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    • GENERAL PROPERTIES OF AQUEOUS SOLUTIONS Solutions • Solutions are defined as homogeneous mixtures of two or more pure substances. • The solvent is present in greatest abundance. • All other substances are solutes. • Aqueous Solution: A solution which water is the dissolving medium.
    • GENERAL PROPERTIES OF AQUEOUS SOLUTIONS Electrolytic Properties• Electrolyte : A substance whose aqueous solutions contain ions : Substances that dissolves in water to give an electrically conducting solution.•Nonelectrolyte : A substance that does not form ions in solution : Substances that dissolves in water to give an a nonconducting or very poorly condusting solution.
    • GENERAL PROPERTIES OF AQUEOUS SOLUTIONS Ionic Compounds in Water• Water is a very effective solvent of ionic compounds• Electrically neutral molecule, but one end of the molecule (Oatom) rich in electron and has a partial negative charge (δ-)• The other end (H atoms), has a partial positive charge (δ+)
    • GENERAL PROPERTIES OF AQUEOUS SOLUTIONS• Cation (+) are attractive to (δ-) and anions (-) are attractive to(δ+)• As an ionic compound dissolves, the ions become surrounded byH2O molecules.• The ions are said to be solvated.• The solvation process helps stabilize ions in solution and preventscations and anions from recombining.• The ions and their shells (water molecules) are free to move, theions become dispersed uniformly throughout the solution.
    • GENERAL PROPERTIES OF AQUEOUS SOLUTIONS Molecular Compounds in Water• Dissolves in water, solution usually consist of intact moleculesdispersed throughout the solution.• So, molecular compounds are non-electrolytes.• Example: Table sugar (sucrose) and methanol.• A few molecular substances have aqueous solutions that containions.• Example: HCl – Ionizes and dissociates into H+(aq) and Cl-(aq)ions.
    • GENERAL PROPERTIES OF AQUEOUS SOLUTIONS Strong and Weak Electrolytes• Two categories: Weak and strong- Differ in the extend to whichthey conduct electricity.• Strong electrolytes: Dissociates completely when dissolved inwater.• All soluble ionic compounds are electrolytes Eg: NaCl• Strong acid, strong base and soluble ionic compounds.• Weak electrolytes: Only dissociates partially when dissolved inwater.• E.g: CH3COOH: Most of the solute is present as CH3COOHmolecules. Only small fraction of the acid is present as H +(aq) andCH3COO-(aq) ions.
    • GENERAL PROPERTIES OF AQUEOUS SOLUTIONS• For strong electrolyte: Single arrow represent the ionization ofstrong electrolytes. HCl (aq)  H+ (aq) + Cl- (aq) • The absence of a reverse arrow indicates that the H+ and Cl- ions have no tendency to recombine in water.
    • PRECIPITATION REACTIONS• Precipitation Reaction: Reaction thatresult in the formation of an insolubleproduct.• Precipitate: An insoluble solid formed bya reaction in solution. Pb(NO3)2 (aq) + 2KI (aq)  PbI2 (s) +2KNO3 (aq)• To predict whether certain combinationsof ions form insoluble compounds, wemust consider some guidelines concerningthe solubilities of common ioniccompounds.
    • PRECIPITATION REACTIONS Solubility Guidelines for Ionic Compounds• Solubility: The amount of substance that can be dissolved in agiven quantity of solvent at the given temperature.• Insoluble: The attraction between the oppositely charged ions inthe solid is too great for the water molecules to separate the ionsto any significant extend- substance remains undissolved.• No rules based on physical properties to predict the solubility ofcompound.
    • PRECIPITATION REACTIONS• Based on the experimental observation only.• All common ionic compounds of alkali metals ions and the NH4+are soluble in water.
    • PRECIPITATION REACTIONS• To predict the forming of precipitation: 1) Note the ion present in the reactant 2) Consider the possible combinations of the cations and anions 3) Determine whether the any of the product is insoluble Example: Mg(NO3)2 react with NaOH. Will the precipitate form? 1) Existing ions: Mg2+, NO3-, Na+ and OH- 2) Possible reaction: Mg2+ and OH- ; Na+ and NO3- 3) Products: Mg(OH)2 and NaNO3 - Mg(OH)2 is insoluble. - NaNO3 is soluble 4) Balanced equation: Mg(NO3)2 (aq) + 2NaOH(aq)  Mg(OH)2 (s) + NaNO3 (s)
    • PRECIPITATION REACTIONS Exchanged (Metathesis) Reactions• Metathesis comes from a Greek word that means “to transpose.”• AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)• To complete and balance a metathesis reactions: 1. Use the chemical formulas of the reactants to determine the present ions 2. Write the chemical formulas of the products by combining the cation from one reactant with anion from another reactant. 3. Balance the equation.
    • PRECIPITATION REACTIONS Writing Equations for Aqueous Ionic ReactionsThe molecular equationShows all of the reactants and products as intact, undissociatedcompounds. AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)The total ionic equationShows all of the soluble ionic substances dissociated into ions. Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)  AgCl (s) + K+ (aq) + NO3(aq)
    • PRECIPITATION REACTIONS Writing Equations for Aqueous Ionic ReactionsThe net ionic equation• Eliminates the spectator ions and shows the actual chemicalchange taking place.• Those things that didn’t change (and were deleted from thenet ionic equation) are called spectator ions Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq) → AgCl (s) + K+(aq) + NO3-(aq) Ag+(aq) + Cl-(aq) → AgCl (s)
    • PRECIPITATION REACTIONS Writing Net Ionic Equations1. Write a balanced molecular equation.2. Dissociate all strong electrolytes.3. Cross out anything that remains unchanged from the left side to the right side of the equation.4. Write the net ionic equation with the species that remain.
    • PRECIPITATION REACTIONSEXAMPLEPredict whether a reaction occurs when each of the followingpairs of solutions are mixed. If a reaction does occur, writebalanced molecular, total ionic, and net ionic equations, andidentify the spectator ions. (a) potassium fluoride(aq) + strontium nitrate(aq)  (b) ammonium perchlorate(aq) + sodium bromide(aq) 
    • PRECIPITATION REACTIONS Using Molecular Depictions to Understand a Precipitation ReactionConsider the molecular views of the reactants for a precipitationreaction.(a) A: KCl, Na2SO4, MgBr2, or Ag2SO4?(b) B: NH4NO3, MgSO4, Ba(NO3)2, or CaF2?(c) Name precipitate and speactator ions from reaction. Write balancedmolecular, total ionic, and net ionic equations.
    • ACID-BASE REACTIONS• Acids: Substance that ionize in aqueoussolutions to form H+ ions, thereby increasingthe concentration of H+ (aq) ions.• Acids : Called as proton donors.•Molecules from different acids can ionize toform different numbers of H+ ions.
    • ACID-BASE REACTIONS There are only seven strong acids: • Hydrochloric (HCl) • Hydrobromic (HBr) • Hydroiodic (HI) • Nitric (HNO3) • Sulfuric (H2SO4) • Chloric (HClO3) • Perchloric (HClO4)
    • ACID-BASE REACTIONS• Base: Substance that accept or react with H+ ions.• Produce OH- ions when dissolve in water• The strong bases are the soluble metal salts of hydroxide ion Example: Ca(OH)2• Compounds that do not contain OH- also can be base. Example: NH3 is common base. When added to water, it accepts and H+ ion and produces OH- ion. NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) - NH3 is a weak electrolyte. Because only small fraction of NH3 forms NH4+ and OH- ions.
    • ACID-BASE REACTIONS ACID BASE Substances that Substances that increase the increase the Arrhenius concentration of H+ concentration of when dissolved in OH− when dissolved water in waterBrønsted and Proton donors Proton acceptors Lowry
    • ACID-BASE REACTIONS Neutralization Reactions and Salts• Properties of acidic solutions are different from the basic one.• Neutralization reaction occurs when a solution of acid andbase are mixed.• The products of the reaction have none of the characteristicproperties of either the acidic or basic solution.• Neutralization process will produce salt and water asproducts.• CH3COOH (aq) + NaOH (aq) →CH3COONa (aq) + H2O (l)
    • ACID-BASE REACTIONS Neutralization Reactions and Salts• When a strong acid reacts with a strong base, the net ionicequation is: HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq) → Na+ (aq) + Cl- (aq) + H2O (l) H+ (aq) + OH- (aq) → H2O (l)
    • ACID-BASE REACTIONS Acid-Base Reactions with Gas Formation• Some metathesis reactions do not give the product expected.• When a carbonate or bicarbonate reacts with an acid, first givesthe carbonic acid (H2CO3) and salt. HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2CO3• Carbonic acid (H2CO3) then decomposed and the productsbecome salt, carbon dioxide, and water H2CO3(aq)  H2O(l) + CO2(g)
    • ACID-BASE REACTIONS• The overall reaction is summarized:Molecular eq: NaHCO3 (aq) + HBr (aq) →NaBr (aq) + CO2 (g) + H2O (l)Total ionic eq: Na+(aq) + HCO3-(aq) + H+(aq) + Br-(aq)  Na+(aq) + Br-(aq) + CO2 (g) + H2O (l)Net ionic eq: HCO3-(aq) + H+(aq) +  CO2 (g) + H2O (l)
    • OXIDATION-REDUCTION (REDOX) REACTIONS• Which electrons are transferred between reactants.• Corrosion: The conversion of a metal into a metal compound by a reaction between the metal and its enviroment.• When a metal corrodes, it loses electrons and form cations.• Example: Calcium vigorously attacked by acids to form calcium ion: Ca(s) + 2H+ (aq)  Ca2+ (aq) + H2 (g) • Oxidation: Loss of electron by a substance • The term oxidation is used because the 1st reactions of this sort to be studied thoroughly were reactions with oxygen. • Many metals react directly with O2 in air to form metal oxide.
    • OXIDATION-REDUCTION (REDOX) REACTIONS• Example: 2Ca(s) + O2(g)  2CaO(s)• Ca losses e- to produce CaO which is ionic compound.• Ca is oxidized (lose e-) while O2 gain e- and transformed to O2- ion. O2 is said to be reduced.• Reduction: The gain of electron by a substance.• The oxidation and reduction process in in-situ process.
    • OXIDATION-REDUCTION (REDOX) REACTIONS Oxidation Numbers• To determine if an oxidation-reduction reaction has occurred,we assign an oxidation number (also called oxidation states) toeach element in a neutral compound or charged entity.• The oxidation numbers of certain atoms change in anoxidation-reduction reaction.• Oxidation occurs when the oxidation number increase, whilereduction occurs when the oxidation number decreases.
    • OXIDATION-REDUCTION (REDOX) REACTIONS• We use the following rules for assigning the oxidationnumbers: • Elements in their elemental form have an oxidation number of 0. • The oxidation number of a monatomic ion is the same as its charge. (e.g; K+ has an oxidation number of +1) – We write the sign +/- 1st before write the number. • Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. • Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1. • Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.
    • OXIDATION-REDUCTION (REDOX) REACTIONS • Fluorine always has an oxidation number of −1. • The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.• The sum of the oxidation numbers in a neutral compoundis 0.• The sum of the oxidation numbers in a polyatomic ion isthe charge on the ion.
    • OXIDATION-REDUCTION (REDOX) REACTIONSEXAMPLE:Determine the oxidation number (O.N.) of each element inthese compounds:(a) zinc chloride (b) sulfur trioxide (c) nitric acid
    • OXIDATION-REDUCTION (REDOX) REACTIONSEXERCISE:Which reactions are redox reactions: (a) CaO(s) + CO2(g) CaCO3(s) (b) 4 KNO3(s) 2 K2O(s) + 2 N2(g) + 5 O2(g) (c) NaHSO4(aq) + NaOH(aq) Na2SO4(aq) + H2O(l)
    • OXIDATION-REDUCTION (REDOX) REACTIONS Oxidation of Metals by Acids and Salts• The reaction of an acid or metal salt conforms to the generalfollowing pattern: A + BX  AX + B Zn(s) + 2 HBr(aq)  ZnBr2(aq) + H2(g) • These reactions are called displacement reactions because the ionic solution is displaced or replaced through oxidation of an element. • Many metals undergo displacement reactions with an acid to produce salt and H2 gas.
    • OXIDATION-REDUCTION (REDOX) REACTIONS • In displacement reactions, ions oxidize an element. • The ions, then, are reduced. • To show that the redox reaction have occurred, the oxidation number is shown below: Mg (s) + 2HCl(aq)  MgCl2(aq) +H2(g) 0 +1-1 +2 -1 0
    • OXIDATION-REDUCTION (REDOX) REACTIONS• Metal can also be oxidized by aqueous solutions of varioussalt.• Example: Molecular Eq: Fe(s) + Ni(NO3)2(aq)  Fe(NO3)2(aq) + Ni(s) Net ionic Eq: Fe(s) + Ni2+ (aq)  Fe2+ (aq) + Ni (s)• The oxidation of Fe to form Fe2+ in this reaction isaccompanied by the reduction of Ni2+ to Ni.• Whenever a substance is oxidized, some other substancemust be reduced.
    • OXIDATION-REDUCTION (REDOX) REACTIONS The Activity Series• A list of metals arranged in order of decreasing ease of oxidationis called: Activity series• TOP: Most easily oxidized (1A) : Active metals• BOTTOM: Stable (8B n 1B) : Noble metal• Can be used to predict theoutcome of reactions betweenmetals and either metal salts oracids.• Any metal on the list can beoxidized the ions of elementsbelow it.
    • OXIDATION-REDUCTION (REDOX) REACTIONS• For example: Copper is above silver in the series. Thus, coppermetal will be oxidized by silver ions: Cu (s) + 2 Ag+ (aq) → Cu2+ (aq) + 2 Ag (s)
    • OXIDATION-REDUCTION (REDOX) REACTIONS• Only those metals above hydrogen in the activity series are ableto react with acids to form H2. • Example: Ni(s) + 2HCl(aq)  NiCl2 (aq) + H2 (g)• Because elements below hydrogen in the activity series are notoxidized by H+, Cu doesn’t react with HCl(aq). But does react withHNO3.• This reaction however, is not a simple oxidation of Cu by the H +ions of the acid. Instead, the metal is oxidized to Cu2+ by thenitrate ion of the acid, accompanied by the formation of brownNO2(g): Cu(s) + 4 HNO3(aq)  Cu(NO3)2 (aq) + 2H2O (l) + 2 NO2(g)
    • CONCENTRATIONS OF SOLUTIONS Molarity• Two solutions can contain the same compounds but be quitedifferent because the proportions of those compounds aredifferent.• Molarity is one way to measure the concentration of a solution moles of solute Molarity (M) = volume of solution in liters• A molar of solutions (1 M) contains 1 mol of solutes in every 1 Lsolution. What is the difference between 0.5 mol of H2SO4 and 0.5 Molar of H2SO4?
    • CONCENTRATIONS OF SOLUTIONS Laboratory preparation of molar solutionsA•Weigh the solid needed. C Add solvent until the solution•Transfer the solid to a reaches its final volume.volumetric flask that containsabout half the final volume ofsolvent. B Dissolve the solid thoroughly by swirling.
    • CONCENTRATIONS OF SOLUTIONSEXAMPLE:What is the molarity of an aqueous solution that contains 0.715mol of glycine (H2NCH2COOH) in 495 mL? EXAMPLE: How many grams of solute are in 1.75 L of 0.460 M sodium monohydrogen phosphate buffer solution?
    • CONCENTRATIONS OF SOLUTIONSEXERCISE:An experiment calls for the addition to a reaction vessel of0.184 g of sodium hydroxide. How many mililiters of 0.150M NaOH should be added.
    • CONCENTRATIONS OF SOLUTIONSExpressing the Concentration of an Electrolyte• When ionic compound dissolves, the relative concentrations ofthe ions introduced into the solution depend on the chemicalformula of the compound.• Example: 1 M solution of NaCl is 1 M in Na+ ions and 1 M in Cl- ions. : 1 M of Na2SO4 is 2 M of H+ ions and 1 M in SO42- ions. • The conc of an electrolyte solution can be specified either in term of the compound used to make the solution (1 M of Na2SO4) or in terms of the ions that the solution contains (2M Na+ and 1M SO42-)
    • CONCENTRATIONS OF SOLUTIONS Dilution• Stock solutions: Solutions that are routinely used in the lab are often purchased or prepared in concentrated form.• Dilution: The process of obtained lower concentration from high concentration solutions by adding water.• One can also dilute a more concentrated solution by – Using a pipet to deliver a volume of the solution to a new volumetric flask, and – Adding solvent to the line on the neck of the new flask.
    • CONCENTRATIONS OF SOLUTIONS• The molarity of the new solution can be determined from the equation Moles solute before dilution = moles solute after dilution Mc × Vc = Md × Vd, where Mc and Md are the molarity of the concentrated and dilute solutions, respectively, and Vc and Vd are the volumes of the two solutions.• The concentration can be in L or mL as long as the unit is the same in both side. (1.00 M)(Vconc) = (0.100 M)(250 mL) Vconc = 25 mL (1.00 M)(Vconc) = (0.100 M)(0.25 mL) Vconc = 0.025 mL
    • SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS• Problem solving procedure : Outline of the procedure used tosolve stoichiometry problems that involve measured (lab) unitsof mass, solution conc (Molarity) or volume.
    • SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSISEXAMPLE:How many grams of NaOH are needed to neutralize 20.0 mLof 0.150 M H2SO4 solution?
    • SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS Titration• To determine the concentration of particular solute in asolution : Titration.• Titration: Acid-base, precipitation or redox reaction.• Equivalence point: The point which stoichiometry equivalentquantities are brought together.• Indicators: - Used to determine the end point of the reaction. - Color changes of the indicator showing that reaction had occur. - Example: Phenolphthalein: Acidic- Colorless, Basic- pink.
    • SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSISEXAMPLE:• What volume of 0.128 M HCl is needed to neutralize 2.87 gof Mg(OH)2?
    • SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSISEXERCISEA solution of 100 mL of 0.200 M KOH is mixed with a solutionof 200.00 mL of 0.150 M NiSO4.a) Write the balanced chemical equation for the reaction.b) What precipitate forms?c) What is the limiting reactant?d) How many grams of the precipitate form?
    • I am only one;But still I am one;I cannot doeverything;But still I can dosomething.I will not refuse to dosomething I can do. -Helen Keller-