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International University of Sarajevo

Lecture 4 :Aqueous Solutions

Course lecturer : Jasmin Šutković11th March 2014

Chemistry - SPRING 2014

Contents International University of Sarajevo

1. Aqueous solutions 2. Solution Concentrations 3. Stoichiometry – Reactions of Solutions4. Ionic equations 5. Participation reactions 6. Acid- Base reactions 7. Acid RAIN 8. Oxidation – Reduction reactions in aqueous

solutions

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.

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.

H2O - 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).

• 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.

Polar substances cont…

Example : NaCl

NaCls Na+(aq) + Cl-(aq)

Ionic solvent !

H2O

Lets take a closer look

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.

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

Weak electrolytes

Molecular compounds that produce a small concentration of ions when dissolved in H2O

Weak electrolytes only ionize to a small extent so that just a (relatively) few of its molecules produce ions.

For example :

CH3COOH + H2O ⇌ H3O+ + CH3COO-

NH3 + H2O ⇌ NH4+ + OH-

Strong Electrolytes

Exists in solution completely or almost completely as ions All ionic compounds and a few molecular

compounds. (Ex: Strong Acids)

)()()( aqaqaq ClHHCl −+ +→

)()()( aqaqs ClNaNaCl −+ +→

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

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.

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.

Examples

KFNa3PO4NH3CH3CH2OHHClNO2HC2H3O2CH4NH4ClCH3Cl

strong electrolyte

strong electrolyte

strong electrolyte

strong electrolyte

weak electrolytenonelectrolyte

nonelectrolyte

nonelectrolyte

nonelectrolyte

weak electrolyte

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!

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 among volume, molarity, and moles is expressed as

Solution Concentrations

See Example 2 (page 239)

Number of moles (n)= V(l) x M(mol/l)

Example 4.3

In the figure below we have a solution that contains 10g of CoCl2 x 2H2O(cobalt chloride dihydrate),and with a proper amount of ethanol it makes exactly 500ml of solution. WHAT IS IT MOLAR CONCENTRATION ?

We are looking for M ( We are given V=500mL and m=10g )

Formula = M = n / Vn= ?, we calculate n by dividing mass of compound by its molar mass

or molecular mass( Mr ).

n= m / MrSo the molar mass (Mr) of CoCl2 x 2H2O is

165.87g/moln= m / Mr = 10g / 165.84(g/mol) = 0.063molM= n / V = 0.063mol / 0.500L = 0.121 M

Concentration of CoCl2 x 2H2O

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

Example 5 (page 244)

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.

Example

A 50.6g sample of Mg(OH)2 is reacted with 45g of HCL according this reaction:

Mg(OH)2 + 2 HCl --> MgCl2 + 2 H2O

a) What is the theoretical yield of MgCl2? b) What is the limiting reactant ?

More complicated example 8 (page 253)

A typical Breathalyzer contains 3mL of 0.25mg/mL solution of K2Cr2O7 in 50% H2SO4 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 Cr6+ to Cr3+ ?

SOLVED PROBLEM – CHECK THE BOOK

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

2. Complete ionic equation — shows all of the substances present in the form in which they actually exist in solution

( )aqsaq KNOPbIKINOPb 3)(2)(23 22)( +→+

−+++→−+++−++)(32)(2)(2)(2)(2

)(22

)( aqNOaqKsPbIaqIaqKaq

NOaqPb

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

)(2)(2

)( 2 saqaq PbIIPb →+ −+

Types of chemical reactions

Three common kinds of reactions that occur in aqueous solution are

1. precipitation,2. acid-base,3. oxidation-reduction.

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 )

6. Acid-Base Reactions

Acids:

• Ionize in H2O, causes increase in H+ ions.• H+ ions are bare protons.• Acids are proton donor• Reacts with some metals to produce H2• Dissolves carbonate salts, releasing CO2

Acids that can only yield one H+ per molecule upon ionization.

HCl H+ + Cl-

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.

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.

Examples of strong acids

For strong acids, try to remember them, there are 6 :

Strong acids:HCl Hydrogen chlorideHBr Hydrogen bromideHI Hydrogen iodide HClO4 Perochloric acidHNO3 Nitric acidH2SO4 Sulfuric acid

Bases

Bases:

Substances that increase the OH- when added to water. (NaOH)

Strong bases:Any groups in 1A or 2Aelements with OHelements with O2 elements and NH2

Definitions of Acidsand 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

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.

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.

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

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.

Acid + Base Neutralization

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.

Neutralization reactions

HBr(aq) + NaOH(aq) H2O(l) + NaBr(aq)

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 baseb. Reaction of a strong acid with a weak basec. 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

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!!

Example with pure liquid water

Pure liquid water contains low but measurable concentrations of H3O+ and OH- ions produced via auto-ionization reaction in which water acts in the same time as an acid and a base .

H2O (aq) + H2O (l) H3O+(aq) + OH-

(aq)

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 at25ºC; the pH of pure water is – log [1.0 x 10-7] = 7.00.• 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+].

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?

SO4 2- (sulfate) and NO3

- (nitrate) level increase due to the industry production of fossil oils .

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 (SO4

2-) and nitrate (NO3-) 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 objects3. Decreases the pH of natural waters4. Biological effects

Acid Rain cont..

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

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.

HNO3(aq) + H3AsO3(aq) NO(g) + H3AsO4(aq) + H2O(l)

Lets do it together.....

Balancing REDOX equation with Oxidation states method

HNO3 + H3AsO3 NO + H3AsO4+ H2O+1 -2 +1 -2 -2 +1 -2 +1 -2+5 +2 +2 +4

1.Try to balance the atoms by inspection ,but O and H are hard to balance that way2. 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 = -3

4. 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.

Balancing REDOX equation with Oxidation states methodAs 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.

2HNO3(aq) + 3H3AsO3(aq) NO(g) + H3AsO4(aq) + H2O(l)

6. Balance the rest of the equation by inspection.

2HNO3(aq) + 3H3AsO3(aq) --> 2NO(g) + 3H3AsO4(aq) + H2O(l)

Quiz 1 – 18th 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 :229 – 297

ADDITIONAL INFO