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Chemistry 131

Lecture Notes

Chapter 4 Chemical Reactions

1. Ions in Aqueous Solution

2. Types of Chemical Reactions

3. Working with Solutions

4. Quantitative Calculations

 

 

Objectives. Be able to:

Ions in Aqueous Solution

Definitions

  1. solution- a homogeneous mixture of two or more substances.
  2. solvent- the substance present in a solution in the greatest amount.
  3. solute-the substance(s) present in the solution in amounts less than the solvent.
  4. soluble- substances that dissolve to a significant extent in the solvent are said to be soluble
  5. insoluble-substances that do not dissolve to a significant extent in the solvent are said to be insoluble.
  6. solubility- the measure (grams or moles) of substance that can be dissolved in a specific amount of solvent (g/100mL).

Solubility

Ionic theory of solutions and solubility rules


 

 

Strong electrolytes are soluble salts, strong acids and strong bases.

Weak electrolytes are weak acids, weak bases, and weakly ionized salts.

Most other substances(usually molecular) are non-electrolytes.

 

 

See solubility rules table 4.1

1. NH4+ and Group IA salts are soluble

2. CH3COO-1 and NO3-1 salts are all soluble

3. Most Cl-, Br-, and I- salts are soluble Exceptions: Pb+2, Hg+2, Ag+

4. Most sulfates are soluble. Exceptions: Pb+2, Hg+2, Ag+, Ca+2, Sr+2, Ba+2

Rule of thumb - most salts containing +1 or -1 ions are soluble

 

Acids and Bases

Acids and bases are some of the most important electrolytes. (see Table 4.2). They can cause color changes in certain dyes called acid-base indicators.

Household acids and bases. (see Figure 4.7)

Red cabbage juice as an acid-base indicator. (see Figure 4.8)

Definitions:

Acids (strong or weak)

Bases (strong or weak)

 

Know strong acids and strong bases

Strong acids-

Name them!!

Strong bases-

GIA and GIIA below Mg

HCl LiOH
HBr NaOH
HI KOH
HNO3 RbOH
H2SO4 CsOH
HClO3 Ca(OH)2
HClO4 Sr(OH)2
  Ba(OH)2

 

Molecular and ionic equations

Strong Electrolytes, (soluble salts, strong acids and strong bases) are written as ions.

All other compounds or elements (weak and non-electrolytes) are written as a complete formula with no charge.

Example:

molecular

3Na2CO3(aq)+ 2AlCl3(aq)--->6NaCl(aq)+ Al2(CO3)3(s)

total ionic

 

net ionic

 

 

 

2. Types of Chemical Reactions

Precipitation reactions

Metathetical Reactions (Double Replacement)

The compounds, usually ionic, exchange partners-two reactants go to products

Remember that indication of a reaction are:

  • ***know your solubility rules

    NOTE: A gas such as CO2, NH3, NO2, H2S, SO2, or any gaseous element may form in double displacement reactions also.

    NOTE: Water (acid-base rxn) or a weakly ionized substance is formed, heat is often generated in double displacement. A weaker acid or water is formed in a neutralization reaction of an acid and a base***know your strong acids and bases

  • Examples of precipitation rxns, ppt forms

    1. Aluminum nitrate reacts with sodium phosphate.

     

     

     

     

     

    2. silver nitrate reacts with iron (II) chloride

     

     

     

     

     

     

    Acid-base reactions

    Acid-base reactions may form water or a gas. They are often double displacement (or metathetical) reactions.

    Examples in aqueous solution (write net ionic equations as well).

    Note:

    H2CO3 (aq) H2O(l) + CO2 (g) recognize that carbonic acid is in equilibrium and goes to some water and carbon dioxide.

    NH4OH (aq) H2O (l)+ NH3(g) recognize that ammonium hydroxide forms water and ammonia and is often written as NH3(aq)

    1. NaHCO3(aq) + HNO3(aq)

     

     

     

     

     

    2. NaOH(aq) + HCl(aq)

     

     

     

    Oxidation reduction reactions

    Oxidation-reduction reactions involve the transfer of electrons from one species to another.

    Oxidation is defined as the loss of electrons.

    Reduction is defined as the gain of electrons.

    Oxidation and reduction always occur simultaneously.

    The reaction of an iron nail with a solution of copper(II) sulfate, CuSO4, is an oxidation- reduction reaction. (see Figure 4.11)

    The molecular equation for this reaction is:

     

     

     

     

    Review assigning oxidation numbers. See table 4.5

    Oxidation is?

    Reduction is?

     

     

    What is being oxidized and reduced in the following equation. What is the oxidizing agent and reducing agent?

     

    Ca (s) + Ag+ (aq) Ca+2(aq) + Ag(s)

     

     

     

     

    Table 4.6 The activity series.

    How to recognize simple redox reactions: An element and a compound form an element and a compound based upon the activity of each metal or H2.

    Examples:

    Zn(s) + 2 HCl(aq) ==> H2 + ZnCl2

    Cu(s) + AgNO3 (aq)==>

    Cu(s) + HCl (aq)===>

    Summary of reactions

     

    Patterns of Chemical Reactivity

    Using the Periodic Table- members of a family will tend to react in a chemically similar way.

    Alkali reactivity

    Na + Cl2.

    Combustion in Air- Combustion rxns are rapid and produce a flame.

    Simple hydrocarbons and compounds of only C, H, and O, "burn" to form carbon dioxide and water.

    Example: Complete and balance the following combustion reaction:

    __ (CH3)2CO + __O2 ==>

     

    Combination and Decomposition Reactions

    Combination Reactions- two or more substances combine to form one product.

    Formation reactions- elements combining to form one mole of product. Redox if elements forming a compound.

    __ Mg (s)+ __ O2 (g)===>

     

    __ C (s)+ Cl2(g) ===>

     

    Decomposition Reactions- one substance reacts to produce two or more products. May be redox or non-redox. With H and O present, water is a product. With CO3-2 present, CO2 is a product.

    __ H2O2(aq)+ light or heat ===>

     

    __ CaCO3 (s) + heat===>

     

     

    Read the section on balancing redox reactions. I will not test you on this method but you need to remember two rules in balancing all chemical equations.

    1. Mass is conserved. The mass of the reactants must equal the mass of the products.
    2. Charge is conserved. The total number of electrons lost must be the same as that gained. Or total charge is the same on each side of the equation.

     

    3. Working with Solutions

    Molar concentrations

    Molar concentration, or molarity (M), is defined as the moles of solute dissolved in one liter (cubic decimeter) of solution.

    M= moles solute/liter of solution.

    Molarity can be used as a conversion factor. A 0.250 M solution of NaOH means

    0.250 mol NaOH = 1 L solution

    Example:

    1. A sample of 0.0341 mol iron(III) chloride, FeCl3, was dissolved in water to give 25.0 mL of solution. What is the molarity of the solution?


    2. How many liters of 0.450 M FeCl3 solution is needed to obtain 0.0368 moles FeCl3?

     

     

    Diluting solutions

    Find the number of mL of 0.100 M HCl that can be prepared from 100.0 mL of 3.00M HCl?

     

    4. Quantitative Calculations

    Gravimetric analysis is a type of quantitative analysis in which the amount of a species in a material is determined by converting the species into a product that can be isolated and weighed.

    Example:

    A solution contains an unknown concentration of citric acid (H3C6H5O7), a triprotic acid. It takes 41.28 mL of 0.2500 M standard sodium hydroxide solution to titrate (react completely) with 10.00 mL of citric acid solution. What is the molarity of the citric acid solution?

     

     

     

     

    MA–> moles A –>moles B –>MB

    Summary

     

     

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