Chapter 4 Reactions in Aqueous Solutions
4.1 General Properties of Aqueous Solutions
What is a solution?
A solution is a homogeneous mixture of substances. If there is much more of one
substance than any others, then that substance is referred to as the solvent, all other
substances in the mixture are called solutes.

An aqueous solution is one whose solvent is water. Many properties of an aqueous
solution are similar to the properties of water
1 (e.g. density, viscosity, melting point,
boiling point). Most properties will be slightly changed by the presence of solutes, and
some may be changed drastically, such as color and electrical conductivity.

The properties of the solutes are greatly changed and we define the aqueous state as
distinct from the solid, liquid and gaseous states of matter. The solute molecules no
longer interact with each other
2, as they are spread out in the solution. The solutes are
solvated by water molecules. That is, they are surrounded by solvent molecules, usually
with some degree of order, imposed by the polarity of the water molecules and the
polarity or charge of the solute.

Electrolytes
Some solutes remain molecular (e.g. sugar, ethanol). In order to be soluble in water they
must be somewhat polar molecules
3. These solutes have no effect on the electrical
properties of the solution and are classified as non-electrolytes.
Some solutes are dissociated into ions when they dissolve in water, (e.g. NaCl, HBr).
These solutes increase the electrical conductivity of water and are classified as
electrolytes. Ions increase the conductivity because electrical current is carried by the
highly mobile and electrically charged aqueous ions.

We can further subdivide electrolytes into strong and weak.

Strong Electrolytes: completely dissociate into ions when in aqueous solution.
The resulting solution is a strong conductor of electricity.
Weak Electrolytes: partially dissociate into ions when in aqueous solution.4
The resulting solution is a weak conductor of electricity.
Non-Electrolytes: do not dissociate into ions when in aqueous solution.
The resulting solution does not conduct electricity.

1. All ionic compounds are strong electrolytes.
2. Some acids and bases are strong electrolytes; many others are weak.
3. Most other molecules are non-electrolytes.

Hydration: The process in which an ion is surrounded by water molecules arranged in a specific manner.


4.2 Precipitation Reactions
A precipitation reaction occurs when water solutions of two different ionic compounds are mixed and an insoluble solid separates out of solution.
The precipitate is itself ionic; the cation comes from one solution and the anion from another. To predict the occurrence of these reactions, we must know which ionic substances are insoluble in water.

In a chemical sense precipitation is the formation of an insoluble product, a precipitate,
usually a solid, which separates from the solution. A common precipitate is an insoluble
salt.

Solubility = maximum amount of a solute that will dissolve in a given amount of a
solvent at a specific temperature. Solubility is typically given in grams of solute per
litre of solvent. (We may also express a salt’s ability to dissolve in terms of its solubility
product, K
sp.11
) Salts and other solutes fall into two (or three) qualitative categories:
soluble, (slightly soluble) and insoluble.

Soluble: ‘fair amount’ visibly dissolves when added to water.
Insoluble or slightly soluble: no noticeable amount dissolves.



Solubility Rules
Solubility Rules


Writing Balanced Reaction Equations
Molecular, Ionic and Net Ionic Equations; Spectator Ions
Molecular equation:
Pb(NO
3)2 (aq) + 2 NaI (aq) à PbI2 (s) + 2 NaNO3 (aq)
Ionic equation:
Pb
2+ (aq) + 2 NO3- (aq) + 2 Na+ (aq) + 2 I- (aq) à PbI2 (s) + 2 Na+ (aq) + 2 NO3- (aq)
Pb
2+ (aq) + 2 NO3- (aq) + 2 Na+ (aq) + 2 I- (aq) à PbI2 (s) + 2 Na+ (aq) + 2 NO3- (aq)
Net ionic equation:
Pb
2+ (aq) + 2 I- (aq) à PbI2 (s)
Spectator Ions:
Na
+ (aq) and NO3- (aq)
From the solubility rules you should be able to predict the products of a precipitation
reaction and write a balanced reaction equation.



4.3 Acid-Base Reactions
Acid-Base Concepts
• Svante Arrhenius proposed the following definitions for acids and
bases in 1884:
– An Arrhenius acid is a substance that ionizes in water to produce hydrogen
ions (H+).
HA (aq) → H+ (aq) + A- (aq)
– An Arrhenius base is a substance that ionizes in water to release hydroxide
ions (OH-).
MOH (aq) → OH- (aq) + M+ (aq)
• For example, HCl is an Arrhenius acid and NaOH is an Arrhenius
base.
• To determine whether you have an Arrhenius acid or base, write
the ionization or dissociation reaction of your compound.
• The Brønsted-Lowry definitions of acids and bases are broader
than the Arrhenius definitions.

A Brønsted-Lowry acid is a substance that donates a hydrogen ion
(H+) to any other substance.
– It is a proton donor.
– All Arrhenius acids are classified as acids by the B-L definition

Brønsted-Lowry acid
Brønsted-Lowry acid


A Brønsted-Lowry base is a substance that accepts a
hydrogen ion (H+).
– It is a proton acceptor (So no OH- required!).
– All Arrhenius bases are classified as bases by the B-L definition


Brønsted-Lowry base
Brønsted-Lowry base


Acids and bases have varying strengths.
• The strength of an acid is measured by the amount of H+ that are produced for each mole
of acid that dissolves
Ionization is the process where polar compounds separate into cations and anions in solution.
– The acid HCl ionizes into H+ and Cl– ions in solution.
• The strength of a base is measured by the degree of dissociation in solution.
Dissociation is the process where cations and anions in an ionic compound separate in solution.
– A formula unit of NaOH dissociates into Na+ and OH– ions in solution.


Common Acids and Bases
Common Acids and Bases


Neutralization Reactions
A neutralization reaction is the reaction of an acid and a base to produce a salt (BX) and water (HOH). HX + BOH
BX + HOH –This is a specialized “double replacement” reaction • The driving force of this reaction is the formation of the stable water molecule. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

4.4 Oxidation-Reduction Reactions
An Oxidation–Reduction (Redox) Reaction is a process in which one or more electrons are transferred between reaction partners.
–The formation of ionic solids from elements is always a Redox reaction.
• The driving force of this reaction is the decrease in electrical potential.
Mg(s) + I2(g) → MgI2(s)

• The transfer of electrons between or among reactants is called the oxidation or reduction of species depending on which way the electrons are flowing.

  • Reducing agent donates electrons
  • Oxidizing agent accepts electrons
  • An atom’s oxidation number, also called oxidation state, signifies the number of charges the atom would have in a molecule if electrons were transferred completely. They enable us to identify elements that are oxidized and reduced at a glance.
  • We use the following rules to assign oxidation numbers:
    • In free elements each atom has an oxidation number of zero.
    • For ions composed of only one atom, the oxidation number is equal to the charge on the ion.
    • The oxidation number of oxygen in most compounds is -2, but in hydrogen peroxide and peroxide ion, it is -1.
    • The oxidation number of hydrogen is +1, except when it is bonded to metals in binary compounds.
    • Fluorine has an oxidation number of -1 in ALL its compounds. Other halogens have negative oxidation numbers when they occur as halide ions in their compounds.
    • In a neutral molecule, the sum of the oxidation numbers of all the atoms must be 0. In a polyatomic ion, the sum of oxidation numbers of all the elements in the ion must be equal to the net charge of the ion.
    • Oxidation numbers do not have to be integers.
Types of Redox Reactions:
  • Combination Reactions: A + B à C
  • Decomposition Reactions: C à A + B
  • Displacement Reactions: A + BC à AC + B
    • Hydrogen Displacement: Ca + 2H2O à Ca(OH)2 + H2
    • Metal Displacement: V2O5 5Ca à 2V + 5CaO
    • Halogen Displacement: Cl2 + 2KBr à 2KCl + Br2
Disproportionation Reaction: an element in one oxidation state is simultaneously oxidized and reduced; Cl2 + 2OH- à ClO- + Cl- + H2O

4.5 Concentration of Solutions

The concentration of a solution is the amount of solute present in a given quantity of solvent or solution.
Molarity (M), or molar concentration, is the number of moles of solute in 1 liter (L) of solution
Dilution is the procedure for preparing a less concentrated solution from a more concentrated one.


--> Moles of solute before dilution = moles of solute after diliution
--> Moles of solute/liters of solution x volume of solution (in liters) = moles of solute
--> MV = moles of solute
--> MiVi = MfVf

Quantitative analysis is the determination of the amount or concentration of a substance in a sample.

The quantitative determination of a substance by precipitation followed by isolation and weighing of the precipitate is called gravimetric analysis.



4.6 Gravimetric Analysis

Gravimetric analysis is an analytical technique based on the measurement of mass.


4.7 Acid-Base Titrations

In a titration a solution of accurately known concentration, called standard solution, is added gradually to another solution of unknown concentration, until the chemical reaction between the two solutions is complete.
The equivalent point is the point at which the acid has completely reacted with or been neutralized by the base.
In acid-base titrations, indicators are substances that have distinctly different colors in acidic and basic media.


4.8 Redox Titrations
· Like acid-base titrations, redox titrations normally require an indicator that clearly changes color. In the presence of large amounts of reducing agent, the color of the indicator is characteristic of its reduced form; in large amounts of oxidizing agent, oxidized form; and at equivalence point, oxidized or reduced.
· Potassium dichromate and potassium permanganate are internal indicators in a redox titration because they have distinctly different colors in the oxidized and reduced forms.




LAB: Microscale Acid-Base Titration

Materials to be used
Provided materials
· 5mL unknown acid sample
· 25mL .01 sodium hydroxide solution
· 10 x 10 graph paper
Additional materials

· Safety materials: Goggles and aprons
· Plastic experimental wells, about 5mL deep
· Pipette
· Roll of litmus paper

Procedure
· Wearing goggles and aprons to protect against splashes and burns, transfer about 1mL of the known base into a clean well.
· Take an initial pH reading with a strip of litmus paper. Record this in a table.
· With a pipette, add one drop of the unknown acid solution to the well. Use another strip of litmus paper to test the new pH and record in the table.
· Continue to add a drop of acid and measure the pH, recording all data in the table. Once the equivalence point is reached, plot the data points on the graph paper. The y-axis quantifies acidity and the x-axis indicates the drops of acid added.
Compare the resulting titration curve to the curves given by a strong base titrated with a strong acid and a strong base with a weak acid (since the provided alkali is NaOH, a strong base). From this comparison determine if the unknown acid is strong or weak.




lab_1_chan.JPG
Dilute solution to form .01M NaOH solution


lab_2_chan.JPG
Transfer about 1 mL (if too small an amount add 10 mL) of known base into a new container



lab_3_chan.JPG
Add a drop (if too small an amount add 1 mL) of unknown acid solution. The acid of the sample above is a strong acid, HCl. Record pH with each drop and graph.


strong_base_with_strong_acid.JPG
strong base - strong acid (the graph made should look like this)


lab_4_chan.JPG
The acid of the sample above is a weak acid, HF. Record pH with each drop and graph.


strong_base_with_weak_acid.JPG
strong base - weak acid (the graph made should look like this)





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