LOGO
Aqueous Solutions
International University of Sarajevo
Course lecturer :
Jasmin Šutković
1th April 2015
Contents
1.
2.
3.
4.
5.
6.
7.
8.
International University of Sarajevo
Aqueous solutions
Solution Concentrations
Stoichiometry – Reactions of Solutions
Ionic equations
Participation reactions
Acid- Base reactions
Acid RAIN
Oxidation – Reduction reactions in aqueous
solutions
9. Titration
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).
Polar substances cont…
 Due to the asymmetric charge distribution in the water molecule,
adjacent water molecules are held together by attractive
electrostatic interactions.
 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.
Example : NaCl
NaCls
Ionic solvent !
H2O
Na+ (aq) + Cl-(aq)
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.
Strong Electrolytes
 Exists in solution completely or almost
completely as ions
 All ionic compounds and a few molecular
compounds. (Ex: Strong Acids)
HCl( aq)  H


( aq)
 Cl


( aq)
NaCl( s)  Na (aq)  Cl (aq)
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.
Examples
KF
Na3PO4
NH3
CH3CH2OH
HCl
NO2
HC2H3O2
CH4
NH4Cl
CH3Cl
strong electrolyte
strong electrolyte
weak electrolyte
nonelectrolyte
strong electrolyte
nonelectrolyte
weak electrolyte
nonelectrolyte
strong electrolyte
nonelectrolyte
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!
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 among volume, molarity, and moles is expressed as
Number of moles (n)= V(l) x M(mol/l)
See Example 4.2 (page 150)
Example 4.3,
(page 151)
In the figure below we have a solution that contains 10g of CoCl2 x 2H2O
(cobalt chloride dihydrate),and with a proper amound 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 / 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 CoCl2 x 2H2O is
165.87g/mol
n= m / Mr = 10g / 165.84(g/mol) = 0.063mol
M= 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 !
 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=VXM
Exercise 4.7 (page 158)
What mass of lanthanium(III) oxalate nonahydrate,
La2 (C2O4)3 x 9H2O, can be obtained from 650ml
of 0.0170 M aqueous solution of LaCl3 by adding
sodium oxalate ?
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 4.8 (page 159)
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
Pb( NO3 )2  2KI( aq)  PbI2( s )  2KNO3aq) 
2. Complete ionic equation — shows all of the substances
present in the form in which they actually exist in solution
2





Pb( aq)  2 NO
 2 K ( aq)  2 I ( aq)  PbI 2( s )  2 K ( aq)  2 NO3( aq)
( aq)
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
( aq)
Pb
 2I

( aq)
 PbI 2( s )
Typed 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
 Are a subclass of exchange reactions that occur
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 chloride
HBr
Hydrogen bromide
HI
Hydrogen iodide
HClO4 Perochloric acid
HNO3 Nitric acid
H2SO4 Sulfuric acid
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
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
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
Base properties
• Bitter taste
• Turns red litmus
red
• pH < 7
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 is called the hydronium ion.
 Substances that can behave as both an acid and a base are said to be
amphoteric.
Neutralization reactions
Acid + Base
Neutralization
HBr(aq) + NaOH(aq)
H2O(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.
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
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 at
25ºC; 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 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 (SO42-) 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 objects
3. Decreases the pH of natural waters
4. Biological effects
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)
H3AsO4(aq) + H2O(l)
Lets do it together.....
NO(g) +
Balancing REDOX equation
with Oxidation states method
+1+5 -2
+1 +2 -2
HNO3 + H3AsO3
+2 -2
+1
+4 -2
+1 -2
NO + H3AsO4+ H2O
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
N +5 to +2
Net Change = +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 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.
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)
Thanks …