Course Outline

General Chemistry--Unit 3

04/25/09

�Before You Begin:

To master this material you need to understand chemical formulas and know the common charges for the representative elements from periodic law.

Reaction Equation Review

A chemical reaction equation is a symbolic notation system for describing chemical changes. Reactants are the substances that are present before the chemical change. These are written on the left side of the reaction equation. The products are the substances that are present after the chemical change. These are written on the right side of the reaction equation. An arrow (à) separates the reactants and the products. An equilibrium is a chemical reaction in which the reactants are converted into products at the same time and rate that the products are converted into reactants. The reaction arrow for an equilibrium looks like right and left pointing arrows written one on top of the other ().

Information about the state (solid, liquid, gas) or solution (aqueous, precipitate) is written in parenthesis after the chemical formula.

Coefficients are the relative number of moles of each of the reactants and products. These numbers are written in front of the chemical formulas. In the reaction equation below, the coefficients indicate that there are two moles of sodium hydroxide and water for every one mole of sulfuric acid and sodium sulfate.

2NaOH_(aq) + H₂SO_4(aq) à Na₂SO_4(aq) + 2H₂O_(l)

Students are expected to be able to balance chemical reaction equations (i.e. to fill in missing coefficients).

4Concept Check: What does this reaction equation mean?

Ba(OH)_2(aq) + H₂SO_4(aq) à BaSO_4(ppt) + 2H₂O_(l)

Answer: the reaction equation gives us information about all of the reactants and products via the chemical formulae:

The reactants are Ba(OH)₂ and H₂SO₄. Ba(OH)₂ is barium hydroxide, an ionic compound consisting of one barium +2 ion for every two hydroxide -1 ions. The hydroxide ion has one hydrogen and one oxygen atom bonded together and an extra electron. H₂SO₄ is sulphuric acid, which consists of two hydrogen +1 ions for every one sulphate -2 ion. The sulphate ion has one sulphur and four oxygen atoms bonded together with two extra electrons. The products are BaSO₄ and H₂O. BaSO₄ is barium sulphate an ionic compound consisting of one barium +2 ion for every sulphate -2 ion. Water is a molecular compound consisting of molecules of two hydrogen atoms bonded to an oxygen atom. If you do not remember how to name chemical compounds and are unclear about subscripts, see "Prep Chem" Compounds.

The reaction equation gives us information about the states of the substances:

Barium hydroxide and sulphuric acid are aqueous solutions, meaning that they have been dissolved in water. Barium sulphate is a precipitate, meaning it is an insoluble solid formed by combining aqueous solutions. Water is a liquid meaning it has definite volume but takes the shape of its container. If you do not remember reaction equation symbols, review the stoichiometry page.

The reaction equation gives us information about the relative amounts of the reactants and products:

For every one mole of barium hydroxide, sulphuric acid and barium sulphate, two moles of water are produced. If you cannot fill in missing coefficients, review the balancing chemical reactions tutorial.

Stoichiometry is the process of using information from a balanced chemical reaction equation to deduce relative amounts of reactants and products. One can use the formula weight of the reactants to find the number of moles of reactants. The coefficients are mole ratios, so one can use the coefficients to determine the number of moles of products formed for a given amount of reactants. From moles of products one can determine mass (using formula weight) or volume (using the ideal gas law) or concentration (using solution concentration definitions discussed below). We will use stoichiometry to explore the descriptive chemistry reactions. See the “Prep Chem” section for an introduction to stoichiometry and the tutorial section for more practice.

Types of Reactions

There are many different methods of categorizing chemical reactions. One is to classify reactions based on changes in the number of reactants and products. By this method, most of the chemical reactions studied in beginning chemistry fall into one of four categories: combination, decomposition, single displacement or metathesis.

Combination:

A combination reaction, sometimes called addition reaction, is one in which two reactants form a single product. Some more complicated combination reactions have more reactants and/or products, but the general trend is for the number of different chemical substances to decrease as the reaction progresses. An example of a combination reaction is HCl_(g) + NH_3(g) à NH₄Cl_(s).

Decomposition:

A decomposition reaction is one in which a single reactant breaks apart into more than one product. Decomposition reactions are the reverse of combination reactions. An example of a decomposition reaction is the decomposition of hydrogen peroxide: 2H₂O_2(aq) à H₂O_(l) + O_2(g).

Single Displacement:

During a single replacement reaction, a reactant element takes the place of a chemically similar element that is part of a compound. An element and a compound react to form a different element and compound. An example is Zn_(s) + Cu(NO₃)_2(aq) à Cu_(s) + Zn(NO₃)_2(aq). Note that the zinc metal takes the place of the copper ion in the compound. Zinc is more similar to copper than to nitrogen or oxygen.

Double Displacement:

During a double replacement reaction, two compounds exchange elements to form two different compounds. One type of double replacement reaction is metathesis. In a metathesis reaction, a pair of ionic compounds exchanges ions in an aqueous solution. An example of a metathesis reaction is NaCl_(aq) + AgNO_3(aq) à AgCl_(ppt) + NaNO_3(aq).

4Concept Check: What types of reactions are these?

CH₃Br_(g) + HCl_(g) à CH₃Cl_(g) + HBr_(g)

2Mg_(s) + O_2(g) à 2MgO_(s)

2NaHCO_3(s) à Na₂CO_3(s) + H₂O_(g) + CO2_(g)

2Al_(s) + 3Pb(NO₃)_2(aq) à 2Al(NO₃)_3(aq) + 3Pb_(s

Answer:

CH₃Br_(g) + HCl_(g) à CH₃Cl_(g) + HBr_(g) is a double displacement (not metathesis because it doesn’t take place in an aqueous solution) reaction.

2Mg_(s) + O_2(g) à 2MgO_(s) is a combination reaction.

2NaHCO_3(s) à Na₂CO_3(s) + H₂O_(g) + CO_2(g) is a decomposition reaction.

2Al_(s) + 3Pb(NO₃)_2(aq) à 2Al(NO₃)_3(aq) + 3Pb_(s) is a single displacement reaction.

Oxidation-Reduction

In some reactions, electrons are transferred from one reactant to another. Oxidation is the loss of electrons. Reduction is the gain of electrons. Any of the types of reactions can also be oxidation-reduction. For example, two of the reactions from the previous concept check are oxidation-reduction reactions:

2Mg_(s) + O_2(g) à 2MgO_(s) and

2Al_(s) + 3Pb(NO₃)_2(aq) à 2Al(NO₃)_3(aq) + 3Pb_(s)  

We can tell these are oxidation reduction reactions because the charges on some of the elements change. In the first example, magnesium element (charge zero) is converted into magnesium +2 ion. In the second example, aluminum element becomes aluminum +3 ion and lead +2 ion becomes lead element (charge zero).

When oxidation-reduction reactions involve molecular substances, it is more difficult to see evidence of the electron transfer, since these substances don’t have charges. Oxidation numbers are one way of keeping track of electron changes. Note that oxidation numbers look like charges, but they are an artificial means of keeping track of electrons rather than something ‘real’ like a charge. To keep oxidation numbers from being confused with charges, they need to be labeled in some way. We will write them above the atomic symbol in red font:

1. The oxidation number of an atom in an element in its most stable form is 0. For example, the atoms in the oxygen molecule have an oxidation number of zero.
2. The oxidation number of any monatomic ion is equation to its charge. For example, the aluminum ion has an oxidation number of +3.
3. The sum of the oxidation numbers of the atoms in a molecule or polyatomic ion must equal the charge (zero for a molecule). For example, the atoms in the nitrate ion must have oxidation numbers that total -1.
4. Atoms in a molecule or polyatomic ion often have oxidation numbers equal to their most common charge. This is your best, first guess, especially for elements on the edges of the periodic table. Assign oxidation numbers for these atoms first, then assign the rest so that they obey rule 3. For example, if the oxygen atoms in the nitrate ion have an oxidation number of  -2, the nitrogen must have an oxidation number of +5 because 5 + 3(-2) = -1.

5. Oxygen usually has an oxidation number of -2. Exceptions are elemental oxygen which has an oxidation number of zero; peroxide ion, O₂^-2, which has an oxidation number of -1; and superoxide ion, O₂^-1, which has an oxidation number of -1/2. Yes, that’s right, -1/2! Because an oxidation number is just a bookkeeping device rather than a real charge, it is possible to have fractional oxidation numbers. If the ion has a negative one charge, and the sum of the oxidation numbers has to equal the charge, and the ion has two atoms, then each atom has a -1/2 oxidation number. Superoxide is very reactive and unstable. It is formed in the mitochondria during electron transport. We have an enzyme, superoxide dismutase, designed to convert this dangerous ion to less dangerous hydrogen peroxide. Amyotrophic lateral sclerosis (Lou Gehrig’s Disease) is believed to be due to a faulty gene that codes this enzyme. In this condition, superoxide does progressive damage to neurons because the enzyme doesn’t work properly.
6. Hydrogen usually has an oxidation number of +1. It will have an oxidation number of zero as an element, and it will have an oxidation number of -1 in a hydride. You can recognize the hydrides because they will be ionic compounds with an active metal. An example is lithium aluminum hydride, LiAlH₄. To assign oxidation numbers for lithium aluminum hydride, give the lithium a +1 oxidation number first. That is most common charge for the alkali metals, and this element is the least likely of the three to do anything unusual. Next assign a -1 oxidation number to the hydrogen. This is an ionic compound, and neither the lithium nor the aluminum will take the role of the negative ion, so hydrogen has to do it. That leaves a +3 for the aluminum because 1 + 3 + 4(-1) = 0, and the total charge has to equal zero for a compound.

The purpose of assigning oxidation numbers is to determine if an oxidation-reduction reaction has occurred. Assign numbers to all of the atoms in a reaction equation, then scan to see if the numbers for the reactant elements match the product elements. If the oxidation number decreases, electrons were gained and that element was reduced. If the oxidation number increases, electrons were lost and that element was oxidized. In the following example, nickel is reduced (the oxidation number drops from +4 to +2) and cadmium is oxidized (the oxidation number increases from zero to +2). This is the chemical reaction for the discharge cycle in a nicad battery.

One specific type of oxidation-reduction reaction is combustion, the rapid chemical combination of a substance with oxygen. We can predict the outcome of a combustion reaction for many molecular compounds. Each element in the compound will be converted into its most stable non-metal oxide. For carbon, this is carbon dioxide; for hydrogen, this is water. We predict that, when benzene, C₆H₆, undergoes combustion, it will be converted into carbon dioxide and water:

2C₆H_6(l) + 15O_2(g) à 12CO_2(g) + 6H₂O_(l)

This is an oversimplification. The products of many combustion reactions are much more complicated mixtures resulting from incomplete combustion and inadequate mixing. Carbon monoxide and smoke are almost always produced when carbon containing compounds burn.

Another type of oxidation-reduction reaction is corrosion, the unwanted spontaneous conversion of elemental metals into ionic compounds, usually oxides, upon exposure to moist air. The example 2Mg_(s) + O_2(g) à 2MgO_(s) is a combustion reaction, if it happens in the chemistry lab when magnesium is heated in air until it gives off a bright light (burns). This same reaction is corrosion if it occurs slowly and gradually, when an aging piece of magnesium metal develops a grayish powdery film while sitting on the chemistry stockroom shelf. Both of these chemical changes are oxidation-reduction and single displacement. Combustion, however, is much faster than corrosion.

Electrochemistry is the study of oxidation-reduction reactions that produce or are driven by electrical energy. We will study electrochemistry along with corrosion in more detail when we explore the descriptive chemistry of the metals.

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