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 Reactions in Aqueous Solutions

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عدد المساهمات : 17
تاريخ التسجيل : 26/04/2010
العمر : 28
الموقع : منتدى عبد العزيز هانى العشى

مُساهمةموضوع: Reactions in Aqueous Solutions   الإثنين أبريل 26, 2010 4:39 pm

Chapter 4 Dr. Nabil EL-Halabi
Reactions in Aqueous Solutions
General Properties of Aqueous Solutions:
A solution: is a homogenous mixture of two or more substances.
The solute: is the substance present in the smaller amount.
The solvent: is the substance present in a larger amount.
Aqueous Solutions: in which the solute initially is a liquid or a solid and the solvent is water.

Electrolytes versus Nonelectrolytes:
An electrolyte: is a substance that, when dissolved in water, results in a solution that can conduct electricity.
A nonelectrolyte: is a substance that, when dissolved, results in a solution that does not conduct electricity.
Pure water is a very poor conductor of electricity. However, if we dissolve NaCl (an ionic compound) in water, it breaks up into Na+ and Cl- ions that are responsible for conducting electricity.

Water is a very effective solvent for ionic compound, it has positive and negative poles; for that reason it referred to as a polar solvent.
In solution, each Na+ ion is surrounded by a number of water molecules and each Cl- ion is surrounded by a number of water molecules.
Hydration: is the process in which an ion is surrounded by water molecules arranged in a specific manner.









Conducting electricity in solutions, means we have cations (+) and anions (-).
In this book we use the term dissociation for ionic compound and ionization for acids and bases.
Strong Electrolyte – 100% dissociation (complete ionization). Examples: HCl, HNO3, NaCl.

Weak Electrolyte – not completely dissociated (not completely ionized). CH3COOH, HNO2.

The double ( )arrow in the equation means that the reaction is reversible.
Nonelectrolyte does not conduct electricity. No cations (+) and anions (-) in solution.


Classification of Solutes in Aqueous Solution
Strong Electrolyte Weak Electrolyte Nonelectrolyte
HCl CH3COOH (NH2)2CO urea
HNO3 HF CH3OH methanol
HClO4 HNO2 C2H5OH ethanol
H2SO4 NH3 C6H12O6 glucose
NaOH HCN C12H22O11 Sucrose
Ba(OH)2
Ionic compounds



Precipitation Reactions:

A precipitate: is an insoluble solid that separates from the solution.

Solubility: The maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.

Solubility Characteristics of Ionic Compounds in Water at 25oC
Soluble Compounds Exceptions
Compounds containing alkali metal ions and NH4+
NO3- , HCO3- , ClO4- , ClO3-
Cl- , Br- , I- Halides of Ag+ , Hg22+ , Pb2+
SO42- Sulfates of Ag+, Ca2+, Sr2+, Ba2+, Hg2+, Pb2+
Insoluble Compounds Exceptions
CO32- , PO43- , S2- Compounds containing alkali metal ion and NH4+
OH- Compounds containing alkali metal ion and Ba2+
Example: 4.1

Molecular Equations and Ionic Equations:


Molecular Equation: reactants and products are written as if they were undissociated compounds.
Total Ionic Equation: all soluble ionic substances (strong electrolytes) are dissociated into ions.
Net Ionic Equation: Eliminate spectator ions.
Spectator Ions: ions that are not involved in the overall reaction (These are ions that are unchanged in the chemical reaction.). In this case the Na+ and NO3- are spectator ions.
Another example: when aqueous solution of BaCl2 is added to Na2SO4, a white precipitate of BaSO4 is formed.
The following steps summarize the procedure for writing ionic net ionic equations:
1- Write the balanced molecular equation.
2- Write the ionic equation showing the strong electrolytes as ions.
3- Determine precipitate from solubility rules.
4- Cancel the spectator ions on both sides of the ionic equation.
Example: 4.2

Acid-Base Reactions:
Acids:
1- Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid.
2- React with certain metals to produce hydrogen gas.

3- React with carbonates and bicarbonates to produce carbon dioxide gas.

4- Aqueous acid solutions conduct electricity.

Bases:
1- Have a bitter taste.
2- Feel slippery. Many soaps contain bases.
3- Change the color of litmus from red to blue.
4- Aqueous base solutions conduct electricity.

Brønsted Acids and Bases:
A Brønsted acid is a proton donor.
A Brønsted base is a proton acceptor.
HCl is a Brønsted acid because it donates a proton in water:

A Brønsted acid must contain at least one ionizable proton.
The proton exists in the hydrated form. Therefore, the ionization of HCl should be written as:

H3O+ is called hydrated proton = Hydronium ion.
Keep in mind that both notations H+(aq) & H3O+ represent the same species.
Monoprotic acid; that is, each unit of the acid yields one hydrogen ion upon ionization.

Diprotic acid; that is; each unit of the acid gives up 2 H+ ions, in two separate steps:

Triprotic acid; that is, each unit of the acid gives up 3 H+ ions.


NaOH is strong electrolyte. This means that it is completely ionized in solution:

The OH- ion can accept a proton as follows:

Thus OH- is a Brønsted base.
Thus, when we call NaOH or any other metal hydroxide a base, we actually referring to the OH- species derived from the hydroxide.
NH3 is classified as a Brønsted base because it can accept a H+ ion,
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
NH3 is a weak electrolyte (weak base) because only a small fraction of dissolved NH3 molecules react with water to form NH4+ and OH- ions.
Example: 4.3

Acid-Base Neutralization:
Neutralization Reaction: is a reaction between acid and base.
Acid + base  salt + water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  Na+(aq) + Cl-(aq) + H2O(l)
Therefore, the reaction can be represented by the net ionic equation:
H+(aq) + OH-(aq)  H2O(l)
Both Na+ and Cl- are spectator ions.
If we had started this reaction with equal molar amounts of the acid and the base, at the end of the reaction we would have only a salt and no leftover acid or base. This characteristic of acid-base neutralization reaction.
These are also examples of acid-base neutralization reactions, represented by molecular equation:
HF(aq) + KOH(aq)  KF(aq) + H2O(l)
H2SO4(aq) + 2NaOH(aq)Na2SO4(aq) + 2H2O(l)
HNO3(aq) + NH3(aq) NH4NO3(aq)
NH3(aq) can be expressed as NH4+(aq) and OH-(aq), then the equation becomes:
HNO3(aq) + NH4+(aq) + OH-(aq) NH4NO3(aq) + 2H2O(l)

Oxidation-Reduction reactions:
Oxidation-Reduction (redox) reactions involve a transfer of electrons.
Oxidation: loss of electrons
Reduction: gain of electrons
Reducing agent: substance oxidized
Oxidizing agent: substance reduced2Ca(s) + O2(g) 2CaO(s)
2Ca(s) 2Ca2+ + 4e oxidation half-reaction (lose e)
O2(g) + 4e 2O2- reduction half-reaction (gain e)
_____________________________________________
2Ca(s) + O2(g) + 4e 2Ca2+ + 2O2- + 4e
2Ca(s) + O2(g) 2Ca2+ + 2O2-
Finally, the Ca2+ and O2- combine to form CaO

2Ca(s) + O2(g) 2CaO(s)
Ca is reducing agent and O2 is oxidizing agent.

Oxidation number: the number of charges an atom would have in a molecule (or an ionic compound) if electrons were transferred completely.
An increase in oxidation number indicates oxidation.
A decrease in oxidation number indicates reduction.
H2(g) + Cl2(g)2HCl(g)
Oxidation No.: 0 0 +1 –1
Rules to Assign Oxidation Numbers:
1- Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2- In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3- The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1.
4- The oxidation number of hydrogen is +1 except when it is bonded to metals in binary
compounds. In these cases, its oxidation number is –1.
5- Group IA metals are +1, IIA metals are +2 and fluorine is always –1.
6- The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on
the molecule or ion.
Example: 4.4

Displacement reaction: the reaction in which one ion or atom in the reaction is replaced (displaced) by another ion (or atom) of another element.
In the metal-HCl reaction, the H+ ions are displaced by a metal ion.
Displacement reactions are among the most common redox reactions.
Zn(S) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
The net ionic equation is Zn(S) + Cu2+(aq)  Zn2+(aq) + Cu(s)
Similarly, metallic copper displaces silver ions from a solution containing AgNO3
Cu(S) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
The net ionic equation is Cu(S) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)
Activity series: is a convenient summary of the results of many possible displacement reaction.
According to this series, any metal above hydrogen will displace it from water or from an acid, but metals below hydrogen will not react with either water or an acid.
Any species listed in the series will react with a compound containing any species listed below it.
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Ca(s) + 2H2O(l)  Ca(OH)2(s) + H2(g)

For displacement reactions involving halogens the reactivity is
F2 > Cl2 > Br2 > I2
Cl2(g) + 2KBr(aq)  2KCl(aq) + Br2(g)
The net ionic equation is Cl2(g) + 2Br-(aq)  2Cl-(aq) + Br2(g)

Concentration of Solutions: is the amount of solute present in a given quantity of solvent or solution.
Molarity (M), or Molar Concentration: the number of moles of solute in 1 liter (L) of solution
M = molarity = moles of solute / liters of solution
A 1.46 molar solution contains 1.46 mole of the solute in 1 L of the solution; or 0.73 moles of the solute in 0.500 L of solution.
Example 4.5

Dilution of Solutions:
Dilution: is the procedure for preparing a less concentrated solution from a more concentrated one.
Moles of solute before dilution = moles of solute after dilution
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Reactions in Aqueous Solutions
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