Chapter 6 notes Chemical Reactions and Equations

6.1 CHEMICAL EQUATIONS

 

REPRESENTING CHEMICAL CHANGES

Scientists relay on a variety of shorthand methods for expressing chemical information. You have already seen how chemical symbols are used for the names of elements and chemical formulas for the names of compounds. A chemical equation is a shorthand expression that represents a chemical reaction. A chemical reaction is the process by which one or more substances are changed into one or more new substances. A chemical equation shows the relative amount of each substance taking place in a chemical reaction.

Observable macroscopic changes that indicate that a chemical reaction has occurred:

        Change in color or odor, gas release,

        Production or absorption of heat or light,

        Formation of a precipitate

All chemical reactions follow the LAW OF CONSERVATION OF MASS
The starting substances in a chemical reaction are called reactants. The substances that are formed are called products. The general format for a chemical equation is as follows.

CO2(g)   + H2O(l) à     H2CO3(aq)

    reactants       yield     products

Reactants are generally written on the left side of chemical equations; products are written on the right side.

The letters in parentheses indicate the physical state of each substance involved in the reaction. The following symbols should be used in your work. (g) gas    (s) solid      (l) liquid          (aq) water solution

COEFFICIENTS

The formula of a compound represents a definite amount of that compound. This amount of an ionic compound is called a formula unit. It may be one molecule (covalent) or the smallest number of particles giving the true proportions of the elements in the compound.

When we wish to represent two molecules of water we write 2H2O. The number prefixed as a multiplier is called the coefficient. For example, when the coefficient 3 is written before the formula unit Fe2O3, it means three times everything in the formula. In 3Fe2O3 there are 6 iron atoms and 9 oxygen atoms.

BALANCING EQUATIONS

The first step in writing a chemical equation is writing a word equation. It is composed of the names of the substances that are involved in a chemical reaction.

In a reaction involving copper(I) chloride dissolved in water reacts to hydrogen sulfide gas yields hydrochloric acid and the precipitate copper(I) sulfide.

copper(I) chloride + hydrogen sulfide à copper(I) sulfide + hydrochloric acid

 

The second step is writing a skeleton equation. This equation includes the chemical symbols and formulas for all the reactants and products identified in the word equation.

CuCl(aq) + H2S(g) à Cu2S(s) + HCl(aq)

 

The third step in writing a chemical equation is balancing the equation. The balanced equation includes the coefficients, numbers placed directly in front of the chemical formulas and symbols. The coefficients indicate the relative proportions of each substance involved in the chemical reaction.

2CuCl(aq) + H2S(g) à Cu2S(s) + 2HCl(aq)

 

This equation states that two units of CuCl(aq) react with one unit of H2S(g) producing one unit of Cu2S(s) and two units HCl(aq).

The fourth step is to check and make sure that it is balanced.

Reactants       Products

Cu            2                      2

Cl             2                      2

H              2                      2

S              1                      1

 

Chemical equations must be balanced according to the law of conservation of mass, which states that matter cannot be created or destroyed in a chemical reaction.

EXAMPLE

Sodium reacts with water to produce a metallic hydroxide and hydrogen gas.

Write a balanced equation for the reaction.

Solving Process:

Step 1. Write the word equation. Determine the products and reactants.

sodium + water à sodium hydroxide + hydrogen

Step 2. Write a skeleton equation. Since hydrogen is a diatomic gas, its formula is H2. The formula for water may be written as HOH; this may make it easier to balance the equation.

Na + HOHà NaOH + H2

Step 3. Balance the equation. The metallic element sodium is balanced. One atom of sodium is on each side of the equation. There is one hydrogen atom on the reactant side (the H in OH has been accounted for) and 2 hydrogen atoms on the product side. Place a 2 in front of the HOH to balance the hydrogen atoms.

Na + 2HOH à NaOH + H2

There are now 2OH on the left and 1 on the right. Place a 2 in front of the NaOH to give the same number of OH on each side.

Na + 2HOH à 2NaOH + H2

Put a 2 in front of the sodium metal. The balanced equation reads

2Na(s) + 2HOH(l) à 2NaOH(aq) + H2(g)

Step 4. Check to see if the equation is balanced.

Reactants       Products

Na            2                      2

H              4                      4

O             2                      2

6.2 TYPES OF REACTIONS

CLASSIFYING CHEMICAL CHANGES

The products of a chemical reaction may often be predicted by applying known facts about common reaction types. While there are hundreds of different “kinds” of chemical reactions, only five general types of reactions will be considered: synthesis, decomposition, single displacement, double displacement, and combustion.

 

a. Synthesis. In a synthesis reaction two or more substances are combined to form one new and more complex substance.

 The general form is as follows.

element/compound + element/compound à compound

a + bà ab

The following are some general types of synthesis reactions.

1. Two or more elements combine to form a compound.

Fe(s) + S(l) à FeS(s)

2. An acid anhydride, nonmetallic oxide, combines with water to give an acid.

SO2(g) + H2O(l) à H2SO3(aq)

3. A basic anhydride, metallic oxide, combines with water to form a base.

Na2O(s) + H2O(l) à 2NaOH(aq)

4. A basic oxide combines with a nonmetallic oxide to form a salt,

CO2(g) + Na2O(s) à Na2CO3(s)

 

b. Decomposition. When energy in the form of heat, electricity, light, or mechanical shock is supplied, a compound may decompose to form simpler substances.

The general form for this type of reaction is as follows.

compound à two or more elements/compounds

ab à a + b

The following are some general types of decomposition reactions.

1. When some acids are heated, they decompose to form water and an acidic oxide.

H2CO3(aq) à CO2(g) + H2O(l)

2. When some metallic hydroxides are heated, they decompose to form a metallic oxide and water.

Ca(OH)2(s)à CaO(s) + H2O(g)

3. When some metallic carbonates are heated, they decompose to form a metallic oxide and carbon dioxide.

Li2CO3(s) à Li2O(s) + CO2(g)

4. When metallic chlorates are heated, they decompose to form metallic chlorides and oxygen.

2KClO3(s) à 2KCl(s) + 3O2(g)

5. Most metallic oxides are stable, but a few decompose when heated.

2HgO(s) à 2Hg(l) + O2(g)

6. Some compounds cannot be decomposed by heat, but can be decomposed into their elements by electricity.

2NaCl(l) à 2Na(s) + Cl2(g)

 

C.Single Displacement. One element displaces another element in a compound. A single displacement has this general form.

element a + compound bc à element b + compound ac

a + bc à b + ac

element d + compound bc à element c + compound bd

d + bc à c + bd

The following are some general types of single displacement reactions.

1. An active metal will displace the metallic ion in a compound of a less active metal.

Fe(s) + Cu(NO3)2(aq)à Fe(NO3)2(aq) + Cu(s)

2. Some active metals, such as sodium and calcium, will react with water to give a metallic hydroxide and hydrogen gas.

Ca(s) + 2H2O(l) à Ca(OH)2(aq) + H2(g)

3. Active metals, such as zinc, iron, and aluminum, will displace the hydrogen in acids to give a salt and hydrogen gas.

Zn(s) + 2HCl(aq) à ZnCl2(aq) + H2(g)

4. An active nonmetal will displace a less active nonmetal.

Cl2(g) + 2NaBr(aq) à 2NaCl(aq) + Br2(aq)

 

d. Double Displacement. The positive portions of two ionic compounds are interchanged in a double displacement reaction. The form of these reactions is easy to recognize.

compound ac + compound bd à compound ad + compound bc

ac + bd à ad + bc

The following are some general types of double displacement reactions.

1. A reaction between an acid and a base yields a salt and water. Such a reaction is a neutralization reaction.

2KOH(aq) + H2SO4(aq) à K2SO4(aq) + 2H2O(l)

2. Reaction of a salt with an acid forms a salt of the acid and a second acid that is volatile.

2KNO3(aq) + H2SO4(aq) à K2SO4(aq) + 2HNO3(g)

This same reaction of a salt with an acid or base may yield a compound that can be decomposed. H2CO3, H2SO3, and NH3(aq) decompose to give a gas and H2O.

CaCO3(aq) + 2HCl(aq) à CaCl2(aq) + H2CO3(aq)

H2CO3(aq) à CO2(g) + H2O(l)

3. Reactions of some soluble salts produce an insoluble salt and a soluble salt.

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

 

e. Combustion. A substance combines with oxygen to form one or more oxides.

Combustion has this general form.

element/compound + oxygen à oxide(s)

a + O à aO

The following are some general types of combustion reactions.

1. A metal will combine with oxygen to produce a metallic oxide.

2Mg(s) + O2(g) à 2MgO(s)

2. In a substance containing hydrogen, water is always one of the products.

4NH3(g) + 7O2(g) à 4NO2(g) + 6H2O(l)

3. Hydrocarbons (compounds made of carbon and hydrogen) will react with oxygen to produce carbon dioxide (oxide of carbon) and water.

CH4(g) + 2O2(g) à CO2(g) + 2H2O(l)

C6H12O6(s) + 6O2(g) à 6CO2(g) + 6H2O(l)

4. Certain non-metals will burn with oxygen.

S(s) + O2(g) à SO2(g)

6.3 NATURE OF REACTIONS

In the previous chapters we have assumed that the reactions have gone to completion (reacted until at least one of the reactants was completely used up, and then stopped).

 Reactions tend to go to completion because of the formation of a gas (e.g., CO2, SO2), a precipitate (e.g., AgCl, PbSO4), or a slightly ionized substance (e.g., H2O, HF). Formation of these or similar species causes the elements from the initial reactants to be removed from the reaction.

2KClO3(s) à 2KCl(s) + 3O2(g)

(formation of a gas)

Na+(aq) + F-(aq) + H+(aq) + Cl-(aq) à Na+(aq) + Cl-(aq) + HF(aq)

(formation of a slightly ionized substance, hydrofluoric acid)

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

(formation of a precipitate)

REVERSIBLE REACTIONS

It has been determined experimentally that the conversion of some reactants to products is incomplete, regardless of the reaction time. Initially the reactants are present at a definite concentration. As the reaction proceeds, the reactant concentration decreases as the product is produced. However, a point is reached at which the reactant concentration levels off and becomes constant. The concentration levels for the reactants and products no longer change. A state of chemical equilibrium is established.

An example of a reaction that can proceed in either direction is the equilibrium system involving nitrogen, hydrogen, and ammonia gases. The reversible reaction is written as follows.

                      forward

N2(g) + 3H2(g) ßà 2NH3(g) + energy

                      reverse

A reversible chemical reaction is in chemical equilibrium when the rates of the opposing reactions are equal and the overall concentrations remain constant. Thus a state of chemical equilibrium is considered to be dynamic.

 

LE CHATELIER’S PRINCIPLE AND REACTANTS

Sometimes, systems initially at equilibrium are subjected to an outside influence or disturbance. Concentration, pressure, and temperature changes affect equilibrium because they produce a disturbance.

Le Chatelier’s principle states: If a system in equilibrium is subjected to a disturbance, the equilibrium will shift in an attempt to reduce the disturbance and regain equilibrium.

To see how these variables affect the equilibrium, consider the reaction between nitrogen and hydrogen to form ammonia.

If more reactant is added to the system in equilibrium, the reaction shifts to the right (the product side) and more product is formed. For example, in the ammonia equation

N2(g) + 3H2(g) ß à 2NH3(g) + energy

the addition of N2 disturbs the system. The system can relieve this disturbance by consuming N2. The system shifts to the right to consume N2, and in the process, produces more NH3. If a reactant is removed, the reaction shifts to the left. In the ammonia synthesis, if we remove some H2, the system can relieve the disturbance by producing H2. When the system shifts left to replace the missing H2, it also produces more N2 and consumes NH3.

Pressure affects only gaseous equilibrium systems. As pressure on the reactant gases is increased, the reaction shifts toward the side with the least volume. In the ammonia synthesis, an increase of pressure would shift the equilibrium to the right. In the process of shifting, four particles

(N2 + 3H2) are converted to two particles (2NH3). The number of particles colliding is thereby reduced, which also reduces the pressure. Lowering the pressure relieves the disturbance.

If temperature is increased, the reaction shifts in such a way that the endothermic reaction is favored. In the ammonia synthesis, the reaction from left to right is exothermic, while the reaction from right to left is endothermic. Consequently, a rise in temperature will shift the reaction to the left.

Catalysts speed up the reaction by lowering the activation energy, but this does not cause the equilibrium to shift. Catalysts do not increase the amount of product being produced. They simply just speed up the production process. In the same way, inhibitors do not shift the equilibrium. They just slow down the reaction.

In the end, the same amount of product will be produced with the inhibitor as without the inhibitor.