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Chapter 4

Types of Chemical Reactions and

Solution Stiochiometry

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Preview

the contents of this chapter will introduce you to the following topics:

Water, Nature of aqueous solutions, types of electrolytes, and dilution

Types of chemical reactions: precipitation, acid base reactions and oxidation-reduction reaction

Stoichiometry of reactions and balancing the chemical equations

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4.1 Water, the Common Solvent

Water is one of the most important substances on earth:– Cooling – engine, nuclear power plants, and many.– Transportation …etc.

Water dissolve many different substances,

e.g. salts, sugar, and many other

To understand this process, we need to consider the nature of water:

– As molecule, is H2O

– Shape is V-shape with angle 105o

– Band type of each O – H is covalent band and polar [electrons are not equirdlntty shared]

– Polarity is polar with +ve charges of hydrogen and -ve on oxygen.

This Polarity of water gives it the greatest ability to dissolve compounds.

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4.1 Water, the Common Solvent

• This Polarity of water gives it the greatest ability to dissolve compounds. Figure 4.2 shows schematic ionic solid dissolving in water. This process is called "Hydration".

Note: • The stronger the ion-water attraction, the higher the solubility.• Therefore, not all the solid have the same solubility [chapter 11].

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4.1 Water, the Common Solvent

Water also dissolve non-ionic substances e.g. alcohols – "polar" and compatible structure to water: "like – dissolve – like"

Figure 4.3

many substances don't dissolve in water e.g. animal fats "non-polar"

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Solution – composite of solute (substance to be dissolved) + Solvent (Major substance e.g. water).

Solute:

Solvent:

4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes & non–Electrolyte

dissolves in water (or other “solvent”) changes phase (if different from the solvent) is present in lesser amount (if the same phase as the solvent)

retains its phase (if different from the solute) is present in greater amount (if the same phase as the solute)

Solution Solvent Solute

Soft drink (l)

Air (g)

Soft Solder (s)

H2O

N2

Pb

Sugar, CO2

O2, Ar, CH4

Sn

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One major property for characterizing aqueous solutions is its "Electrical conductivity" or its ability to conduct an electric current. Three solutions are observed: non electrolytes, weak, electrolytes and strong electrolyte. figure 4.4

The basis for conductivity properties of solutions was first correctly identified by Svante Arrhenius (1859 – 1927)

4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes & non–Electrolyte

nonelectrolyte weak electrolyte strong electrolyte

An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.

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Arrhenius postulated that the electric current depend directly on the number of ions present.

4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes & non–Electrolyte

Strong Electrolyte – 100% dissociation

NaCl (s) Na+ (aq) + Cl- (aq)H2O

Weak Electrolyte – not completely dissociated

CH3COOH CH3COO- (aq) + H+ (aq)Nonelectrolyte does not conduct electricity?

No cations (+) and anions (-) in solution

C6H12O6 (s) C6H12O6 (aq)H2O

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4.3 The Composition of Solutions:

To perform stoichiomctric calculations at any chemical reactions you must know two things:– The nature of the reaction – exact forms of chemical in solutions.

– The amounts of chemicals – "concentration"

concentration of a solution can be described in many different ways, % , molar, molal, mole. faction .. etc.

We will consider here the unit "molar"(M) or molarity: "the method used to prepare molar solutions".

M

M

molaritymoles of soluteliters of solution

HClmoles of HCl

liters of solution3

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Molarity (M) = moles of solute per volume of solution in liters:

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4.3 The Composition of Solutions:Notes:

For ionic systems the solution prepared will contain the number of moles prepared but for ionic species its different story

e.g. 1.0M NaCl contains 1.0 mole NaCl or more accurate 1.0 mole of Na+ and 1.0 mole of Cl-.

Molarity can be used to determine number of moles per certain volumes where:

Moles = Liters of solution x Molarity.

Example 4.3Gives the concentration of each type of ion in the solutions of0.50M Co(NO3)2

Example 4.4 Calculate the number of moles of Cl- ions in 1.75 L of 1.0x10-3M ZnCl2

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4.3 The Composition of Solutions:

Dilution:it is the procedure to get low concentrated solution (diluted) from a high concentrated one.

Moles of solute after dilution = moles of solute before dilution.(M x V) after = (M x V) before.

Example 4.7 What volume of 16 M sulfuric Acid must be used to prepare 1.5 L of a 0.10M H2sO4 solution?

Standard Solution:Solution used in chemical analysis. it has accurately known concentration

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4.4 Types of Solution Reactions

Millions of possible chemical reaction needs system for grouping them into classes. The commonly used by chemists:

•Precipitation reactionsAgNO3(aq) + NaCl(aq) AgCl(s) +

NaNO3(aq)

•Acid-base reactionsNaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

•Oxidation-reduction reactionsFe2O3(s) + Al(s) Fe(s) + Al2O3(s)

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4.5 Precipitation Reactions

Precipitate – insoluble solid that separates from solution

molecular equation

ionic equation

net ionic equation

Na+ and NO3- are spectator ions

PbI2

Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2NaNO3 (aq)

precipitate

Pb2+ + 2I- PbI2 (s)

Pb2+ + 2NO3- + 2Na+ + 2I- PbI2 (s) + 2Na+ + 2NO3

-

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Simple Rules for Solubility of Aq. Solutions

1. Most nitrate (NO3) salts are soluble.

2. Most alkali (group 1A) salts and NH4+ are soluble.

3. Most Cl, Br, and I salts are soluble (NOT Ag+, Pb2+, Hg22+)

4. Most sulfate salts are soluble (Except BaSO4, PbSO4, HgSO4, CaSO4)

5. Most OH salts are only slightly soluble (Except NaOH, KOH are soluble)

6. Most S2, CO32, CrO4

2, PO43 salts are only slightly soluble, i.e., Not

soluble.

4.5 Precipitation Reactions

Exercise 4.8:Using the solubility rules in table 4.1, predict what will happen when the following pairs of solutions are mixed.

a. KNO3(aq) & BaCl2(aq)b. Na2SO4(aq) & Pb(NO3)2(aq)c. KOH(aq) & Fe(NO3)(aq)

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Writing Net Ionic Equations

AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)

Ag+ + NO3- + Na+ + Cl- AgCl (s) + Na+ + NO3

-

Ag+ + Cl- AgCl (s)

Write the net ionic equation for the reaction of silver nitrate with sodium chloride.

1. Write the balanced molecular equation.

2. Write the ionic equation showing the strong electrolytes

3. Determine precipitate from solubility rules

4. Cancel the spectator ions on both sides of the ionic equation

4.6 Describing Reactions in Solution

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4.7 Stoichiometry of Precipitation Reactions

The procedures for calculating quantities of reactants and products involved in chemical reaction. The following steps summarized the procedure:Step 1: Identify the present in the combined solution, and

determine what reaction occurs.Step 2: write the balanced net ionic equation.Step 3: calculate the moles of reactants.Step 4: determine which reactant is limiting.Step 5: calculate the moles of product or product as required.Step 6: convert to grams or other units, as requires.

Example 4.11When aqueous solutions of Na2SO4 and Pb(NO3)2 are mixed, PbSO4 precipitates. Calculate the mass of PbSO4 formed when 1.25L of 0.0500M Pb(NO3)2 and 2.00L of 0.0250M Na2SO4 are mixed.

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4.8 Acid – Base Reactions:

Acids

Have a sour taste. Taste of vinegar is due to acetic acid.

Citrus fruits contain citric acid.

React with certain metals to produce hydrogen gas.

React with carbonates and bicarbonates to produce carbon dioxide gas

Have a bitter taste.

Feel slippery. Many soaps contain bases.

Bases

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4.8 Acid – Base Reactions:

Arrhenius acid is a substance that produces H+ (H3O+) in water

Arrhenius base is a substance that produces OH- in water

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4.8 Acid – Base Reactions:

A Brønsted acid is a proton donorA Brønsted base is a proton acceptor

acidbase acid base

A Brønsted acid must contain at least one ionizable proton!

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4.8 Acid – Base Reactions:

acid + base salt + water

HCl (aq) + NaOH (aq) NaCl (aq) + H2O

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O

H+ + OH- H2O

Describing Reactions in Solution

They are also called neutralization reaction

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4.8 Acid – Base Reactions:

x acid + y base salt + water

They are also called neutralization reaction

Main reaction is titration, the key terms are:Titrant - solution of known concentration used in titrationAnalyte - substance being analyzed Equivalence point - enough titrant added to react exactly with the analyte Endpoint - the indicator changes color so you can tell the equivalence point has been reached.

a. Write the correct balanced acid–base reaction.

b. Use the following equation:

y. (M.V)acid = x. (M.V)base

The neutralization reaction calculation:

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4.8 Acid – Base Reactions:In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.

Equivalence point – the point at which the reaction is complete

Indicator – substance that changes color at (or near) the equivalence point

Slowly add baseto unknown acid

UNTIL

the indicatorchanges color

4.7

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4.9 Oxidation-Reduction Reactions:

(electron transfer reactions)

2Mg (s) + O2 (g) 2MgO (s)

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

Oxidation half-reaction (lose e-)

Reduction half-reaction (gain e-)

2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e-

2Mg + O2 2MgO

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4.9 Oxidation-Reduction Reactions:

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4.9 Oxidation-Reduction Reactions:Rules for Assigning Oxidation States

1. Oxidation state of an atom in an element = 02. Oxidation state of monatomic element ions = charge3. Oxygen =-2 in covalent compounds

(except in peroxides where it = -1)4. H = +1 in covalent compounds5. Fluorine = -1 in compounds6. Sum of oxidation states = 0 in compounds Sum of oxidation states

= charge of the ion

IF7

F = -17x(-1) + ? = 0

I = +7

Oxidation numbers of the elements in the following ?

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4.10 Balancing Oxidation – Reduction Equations

Balancing by Half-Reaction Method (Acidic)1. Write separate reduction, oxidation reactions

2. For each half-reaction:• Balance elements (except H, O)

• Balance O using H2O

• Balance H using H+

• Balance charge using electrons

3. If necessary, multiply by integer to equalize electron count

4. Add half-reactions

5. Check that elements and charges are balanced

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4.10 Balancing Oxidation – Reduction Equations

Half-Reaction Method - Balancing in Base

1. Balance as in acid.

2. Add OH- that equals H+ ions (both sides!)

3. Form water by combining H+, OH-

4. Check elements and charges for balance

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4.10 Balancing Oxidation – Reduction Equations

MnO4- + ClO2

- → MnO2 + ClO4-