Types of Chemical Reactions and Solution Stoichiometry

Types of Chemical Reactions and Solution Stoichiometry

Aqueous Solutions : Water

SolutionsSolution- Homogeneous substanceSolute- dissolving substance; in lesser quantity; changes phaseSolvent- dissolving substance; in greater quantity

Solubility of ionic substances in water depends upon the relative attractions of the ions for each other and the attractions of the ions for water. Generally if the substance is polar, most likely it will dissolve in water. Like dissolves like.

ElectrolytesElectrolytes- ions in waterStrong electrolyte- conducts electricitySoluble salts- ionic soluble solids break down to form ionsStrong acids- completely dissociates its ionsBinary acids- HCl, HBr, HI; all others are weakTernary acids- If the number of oxygen out number the number of hydrogen by two or more the acid is strong; H2SO4, HClO4, HClO3, HNO3Strong Bases- soluble hydroxides; Group IA, Group IIA from calcium down

Weak and Non-electrolytesWeak electrolyte- weakly conducts electricityWeak acidsWeak basesNonelectrolyte- does not conduct electricityHighly insoluble salts

Arrhenius DefinitionArrhenius Acid- substance containing ionizable H+Base- contains OH-

Ex. Write an equation for the dissociation of sodium hydroxide.

Ex. Write and equation for the dissociation of acetic acid.

ExamplesEx. Write the net ionic equation for:Aqueous potassium chloride is added to aqueous silver nitrate to form silver chloride plus aqueous potassium nitrateAqueous potassium hydroxide is mixed with aqueous iron (III) nitrate to form a precipitate of iron (III) hydroxide and aqueous potassium nitrate

ConcentrationsCommon Terms of Solution Concentration

Stock- routinely used solutions prepared in concentrated form

Concentrated- relatively large ratio of solute to solvent

Dilute- relatively small ratio of solute to solvent ConcentrationsMolarity- unit of concentration

M = moles of solute/ L of solution Ex. 1 M NaCl (given before dissociation)

NaCl Na+ + Cl-1 M 1 M 1 M

Co(NO3)2 Co2+ + 2 NO3-0.50 M 0.50 M 1. 0 M

ExampleEx. What are the concentrations of each ion in a solution of 0.32 M Ba(NO3)2?

DilutionM = mols/LMol = M x L or MVDilution

Moles before dilution = moles after dilution

M1V1= M2V2

Precipitation ReactionsPrecipitation reaction- precipitate forms- insoluble substance formed from a reaction

Solubility rules

ConcentrationWeight percent- concentration unit; used in medical field

# of grams of solute/ 100 g of water

Ex. How would you prepare a 5% glucose solution?

Net Ionic EquationsMolecular EquationIonic EquationNet Ionic EquationSpectator Ions

Stoichiometry of Chemical Reactions

Ex. Calculate the mass of solid KCl necessary to precipitate all of the Ag+ ions from a solution of 0.75 L of a 0.15 M AgNO3 solution.

StoichiometryLimiting Reactant- Must take into consideration which one is limiting.Ex. When aqueous solutions of Na2SO4 and Pb(NO3)2 are mixed, PbSO4 precipitates. Calculate the mass of PbSO4 formed when 1.25 L of 0.0500 M Pb(NO3)2 and 2.00 L of 0.0250 M Na2SO4 are mixed.

Bronsted-Lowry Acid-Base TheoryKOH + HC2H3O2 KC2H3O2 + H2O

The hydroxide ion is a very strong base and can strip the proton from a weak electrolyte like HC2H3O2OH- is base and HC2H3O2 is acid.

Acid-Base Reaction= Neutralization ReactionWhen exactly the same amount of acid and base has been added (according to the coefficients in the balanced equation) the acid has been neutralizedExampleEx. What volume of 0.100 M HCl solution is needed to neutralize 25.0 mL of a 0.350 M NaOH solution?

ExampleEx. In a certain experiment 28.0 mL of 0.250 M HNO3 and 53.0 mL of 0.125 M KOH are mixed. Calculate the amount of water formed in the resulting reaction. What is the concentration of H+ and OH- ions in excess after the reaction goes to completion?Monoprotic and Diprotic AcidsMonoprotic acid- contains only one ionizable hydrogen Ex. HCl, HNO3, HC2H3O2

Diprotic acid- contains two ionizable hydrogens Ex. H2SO4

TitrationsTitration- delivery of a measured volume of a solution of known concentration (titrant) into a solution containing the substance being analyzed (analyte).Point where enough titrant has been added is the point where the indicator actually changes color is called endpoint.Use an indictor to determine where the point isAcid/Base TitrationCommon indicator for strong acids and bases is phenolphthaleinTurns from colorless (acid) to pink (base)Before beginning the titration, the titrant is measured for exact concentration, called standardizing the solution. Use a solid to get exact number of moles of acid or base. Does not matter how much water is used.

ExampleEx. A sample of an analgesic drug was analyzed for aspirin, a monoprotic acid, HC9H7O4, by titration with a base. In a titration, a 0.500 g sample of the drug required 21.50 mL of 0.100 M NaOH for complete neutralization. What percent by mass of the drug was aspirin?

ExampleEx. A student carries out an experiment to standardize a sodium hydroxide solution. To do this the student weighs out a 1.3009 g sample of potassium hydrogen phthalate (KHP, MM=204.22), a compound with the formula KHC8H4O4 which has one acidic hydrogen. The student dissolves the KHP in distilled water and titrates it with the sodium hydroxide solution. The difference between the final and initial buret readings indicates that 41.20 mL of sodium hydroxide solution is required to react with the 1.3009 g of KHP. Calculate the concentration of the sodium hydroxide solution.

ExampleEx. In a reaction involving Ba(OH)2 and HCl, determine the number of mL of a 0.50 M solution of Ba(OH)2 necessary to neutralize 25.0 mL of a 0.35 M solution of HCl.

RedoxElectrons are transferred- called oxidation-reduction reactions or redox.

Use oxidation numbers (states) to keep track of electron transfer.

Oxidation numbersRules for assigning oxidation states:The oxidation state of an atom in an element is 0.The oxidation state of a monatomic ion is the same as its charge.In its compounds, fluorine is always assigned an oxidation state of 1.Oxygen is usually assigned an oxidation state of 2 in its covalent compounds. Exceptions are peroxides O22- and in OF2 where oxygen is +2.In its covalent compounds with nonmetals, hydrogen is assigned an oxidation state of +1.The sum of the oxidation states must be 0 for an electrically neutral compound. For an ion, the sum of the charges must equal the charge on the ion.

ExamplesEx. Assign oxidation numbers for:CH4, SO3, SO42-, Fe3O4, C12H22O11

Oxidation Reductions ReactionsOxidation increase in oxidation state ( a loss of electrons)

Reduction- decrease in oxidation state ( a gain of electrons)

Oxidizing agent- electron acceptor- substance being reducedReducing agent- electron donor- substance being oxidized.

RedoxWhen identifying what is oxidized or reduced use specific element.When identifying oxidizing or reducing agent identify whole compound.

ExampleEx. 2Al + 3 I2 2 AlI3Identify the substance oxidized, reduced, oxidizing agent and reducing agent.

Balancing Redox Equations- Acid SolutionsBalancing Redox Reactions

In acidic solutions, there are excess hydrogen ions and water molecules:Break into two half reactions: one reduction, one oxidationWrite half reactions.Balance all elements other that H and O.Balance oxygen using H2O.Balance hydrogen by using H+.Balance charge using electrons.If needed multiply to make sure all electrons lost by one substance are gained by another. Same number of electrons lost and gained.

ExampleEx. MnO4- + Fe2+ Fe3+ + Mn2+

Balancing Redox Equations- Basic SolutionsIn basic solutions, there are excess hydroxide ions and water molecules:Balance according to acid rules.For every H+ add OH- to both sides of the equation.Form water on side containing H+ and OH-.Add together and eliminate species on both sidesExampleEx. Ag + CN- + O2 Ag(CN)2-