic reaction Mg(s) + 2 H+(aq) ---> H2(g) + Mg2+(aq) Net Ionic Equations

19. Classifying Chemical Reactions • Exchange reactions • Single Displacement Reactions – one element displaces another from a compound – AB + C  AC + B • Metathesis Reactions - Exchange Reactions – AB + CD  AD + CB • Precipitation: products include an insoluble substance which precipitates from solution as a solid • Acid-base neutralization: product is a salt and water • Gas formation – primarily the reaction of metal carbonates • Condensation reactions (and the reverse, cleavage reactions) • Combination Reactions (Condensation) – More than one reactant, one product. Some condensation reactions are redox reactions. – A + B  AB • Decomposition Reactions (Cleavage) – Single reactant, more than one product – AB  A + B • Redox (Oxidation Reduction Reactions) – Electrons are transferred between reactants. Oxidation numbers of some elements change; at least one element must increase and one must decrease in oxidation number. • Single Displacement Reactions are always Redox reactions as well. • oxidant + reductant  reduced oxidant + oxidized reductant

20. Aqueous Chemical Reactions: Metathesis EXCHANGE REACTIONS The anions exchange places between cations. Precipitation Pb(NO3)2(aq) + 2 KI(aq) → PbI2(s) + 2KNO3(aq) Pb2+(aq) + 2 I-(aq) → PbI2(s) Neutralization: NaOH(aq) + HCl(aq)→ NaCl(aq)+ H2O(l) OH-(aq) + H+(aq) → H2O(lq) Gas Formation MgCO3(s)+ 2HCl(aq) → 2Mg(Cl)2(aq)+ H2O(l) + CO2(g)

21. Precipitation Reactions • A reaction that yields an insoluble product, a precipitate, when two solutions are mixed • Are a subclass of exchange reactions that occur between ionic compounds when one of the products is insoluble • Used to isolate metals that have been extracted from their ores and to recover precious metals for recycling

22. Water Solubility of Ionic Compounds If one ion from the “Soluble Compound” list is present in a compound, the compound is water soluble.

23. Precipitation Reactions In addition to understanding solubility. It is equally important to know if a reaction will occur. K2Cr2O7(aq) + AgNO3(aq) → 2KNO3(aq) + Ag2Cr2O7(s) KBr(aq) + NaCl(aq) → KCl(aq) + NaBr(aq)

24. Neutralization Reactions • A reaction in which an acid and a base react in stoichiometric amounts to produce water and a salt • Strengths of the acid and base determine whether the reaction goes to completion 1. Reactions that go to completion • Reaction of any strong acid with any strong base • Reaction of a strong acid with a weak base • Reaction of weak acid with a weak base 2. Reaction that does not go to completion is a reaction of a weak acid or a weak base with water

25. Neutralization Reactions A brief history of Acid-Base Identification Systems

26. StrongBrönsted-Lowry acids are strong electrolytes HCl hydrochloric H2SO4 sulfuric HClO4perchloric HNO3 nitric ACIDS Acetic acid HNO3 A Brönsted-Lowry Acid → H+ in water Weak Brönsted-Lowry acids are weak electrolytes CH3CO2H acetic acid (CH3COOH) H2CO3 carbonic acid H3PO4 phosphoric acid HF hydrofluoric acid Carbonic Acid

27. Polyprotic Acids • Acids differ in the number of hydrogen ions they can donate. – Monoprotic acids are compounds capable of donating a single proton per molecule. – Polyprotic acids can donate more than one hydrogen ion per molecule.

28. Ammonia, NH3 an Important weak Base NaOH(aq) → Na+(aq) + OH-(aq) BASES Brönsted-LowryBase → OH- in water NaOH is a strong base

29. ACIDS Nonmetal oxides can be acids CO2(aq) + H2O(l) → H2CO3(aq) SO3(aq) + H2O(l) → H2SO4(aq) NO2(aq) + H2O(l) → HNO3(aq) Acid Rain is an example of nonmetal oxides behaving as acids. This process can result from burning coal and oil. BASES • Metal oxides can be bases • CaO(s)+H2O(l) → Ca(OH)2(aq) CaO in water. Phenolphthalein indicator shows a of calcium oxide solution is basic.

30. pH, a Concentration Scale pH: a way to express acidity -- the concentration of H+ in solution. Low pH: high [H+] High pH: low [H+] Acidic solution pH < 7 Neutral pH = 7 Basic solution pH > 7

31. The pH Scale pH = log (1/ [H+]) = - log [H+] In a neutral solution, [H+] = [OH-] = 1.00 x 10-7 M at 25 oC pH = - log [H+] = If the [H+] of soda is 1.6 x 10-3 M, the pH is ____. If the pH of Coke is 3.12, it is _____.

32. Acid-Base Strength Identification You should know the strong acids & bases

33. Oxidation-Reduction Reactions in Solution • Oxidation-reduction reactions — electrons are transferred from one substance or atom to another. • Oxidation-reduction reactions that occur in aqueous solution are complex, and their equations are very difficult to balance. • Two methods for balancing oxidation-reduction reactions in aqueous solution are: • Oxidation states — overall reaction is separated into an oxidation equation and a reduction equation • Half-reaction

34. Oxidation-Reduction Reactions • The term oxidation was first used to describe reactions in which metals react with oxygen in air to produce metal oxides. – Metal acquires a positive charge by transferring electrons to the neutral oxygen atoms of an oxygen molecule. – Oxygen atoms acquire a negative charge and form oxide ions (O2-). – Metals lose electrons to oxygen and have been oxidized—oxidation is the loss of electrons. – Oxygen atoms have gained electrons and have been reduced—reduction is the gain of electrons.

35. Oxidation-Reduction Reactions • Oxidation and reduction reactions are now characterized by a change in the oxidation states of one or more elements in the reactants. • Oxidation states of each atom in a compound is the charge that atom would have if all of its bonding electrons were transferred to the atom with the greater attraction for electrons. Atoms in their elemental form are assigned an oxidation state of zero. • Oxidation-reduction reactions are called redox reactions, in which there is a net transfer of electrons from one reactant to another. The total number of electrons lost must equal the total number of electronsgained. • Oxidants and reductants • Oxidants – Compounds that are capable of accepting electrons are called oxidants, or oxidizing reagents, because they can oxidize other compounds. • An oxidant is reduced in the process of accepting electrons. • Reductants – Compounds that are capable of donating electrons are called reductants, or reducing agents, because they can cause the reduction of another compound. • A reductant is oxidized in the process of donating electrons.

36. OXIDATION NUMBERS NH3 ClO- H3PO4 MnO4- Cr2O72-

37. Recognizing a Redox Reaction 2 Ag(s) + Cu2+(aq) → Al+(aq) + Cu(s) Hydrogen Fuel Cell 2 H2(g) + O2(g) → 2H2O(l) Thermite reaction Fe2O3(s) + 2Al(s) → 2 Fe(s) + Al2O3(s)

38. Balancing Redox Equations Using Oxidation States Will be covered in Chem 102

39. Quantitative Analysis: Titrations • Quantitative analysis — used to determine the amounts or concentrations of substances present in a sample by using a combination of chemical reactions and stoichiometric calculations • Titration – A method in which a measured volume of a solution of known concentration, called the titrant, is added to a measured volume of a solution containing a compound whose concentration is to be determined (the unknown) – Reaction must be fast, complete, and specific (only the compound of interest should react with the titrant) – Equivalence point — point at which exactly enough reactant has been added for the reaction to go to completion (computed mathematically)

40. Acid-Base Titrations • Most common acids and bases are not intensely colored – Rely on an indicator • Endpoint — point at which a color change is observed, which is close to the equivalence point in an acid-base titration

41. Standard Solutions • A solution of a primary standard whose concentration is known precisely • A primary standard is non-hygroscopic, has a high mass, is fairly inexpensive compound, is of known reactive ability that can be accurately weighed for use as a titrant. • A standard solution is used to determine the concentration of the titrant. • Accuracy of any titration analysis depends on accurate knowledge of the concentration of the titrant. • Most titrants are first standardized—their concentration is measured by titration with a standard solution.

42. Titration 1. Add titrant solution from the buret. 2. Reagent (base) reacts with compound (acid) in solution in the flask. 3. Indicator shows when exact stoichiometric reaction has occurred. 4. Net ionic equation H++ OH- → H2O 5. At equivalence point moles H+= moles OH-

43. ACID-BASE REACTIONSTitrations H2C2O4(aq) + 2 NaOH(aq) → Na2C2O4(aq) + 2 H2O(l) acidbase Carry out this reaction using a TITRATION. Oxalic acid, H2C2O4

44. PROBLEM: Standardize a solution of NaOH — i.e., accurately determine its concentration. 1.065 g of H2C2O4 (oxalic acid) requires 35.62 mL of NaOH for titration to an equivalence point. What is the concentration of the NaOH? Quantitative Analysis: Titrations

45. PROBLEM : Use standardized NaOH to determine the amount of an acid in an unknown. Apples contain malic acid, C4H6O5. 76.80 g of apple requires 34.56 mL of 0.663 M NaOH for titration. What is weight % of malic acid? Quantitative Analysis: Titrations HOOCCH2COHCOOH(aq) + 2NaOH(aq) → Na2C4H4O5(aq) + 2H2O(l)

46. Energy & Chemistry 2H2(g) + O2(g) → 2H2O(g) + heat and light This can be set up to provide ELECTRIC ENERGY in a fuel cell. Oxidation: 2 H2 → 4 H+ + 4 e- Reduction: 4 e- + O2 + 2 H2O → 4 OH- H2/O2 Fuel Cell Energy

47. Energy & Chemistry ENERGYis the capacity to do work (w) or transfer heat (q). HEATis the thermal energy that can be transferred from an object at one temperature to an object at another temperature – Net transfer of thermal energy stops when the two objects reach the same temperature. Other forms of energy — • Radiant (light)— energy in light, microwaves, and radio waves • Thermal (kinetic and potential) — results from atomic and molecular motion – Temperature of an object is a measure of the thermal energy content • Chemical — results from the particular arrangement of atoms in a chemical compound; radiant and thermal energy produced in this reaction due to energy released during the breaking and reforming of chemical bonds • Nuclear — radiant and thermal energy released when particles in the nucleus of the atoms are rearranged • Electrical— due to the flow of electrically charged particles

48. Potential & Kinetic Energy Kinetic energy • Potential energy – Energy storedin an object because of the relative positions or orientations of its components – • PE = Fd = mad = mgh = work • Kinetic energy – Energy due to the motion of an object – • KE = ½ mv2

49. Potential & Kinetic Energy

50. Internal Energy (E) • PE + KE = Internal energy (E or U) • Internal Energy of a chemical system depends on • number of particles • type of particles • temperature • The higher the T the higher the internal energy • So, use changes in T (∆T) to monitor changes in E (∆E).