Presentation on theme: "Chapter 4 Reactions in Aqueous Solutions"— Presentation transcript:
1 Chapter 4 Reactions in Aqueous Solutions Many chemical and almost all biological reactions occur in the aqueous mediumSubstances (solutes) that dissolve in water (solvent) can be divided into two categories:ElectrolytesNon-Electrolytes
2 Three Major Types of Reactions Precipitation Reaction – the product an insoluble substance separates from the solutionAcid/Base Reactions – A proton transfer from an acid to a baseOxidation/Reduction (Redox) “the bane of the AP Test” – Electrons are transferred from a reducing agent to an oxidizing agent
3 Solution Stoichiometry Quantitative studies with known concentrations (Molarity) of solutionsGravimetric AnalysisTitrations
4 General Properties of a Solution Solution – a homogenous mixture of two or more substancesSolution may be gaseous (air), solid (alloy) or Liquid (salt water)In this chapter we will deal only with aqueous solutionsMost common Solvent - Water
5 Electrolytes versus Nonelectrolytes Electrolytes – Ionic compounds that completely or partially dissociate in solution with the ability to pass electric current in solutionAcids/Bases will ionize in solution, therefore electricity can be conductedNon-Electrolytes – Molecular compounds that do not dissociate in solution, therefore no electric current can be pass
6 Ionic Compounds in Solution Water is a great solvent for ionic compounds because it is polar, the positive end attracts the Negative Ion and vice versa
7 Acids and Bases as Electrolytes Some acid/bases competely dissociate in solutionHClHNO3H2SO4Ba(OH)2NaOHWhile others only partially dissociateCH3COOHHFHNO2HN3
8 Writing Partial Dissociation Equations Partial dissociation equations are written with a double arrow, indicating a reversible reactionWrite partial dissociation for CH3COOH
9 Precipitation Reactions A double replacement reaction (metathesis) in which a product is insoluble
10 Solubility Rules In water at 25 Degrees All common compounds of Group I and ammonium ions are soluble.All nitrates, acetates, and chlorates are soluble.All binary compounds of the halogens (other than F) with metals are soluble, except those of Ag, Hg(I), and Pb.All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I). The latter three are slightly soluble.Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.
11 Soluble or Insoluble at 25 Degrees Celsius in Water PbSO4BaCO3Li3PO4FeSCa(OH)2Co(NO3)3
12 Net Ionic EquationsWrite the correctly balanced equation and decide on state of each productWrite free state of all ions and insoluble productCancel out spectator ions – anyone not part of the reactionCheck charges and balancing in net ionic
13 Practice Net Ionic Predict, Balance and write net ionic Lead Nitrate and Potassium IodideBarium Chloride and Sodium SulfatePotassium Phosphate and Calcium NitrateAluminum Nitrate and Sodium Hydroxide
14 Acid – Base ReactionsAcids react with metal such as Zn, Mg and Fe to produce hydrogen gasAcids react with carbonates and bicarbonates to produce carbon dioxide gas, water and the salt
15 Bronsted Acid and Bases Bronsted Acid is a proton donorBronsted Base is a proton acceptorHCl (aq) H+ (aq) + Cl-(aq)In water the H+ attracts to the water molecule producing the hydronium ion
16 Monoprotic AcidsEach unit of acid yields one hydrogen ion upon ionization
17 Diprotic AcidsEach unit of the acid gives up two hydrogen ions in two separate steps (they strip)
18 Triprotic AcidsYield three hydrogen ions in three separate steps (they strip)
19 Bronsted Acid is a proton donor Bronsted Base is a proton acceptor Classify each of the following as an Bronsted acid or Bronsted base, explain your reasoning based on the definitionHBrSO-24HIHCO-3NO2
20 Neutralization Reaction Acid and Base will form Salt and Water Write the net ionic for the followingHydrochloric acid and Sodium HydroxideSulfuric acid and Aluminum Hydroxide
21 Acid – Base Reactions Leading to Formation of a Gas Certain Salts – Carbonates, bicarbonates, sulfites and sulfides react with acids to form gaseous products
22 Oxidation NumbersOxidation Reaction – refers to half reaction that involves loss of electronsReduction reaction – refers to a half reaction that involves the gain of electronsThe extent of oxidation in a redox reaction must be equal to the extent of reduction; that is the number of electrons lost by a reducing agent must be equal to the number of electrons gained by an oxidizing agent
23 The half-reactions of a redox reaction or oxidation-reduction reaction
24 Oxidation NumberThe number of charges the atom would have in a molecule if electrons are transfer completely
25 The convention is that the cation is written first in a formula, followed by the anion. For example, in NaH, the H is H-; in HCl, the H is H+.The oxidation number of a free element is always 0.The atoms in He and N2, for example, have oxidation numbers of 0.The oxidation number of a monatomic ion equals the charge of the ion.For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3.The usual oxidation number of hydrogen is +1.The oxidation number of hydrogen is -1 in compounds containing elements that are less electronegative than hydrogen, as in CaH2.The oxidation number of oxygen in compounds is usually -2.Exceptions include OF2, since F is more electronegative than O, and BaO2, due to the structure of the peroxide ion, which is [O-O]2-.The oxidation number of a Group IA element in a compound is +1.The oxidation number of a Group IIA element in a compound is +2.The oxidation number of a Group VIIA element in a compound is -1, except when that element is combined with one having a higher electronegativity.The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl.The sum of the oxidation numbers of all of the atoms in a neutral compound is 0.The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.For example, the sum of the oxidation numbers for SO42- is -2.
26 Assign oxidation numbers to all the elements in the following compounds Na2OHNO2Cr2O7-2PF3MnO4-
27 Arrange the following species in order of increasing oxidation number of the sulfur atoms H2SSO2SO3S8H2SO4S-2HS-
28 Concentration Molarity = moles of solute/liters of solution Example: How many grams of potassium dichromate are required to prepare a 125ml solution whose concentration is 1.83M
29 ConcentrationIn a biochemical assay a chemist needs a to add 4.07g of glucose to a reaction mixture. Calculate the volume in milliliters the volume of a 3.16M glucose she should use
30 Dilution of SolutionsThe procedure of making a less concentrated solution from a high concentration solutionMiVi = MfVf
31 Dilution ProblemDescribe hou you would prepare 2.50 * 102 ml of a 2.25M H2SO4 solution, starting with a 7.41 M stock solution of H2SO4
32 Dilution Problem #2How would you prepare a 200ml of a .866M KOH solution, starting with 5.07M stock solution
33 Acid – Base TitrationsIn a titration a solution of an accurately known concentration, called the standard is added gradually to another solution of unknown until reaction is neutralized (equivalence point)Indicators are used to color the reaction when it is complete
34 Titration ProblemIn a titration experiment, a student finds that 25.46ml of a NaOH solution is needed to neutralize .6092g of KHP. What is the concentration of the NaOH solution?
35 Titration Problem #2How many milliliters of a .836M NaOH solution is needed to neutralized 25ml of a .335M of H2SO4?
36 Solution Stoichiometry When sodium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 150ml 3M of silver nitrate?
37 When Magnesium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 4.5M in 250ml of silver nitrate? What is the name of the other product of the reaction?