AQUEOUS REACTIONS

CHAPTER 4 :
AQUEOUS REACTIONS
AND SOLUTION
STOICHIOMETRY

QBA Miguel A. Castro Ramírez

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.

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.

Ionic Compounds in Water
• Water is a very effective solvent of ionic compounds
• Electrically neutral molecule, but one end of the molecule (O
atom) rich in electron and has a partial negative charge (δ-)
• The other end (H atoms), has a partial positive charge (δ+)

• Cation (+) are attractive to (δ-) and anions (-) are attractive to
(δ+)
• As an ionic compound dissolves, the ions become surrounded by
H2O molecules.
• The ions are said to be solvated.
• The solvation process helps stabilize ions in solution and prevents
cations and anions from recombining.
• The ions and their shells (water molecules) are free to move, the
ions become dispersed uniformly throughout the solution.

Molecular Compounds in Water

• Dissolves in water, solution usually consist of intact molecules
dispersed throughout the solution.
• So, molecular compounds are non-electrolytes.
• Example: Table sugar (sucrose) and methanol.
• A few molecular substances have aqueous solutions that contain
ions.
• Example: HCl – Ionizes and dissociates into H+(aq) and Cl-(aq)
ions.

Strong and Weak Electrolytes

• Two categories: Weak and strong- Differ in the extend to which
they conduct electricity.
• Strong electrolytes: Dissociates completely when dissolved in
water.
• All soluble ionic compounds are electrolytes Eg: NaCl
• Strong acid, strong base and soluble ionic compounds.
• Weak electrolytes: Only dissociates partially when dissolved in
water.
• E.g: CH3COOH: Most of the solute is present as CH3COOH
molecules. Only small fraction of the acid is present as H +(aq) and
CH3COO-(aq) ions.


• For strong electrolyte: Single arrow represent the ionization of
strong 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 that
result in the formation of an insoluble
product.
• Precipitate: An insoluble solid formed by
a reaction in solution.
Pb(NO3)2 (aq) + 2KI (aq)  PbI2 (s) +
2KNO3 (aq)
• To predict whether certain combinations
of ions form insoluble compounds, we
must consider some guidelines concerning
the solubilities of common ionic
compounds.


Solubility Guidelines for Ionic Compounds

• Solubility: The amount of substance that can be dissolved in a
given quantity of solvent at the given temperature.
• Insoluble: The attraction between the oppositely charged ions in
the solid is too great for the water molecules to separate the ions
to any significant extend- substance remains undissolved.
• No rules based on physical properties to predict the solubility of
compound.

• Based on the experimental observation only.
• All common ionic compounds of alkali metals ions and the NH4+
are soluble in water.

• 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)


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.

Writing Equations for Aqueous Ionic Reactions
The molecular equation

Shows all of the reactants and products as intact, undissociated
compounds.
AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)

The total ionic equation

Shows all of the soluble ionic substances dissociated into ions.
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)  AgCl (s) + K+ (aq) +
NO3(aq)

Writing Equations for Aqueous Ionic Reactions
The net ionic equation
• Eliminates the spectator ions and shows the actual chemical
change taking place.
• Those things that didn’t change (and were deleted from the
net 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)


Writing Net Ionic Equations

1. 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.


EXAMPLE

Predict whether a reaction occurs when each of the following
pairs of solutions are mixed. If a reaction does occur, write
balanced molecular, total ionic, and net ionic equations, and
identify the spectator ions.
(a) potassium fluoride(aq) + strontium nitrate(aq) 
(b) ammonium perchlorate(aq) + sodium bromide(aq) 

Using Molecular Depictions to Understand a Precipitation
Reaction
Consider the molecular views of the reactants for a precipitation
reaction.

(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 balanced
molecular, total ionic, and net ionic equations.

ACID-BASE REACTIONS

• Acids: Substance that ionize in aqueous
solutions to form H+ ions, thereby increasing
the concentration of H+ (aq) ions.
• Acids : Called as proton donors.
•Molecules from different acids can ionize to
form 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 water
Brønsted 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 and
base are mixed.
• The products of the reaction have none of the characteristic
properties of either the acidic or basic solution.
• Neutralization process will produce salt and water as
products.
• 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 ionic
equation 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 gives
the carbonic acid (H2CO3) and salt.
HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2CO3
• Carbonic acid (H2CO3) then decomposed and the products
become 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.


• 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 Numbers

• To determine if an oxidation-reduction reaction has occurred,
we assign an oxidation number (also called oxidation states) to
each element in a neutral compound or charged entity.
• The oxidation numbers of certain atoms change in an
oxidation-reduction reaction.
• Oxidation occurs when the oxidation number increase, while
reduction occurs when the oxidation number decreases.


• We use the following rules for assigning the oxidation
numbers:
• 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.


• 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 compound
is 0.
• The sum of the oxidation numbers in a polyatomic ion is
the charge on the ion.


EXAMPLE:
Determine the oxidation number (O.N.) of each element in
these compounds:
(a) zinc chloride (b) sulfur trioxide (c) nitric acid

EXERCISE:
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 of Metals by Acids and Salts

• The reaction of an acid or metal salt conforms to the general
following 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.


• 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


• Metal can also be oxidized by aqueous solutions of various
salt.
• 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 is
accompanied by the reduction of Ni2+ to Ni.
• Whenever a substance is oxidized, some other substance
must be reduced.

The Activity Series
• A list of metals arranged in order of decreasing ease of oxidation
is called: Activity series
• TOP: Most easily oxidized (1A)
: Active metals
• BOTTOM: Stable (8B n 1B)
: Noble metal
• Can be used to predict the
outcome of reactions between
metals and either metal salts or
acids.
• Any metal on the list can be
oxidized the ions of elements
below it.

• For example: Copper is above silver in the series. Thus, copper
metal will be oxidized by silver ions:
Cu (s) + 2 Ag+ (aq) → Cu2+ (aq) + 2 Ag (s)


• Only those metals above hydrogen in the activity series are able
to react with acids to form H2.
• Example: Ni(s) + 2HCl(aq)  NiCl2 (aq) + H2 (g)
• Because elements below hydrogen in the activity series are not
oxidized by H+, Cu doesn’t react with HCl(aq). But does react with
HNO3.
• 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 the
nitrate ion of the acid, accompanied by the formation of brown
NO2(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 quite
different because the proportions of those compounds are
different.
• 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 L
solution.

What is the difference between 0.5 mol of H2SO4 and 0.5
Molar of H2SO4?


Laboratory preparation of molar solutions

A
•Weigh the solid needed.
C Add solvent until the solution
•Transfer the solid to a
reaches its final volume.
volumetric flask that contains
about half the final volume of
solvent.
B Dissolve the solid
thoroughly by swirling.


EXAMPLE:
What is the molarity of an aqueous solution that contains 0.715
mol 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?


EXERCISE:
An experiment calls for the addition to a reaction vessel of
0.184 g of sodium hydroxide. How many mililiters of 0.150
M NaOH should be added.

Expressing the Concentration of an
Electrolyte

• When ionic compound dissolves, the relative concentrations of
the ions introduced into the solution depend on the chemical
formula 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-)

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.


• 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 to
solve stoichiometry problems that involve measured (lab) units
of mass, solution conc (Molarity) or volume.

ANALYSIS
EXAMPLE:
How many grams of NaOH are needed to neutralize 20.0 mL
of 0.150 M H2SO4 solution?

ANALYSIS
Titration
• To determine the concentration of particular solute in a
solution : Titration.
• Titration: Acid-base, precipitation or redox reaction.
• Equivalence point: The point which stoichiometry equivalent
quantities 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.

ANALYSIS

ANALYSIS
EXAMPLE:
• What volume of 0.128 M HCl is needed to neutralize 2.87 g
of Mg(OH)2?

ANALYSIS

EXERCISE

A solution of 100 mL of 0.200 M KOH is mixed with a solution
of 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 do
everything;
But still I can do
something.
I will not refuse to do
something I can do.

-Helen Keller-

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