Lecture 4 :Aqueous Solutions

Transcription

1 LOGO Lecture 4 :Aqueous Solutions International University of Sarajevo Chemistry - SPRING 2014 Course lecturer : Jasmin Šutković 11 th March 2014

2 Contents International University of Sarajevo 1. Aqueous solutions 2. Solution Concentrations 3. Stoichiometry Reactions of Solutions 4. Ionic equations 5. Participation reactions 6. Acid- Base reactions 7. Acid RAIN 8. Oxidation Reduction reactions in aqueous solutions

3 1. Aqueous solutions Reminder : Solution is a homogenous mixture where the substances are in smaller amounts, called SOLUTES ( the substance being dissolved) and if these substances are found in greater amount they are called SOLVELNT ( the substance doing the dissolving) In Aqueous solution the solvent is WATER No-aqueous solution any substance other than water is the solvent Water is essential for life and makes up about 70% of the mass of the human body. Many of the chemical reactions that are essential for life depend on the interaction of water molecules with dissolved compounds.

4 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 and have a partial positive charge, indicated by the symbol δ +. The oxygen atom is more electron-rich and have a partial negative charge, indicated by the symbol δ -. Unequal distribution of charge (sharing of electrons) creates a polar bond.

5 H 2 O - Water A water molecule, a commonly-used example of polarity. The two charges are present with a negative charge in the middle (red shade), and a positive charge at the ends (blue shade).

6 Polar substances cont The water molecules are held together by attractive electrostatic interactions( due to the asymmetric charge ) Energy is needed to overcome these electrostatic interactions Unequal charge distribution in polar liquids, like water, makes them good solvents for ionic compounds. When an ionic solid is dissolves in water, the partially negatively charged oxygen atoms in the water surround the cations of the ionic solid, and the partially positively charged hydrogen atoms in water surround the anions. Individual cations and anions are called hydrated ions.

7 Example : NaCl NaCl s Ionic solvent! H2O Na + (aq) + Cl - (aq)

8 Lets take a closer look

9 Binding of NaCl with H2O Unequal sharing of electrons leads to partial positive and negative charges in a water molecule. These charges attract the ions which causes dissociation of the ionic compound in water.

10 Electrolytes Electrolyte is 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, and 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. More IONS= better electrolyte Less IONS = bad electrolytes

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12 Weak electrolytes Molecular compounds that produce a small concentration of ions when dissolved in H 2 O Weak electrolytes only ionize to a small extent so that just a (relatively) few of its molecules produce ions. For example : CH 3 COOH + H 2 O H 3 O + + CH 3 COO - NH 3 + H 2 O NH 4+ + OH -

13 Strong Electrolytes Exists in solution completely or almost completely as ions All ionic compounds and a few molecular compounds. (Ex: Strong Acids) HCl NaCl H + ( aq) ( aq) + Cl ( aq) ( s) Na + ( aq) + Cl ( aq)

14 7 strong acids remember! HCl -hydrochloric acid, HNO3 -nitric acid, H2SO4- sulfuric acid, HBr- hydrobromic acid, HI- hydroiodic acid, HClO3 - chloric acid HClO4 - Perchloric acid

15 Weak VS Strong Electrolytes The main difference between strong and weak electrolytes is the amount of electricity that is allowed to flow. It is the number of ions in solution that determines the amount of electricity that can flow through a solution.

16 Six Steps for Categorizing Electrolytes So how do we categorize compounds based on their formula? One practical method is outlined below: Step 1 - Is it one of the seven strong acids? Step 2 - Is it of the form Metal(OH) n? Then it's a strong base. Step 3 - Is it of the form Metal(X) n? Then it's a salt. Step 4 - Does it's formula start with 'H'? It's probably a weak acid. Step 5 - Does it have a nitrogen atom? It may be a weak base. Step 6None of those? Call it a nonelectrolyte.

17 Examples KF Na 3 PO 4 NH 3 CH 3 CH 2 OH HCl NO 2 HC 2 H 3 O 2 CH 4 NH 4 Cl CH 3 Cl strong electrolyte strong electrolyte weak electrolyte nonelectrolyte strong electrolyte nonelectrolyte weak electrolyte nonelectrolyte strong electrolyte nonelectrolyte

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19 2. Solution Concentrations Concentration of a solution describes the quantity of a solute that is contained in the solvent or solution! Knowing the concentration of solutes is important in controlling the Stoichiometry of reactant for reactions that occur in solution!

20 Solution Concentrations Molarity (M) Most common unit of concentration 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 Number among volume, of moles (n)= molarity, V(l) x and M(mol/l) moles is expressed as See Example 2 (page 239)

21 Example 4.3 In the figure below we have a solution that contains 10g of CoCl 2 x 2H 2 O (cobalt chloride dihydrate),and with a proper amount of ethanol it makes exactly 500ml of solution. WHAT IS IT MOLAR CONCENTRATION?

22 We are looking for M ( We are given V=500mL and m=10g ) Formula = M = n / V n=?, we calculate n by dividing mass of compound by its molar mass or molecular mass( Mr ). n= m / Mr So the molar mass (Mr) of CoCl 2 x 2H 2 O is g/mol n= m / Mr = 10g / (g/mol) = 0.063mol M= n / V = 0.063mol / 0.500L = M Concentration of CoCl 2 x 2H 2 O

23 3. Stoichiometry of Reactions in Solution Before everything we have to do balancing of the chemical equation! The coefficients in the balanced chemical equation indicate the number of moles of each reactant that is needed and the number of moles of each product that can be produced. It doesn t matter if you are dealing with volumes of solutions of reactants or masses of reactants. Calculating Moles from Volume Number of moles (n)= V(l) x M (mol/l) M = n / V, n = V X M

24 Example 5 (page 244)

25 Limiting Reactants in Solutions Are those reactants 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.

26 Example A 50.6g sample of Mg(OH)2 is reacted with 45g of HCL according this reaction: Mg(OH) HCl --> MgCl H 2 O a) What is the theoretical yield of MgCl 2? b) What is the limiting reactant?

27 More complicated example 8 (page 253) A typical Breathalyzer contains 3mL of 0.25mg/mL solution of K 2 Cr 2 O 7 in 50% H 2 SO 4 as well as a fixed concentration of AgNO3. How many grams of ethanol must be present in 52.5mL of persons breath to convert all of the Cr 6+ to Cr 3+? SOLVED PROBLEM CHECK THE BOOK

28 4. 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 un-dissociated form Pb( NO ) KNO3 3 2 KI( aq) PbI2( s) 2 ( aq) 2. Complete ionic equation shows all of the substances present in the form in which they actually exist in solution Pb NO + 2K + + 2I PbI + 2K + + 2NO ( aq) ( aq) ( aq) 2( s) ( aq) 3( aq) ( aq)

29 Ionic equations cont 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 Pb I PbI ( aq) ( aq) 2( s)

30 Types of chemical reactions Three common kinds of reactions that occur in aqueous solution are 1. precipitation, 2. acid-base, 3. oxidation-reduction.

31 5. Precipitation reactions A reaction that yields an insoluble product, a precipitate, when two solutions are mixed Reaction occures 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! ( PREDICTING SOLUBILITIES - NOT NEEDED )

32 6. Acid-Base Reactions Acids: Ionize in H 2 O, causes increase in H + ions. H + ions are bare protons. Acids are proton donor Reacts with some metals to produce H 2 Dissolves carbonate salts, releasing CO 2 Acids that can only yield one H + per molecule upon ionization. HCl H + + Cl -

33 IONIZATION Ionization is the process of converting an atom or molecule into an ion by adding or removing charged particles such as electrons or ions.

34 What mean actually strong and weak acid/base? The terms "strong" and "weak" do NOT refer to the concentration of the acid or base, but instead, refer to whether the acid or base dissociates completely in water.

35 Examples of strong acids For strong acids, try to remember them, there are 6 : Strong acids: HCl Hydrogen chloride HBr Hydrogen bromide HI Hydrogen iodide HClO4 Perochloric acid HNO3 Nitric acid H2SO4 Sulfuric acid

36 Bases Bases: Substances that increase the OH - when added to water. (NaOH) Strong bases: Any groups in 1A or 2A elements with OH elements with O2 elements and NH2

37 Definitions of Acids and Bases Brønsted Lowry definition of acids and bases A more general definition of acids and bases An acid is any substance that can donate a proton. A base is any substance that can accept a proton. Not restricted to aqueous solutions

38 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.

39 Strengths of Acids and Bases Strong acids react essentially completely with water to give H + and the corresponding anion. Strong bases dissociate essentially completely in water to give OH and the corresponding cation. Both strong acids and strong bases are strong electrolytes.

40 Some Properties of Acids and Bases Acid Properties Sour taste Turn blue litmus red ph < 7 Base properties Bitter taste Turns red litmus blue ph >7 slippery

41 The Hydronium Ion When a strong acid dissolves in water, the proton that is released is transferred to a water molecule that acts as a proton acceptor or base, the Resulting molecule is H3O+ ion, also called as hydronium ion. Substances that can behave as both an acid and a base are said to be amphoteric.

42 Neutralization reactions Acid + Base Neutralization HBr (aq) + NaOH (aq) H 2 O(l) + NaBr (aq) Products of a neutralization reaction have none of the properties of an acid or a base. An acid reacts with a metal hydroxide to form a salt plus water.

43 Neutralization reactions cont.. A reaction in which an acid and a base react 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 a. Reaction of any strong acid with any strong base b. Reaction of a strong acid with a weak base c. 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

44 The ph Scale It is one of the main factors that affects the chemical reaction that occur in dilute solutions. It is a convenient way to express the hydrogen ION (H+) concentration of a solution and enables as to understand if a solution is an acid or base!!

45 Example with pure liquid water Pure liquid water contains low but measurable concentrations of H 3 O + and OH - ions produced via auto-ionization reaction in which water acts in the same time as an acid and a base. H 2 O (aq) + H 2 O (l) H 3 O + (aq) + OH - (aq)

46 The ph scale ph is defined as the negative base -10 logarithm of the hydrogen ion concentration ph = log [H + ] or [H + ] = 10 -ph Hydrogen ion concentration in pure water is 1 x 10-7 M at 25ºC; the ph of pure water is log [1.0 x 10-7 ] = ph decreases with increasing [H + ] adding an acid to pure water increases the hydrogen ion concentration and decreases the hydroxide ion concentration. Adding a base to pure water increases the hydroxide ion concentration and decreases the hydrogen ion concentration ph increases with decreasing [H + ].

47 7. The Chemistry of Acid Rain Acid rains have strong environmental impact! It accelerate the corrosion of metal objects and decreases the ph of natural water To understand acid rain we need to know what are acid base reactions Typical ph in US of acid rain is 4-5. Normal rain became acid as tomato juice and black coffee.. What is the source of this increase of acidity? SO 4 2- (sulfate) and NO 3- (nitrate) level increase due to the industry production of fossil oils.

48 Acid Rain cont.. Acid rain is rainfall whose ph is less than 5.6 due to dissolved carbon dioxide, which reacts with water to give the weak acid carbonic acid. Source of the increased acidity in rain due to the presence of large quantities of sulfate (SO 4 2- ) and nitrate (NO 3- ) ions, which come from nitrogen oxides and sulfur dioxide produced both by natural processes and by the combustion of fossil fuels These oxides react with oxygen and water to give nitric acid and sulfuric acid. Some damages caused by acid rain 1. Dissolves marble and limestone surfaces due to a classic acid-base reaction 2. Accelerates the corrosion of metal objects 3. Decreases the ph of natural waters 4. Biological effects

49 8. Oxidation Reduction reactions in solutions 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: 1. Oxidation states overall reaction is separated into an oxidation equation and a reduction equation 2. Half-reaction

50 Balancing REDOX equation with Oxidation states method Balance the following redox equation using the oxidation number method. Be sure to check that the atoms and the charge are balanced. HNO 3 (aq) + H 3 AsO 3 (aq) NO(g) + H 3 AsO 4 (aq) + H 2 O(l) Lets do it together...

51 Balancing REDOX equation with Oxidation states method HNO 3 + H 3 AsO 3 NO + H 3 AsO 4 + H 2 O 1.Try to balance the atoms by inspection,but O and H are hard to balance that way 2. Is this a redox reaction? The N atoms change from +5 to +2, so they are reduced. This information is enough to tell us that the reaction is redox. (The As atoms, which change from +3 to +5, are oxidized.) 3. Determine the net increase in oxidation number for the element that is oxidized and the net decrease in oxidation number for the element that is reduced. As +3 to +5 Net Change = +2 N +5 to +2 Net Change = Determine a ratio of oxidized to reduced atoms that would yield a net increase in oxidation number equal to the net decrease in oxidation number.

52 Balancing REDOX equation with Oxidation states method As atoms would yield a net increase in oxidation number of +6. (Six electrons would be lost by three arsenic atoms.) 2 N atoms would yield a net decrease of -6. (Two nitrogen atoms would gain six electrons.) Thus the ratio of As atoms to N atoms is 3:2. 5. To get the ratio identified in Step 4, add coefficients to the formulas which contain the elements whose oxidation number is changing. 2HNO 3 (aq) + 3H 3 AsO 3 (aq) NO(g) + H 3 AsO 4 (aq) + H 2 O(l) 6. Balance the rest of the equation by inspection. 2HNO 3 (aq) + 3H 3 AsO 3 (aq) --> 2NO(g) + 3H 3 AsO 4 (aq) + H 2 O(l)

53 ADDITIONAL INFO Quiz 1 18 th March at 11h ( 11h-11.50h) Follow the lecture slides( titles and subtitles) and check the book for more details and exercise the tutorial questions! Lecture 4 : Book pages :