Tro Cpt 7 Introductory Chemistry Class Notes

Chemical Reactions

WRITING CHEMICAL EQUATIONS A chemical equation is the symbolic representation of a chemical reaction utilizing chemical formulas-it is a chemistry "short-hand".

A + B —> C + D

is "read" as "the reaction of A with B "goes to" "gives" "react to form" "yields" C and D. A and B are called the REACTANTS, the arrow is read "goes to", "react to form", "yields" or "gives" and C and D are called the PRODUCTS.

It is customary to write BALANCED equations such that the NUMBER and KIND of atoms appearing on the RIGHT (or reactant) side of the arrow are EQUAL (in number and kind) to those atoms that appear on the LEFT (or product) side.

COEFFICIENTS (numbers that appear before the formula weight of a material in the equation) are added to BALANCE the equation-make the number and kind of atoms (not compounds) that react be the same in number and kind as appear as products.Thus, it is the ATOMS that are balanced-not molecules, compounds.

(Thus, for the balanced equation 2H2 + O2 2H2O the 2 in front of H2 and H2O is a coefficient and says that to have a balanced equation we need 2 H2 units and 2 H2O units. There is an understood “1” in front of the O2 that is not written but we know that it is there. Notice that we can also write this equation by changing the order of hydrogen and oxygen as follows. O2 + 2H2 2H2O. We can also write this going in the opposite direction just as long as the products and reactants are the same. 2H2O 2H2 + O2.)

Further, the "form" or "phase" of the reactants and products may also be indicated where (g) = gas; (l) = liquid, (s) = solid, and (aq) indicates an aqueous (water) solution.

Thus, for the reaction of diatomic (contains TWO atoms of oxygen) oxygen gas and diatomic (contains TWO atoms of hydrogen) hydrogen gas to form gaseous water we have

O2 + H2 -------> H2O

as the "unbalanced" equation where "O2 " and "H2 " are REACTANTS and the "H2 O" is the PRODUCT.

We can "BALANCE" the reaction by placing a "2" in front of "H2 " and in front of "H2O". Since "O2 " appears in the balanced equation only once, by agreement, we "understand" that a "1" is there and do not place the "1" in the equation.

Thus the "completed" equation (including the phases) is as follows: O2 (g) + 2H2 (g) ------> 2H2O (g)Practice "balancing" and writing equations. This is a skill that will be required throughout your chemistry classes-as well as other classes that include chemical reactions.

Lets practice balancing some reactions. To do this we will simply put the equation and then the balanced equation later.

C2H6 + O2 CO2 + H2O

This is an example of a combustion or burning reaction of hydrocarbons, materials with only hydrogen and carbon in them. The products are carbon dioxide and water. The major products are carbon dioxide and water for compounds that contain only C, H, and O such as sugars, wood, and paper.

The almost balanced reaction is

C2H6 + 3.5 O2 2 CO2 + 3 H2O

Balanced reactions must have whole numbers in front of each compound or element. Where there is only one we do not place a one but it is understood to be there. Here we have a 3.5 in front of the O2. We can get rid of the fraction by simply multiplying by two giving us

2 C2H6 + 7 O2 4 CO2 + 6 H2O This is now a balanced reaction.

We might view

4 C2H6 + 14 O2 8 CO2 + 12 H2O

as a balanced reaction since there are no fractions but it is not because we can divide everything by two and get whole numbers.

Thus, a balanced reaction must have only whole numbers and be the smallest combination of small whole numbers.

What is balanced in a chemical equation? It is the number of each kind of element that is balanced.

Let us try

K + O2 K2O

The balanced reaction is

4 K + O2 2 K2O

AQUEOUS SOLUBILITY

A solution contains a solute and solvent. For instance, common table salt, NaCl, dissolves in water forming sodium ions and chloride ions. Solutions that have water as the solvent are called aqueous solutions.

Solutes that dissolve producing ions are called electrolytes. Those that produce lots of ions such as NaCl are called strong electrolytes. Those that produce only a few ions such as acetic acid are called weak electrolytes. And, those that product no ions either because they are covalently bonded such as sugar or because they are not soluble in water are called non-electrolytes. Those solutes that dissolve are said to be soluble and those that do not dissolve to an appreciable extent are called insoluble. When a material is formed within a solution and that material is not soluble, it falls out from solution, it precipitates from solution.

We will use a table that contains which compounds are soluble or not soluble (insoluble) to help us in predicting if an ionic compound is soluble. You will need to memorize it and be able to predict what compounds are soluble or not soluble.

Solubility Rules

Compounds that contain the following Exceptions ions are most SOLUBLE.

Li+1, Na+1, K+1, Rb+1, NH4+1 None

(Hint-these are Group IA ions plus)

NO3-1, C2H3O2

-1 None

Cl-1, Br-1, I-1 (Hint-Group VIIA ions) When combined with Ag+1, Hg+1, Pb+2 compound is not soluble

SO4-2 When combined with Sr+2, Ba+2, Ca+2, Pb+2

compound not soluble(Hint-mostly Group IIA)

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Compounds containing the following Exceptionsions are mostly INSOLUBLE.

OH-1, S-2 When combined with Li+1, Na+1, K+1, Rb+1, NH4

+1 (Hint-Group IA)

When combined with Sr+2, Ba+2, Ca+2 soluble

CO3-2, PO4

-3 When combined with Li+1, Na+1, K+1, Rb+1, NH4

+1 (Group IA)

Be able to predict the solubility when given a compound.

PREDICTING PRECIPITATION REACTIONS

NaBr + AgNO3 Need to see what possible combinations of cations and anions are.

Here there are four. Two have already been given, NaBr and AgNO3. From the solubility table we see that both NaBr and AgNO3 are soluble. The other combinations are NaNO3 which is a sodium salt so it is soluble. The other combination is AgBr. We see that halides such as Br-1 are generally soluble but with some exceptions and one exception is when combined with Ag+1. thus, AgBr is not soluble and we will have a precipitation of the AgBr since it is not soluble.

MOLECULAR AND IONIC EQUATIONS

MOLECULAR EQUATIONS-where the substances are written as if they were molecular in nature even though they may actually exist in solution as ions.

CaCl2 (aq) + Na2CO3 (aq) -----> CaCO3 (s) + 2NaCl (aq)

IONIC EQUATIONS-where all the compounds that form ions in aqueous solution are written as ions. Since acetic acid is a weak acid it is NOT written as an ionic species and as a solid that is not soluble, like AgCl , does not form ionic species it is also not written as an ion. By comparison, H Cl is a strong acid and does appreciably form ions in solution, it is written as ions and NaCl is a soluble salt, so it does form ions when dissolved in water, it is written as ions.

Ca+2 + 2Cl- + 2Na+ + CO3-2 -----> CaCO3 (s) + 2Na+ + 2Cl-

NET IONIC EQUATIONS-where dissolved ions that are NOT involved (that is they are not changed from each side of the arrow “—>”) in the reaction are OMITTED.

Ca+2 + CO3-2 -----> CaCO3 (s)

Ions such as Na+, OH- that are not changed in the chemical reaction are called SPECTATOR IONS. Such “Spectator Ions” are INCLUDED in the MOLECULAR equations but ELIMINATED in IONIC equations.

ACIDS-BASES Rx Aq-6

Arrhenius defined an ACID as a substance that increases the acidity, proton or hydronium content of an aqueous (water) solution when dissolved and a BASE is defined as a substance that increases the basicity or hydroxide content of the aqueous solution when it is dissolved in water.

The terms STRONG ACID, WEAK ACID AND STRONG BASE AND WEAK BASE are directly measurable using conductivity since a STRONG ACID largely dissociates, ionizes forms ions upon dissolving in water; a WEAK ACID is a substance that gives only a small fraction of ions when dissolved in water; a STRONG BASE is a substance that almost totally dissociates forming ions when dissolved in water and a WEAK BASE is a material that forms only a small number of ions when dissolved in water.

Know material in Table 5.2.Memorize what a strong acid is-rest will be weak.

HA -----> H+ + A-BOH ----> B+ + OH-

*where most stays as HA and BOH -un-ionized- they are considered as a WEAK acid or base;

*where most ionizes-they are considered a STRONG acid or base.

ACIDIC/BASIC COMPOUNDS

Acids form acidic solutions while bases form basic solutions. Some compounds form acids and bases when dissolved in water. According to Arrhenius, an acid is a compound that when dissoled increases the proton, hydrogen ion, or hydronium ion content. Note: These three names refer to the same thing- proton, hydrogen ion, hydronium ion.

Thus, nitric acid is an acid since it increases the hydrogen ion, etc. content when dissolved in water.

HNO3 H+ + NO3-1

It is said to be a strong acid because essentially every nitric acid molecule ionizes, forms ions, when dissolved in water; essentially every nitric acid molecule dissolves forming a proton and a nitrate ion. It is also called a strong electrolyte.

Strong acids are also strong electrolytes. They include nitric acid, sulfuric acid, and the HX (group 7A) halides except HF. Weak acids include acetic acid, carbonic acid, and oxalic acid.

Bases, according to Arrhenius, are compounds that when added to water increase the concentration of the hydroxide, OH-1, ion. Thus, sodium hydroxide is a base since it gives an increase in the hydroxide ion when it is dissolved in water.

NaOH Na+ + OH-1

Sodium hydroxide is a strong base since it gives essentially one OH-1 for each sodium hydroxide dissolved. It is also a strong electrolyte. Strong bases are mainly the hydroxides of Group IA and IIA metals such as KOH, and Ba(OH)2. Metal compounds containing the hydroxide, OH-1, unit are bases. But there are lots of materials that contain OH groups that are not bases such as most sugars and alcohols.

Acid Rain

Acid rain is the result of the formation of acids from the reaction of sulfur oxides and nitrogen oxides with water.

Sulfur Trioxide + Water —> Sulfuric AcidSO3 + H2 O —> H2 SO4

Nitrogen Dioxide + Water —> Nitric Acid + Nitrous Acid2NO2 + H2O —> HNO3 + HNO2

Neutralization Reactions

Acids and bases react to give a salt and water. They do so until all of the acid or base is used up. This is a general reaction and one that you should remember.

Acid + Base Salt + H20

The reaction of the

H+ + OH-1 is the net reaction and the product is H2O

HCl + NaOH NaCl + H2O

OXIDATION-REDUCTION OR REDOX REACTIONS

Redox reactions involve an EXCHANGE of ELECTRONS. Where there is oxidation there must also be reduction. It is a series of opposites. To help remember the definitions we can use the oxidation or reaction of iron with oxygen. This, in fact was the basis of redox reactions so it is appropriate that we use it as an example. Before we return to this let us see how we determine the oxidation number. We determined oxidation when we were naming the transition compounds.

We assign, according to some rules that are related to the chemical tendencies and placement within the periodic chart of the specific atoms. These rules allow us to assign what is called an OXIDATION NUMBER to atoms within a compound. These rules are as follows:1. The oxidation number of an element is zero.2. For monoatomic ions, the oxidation number is the charge on the ion. Thus Ca+2 has an oxidation number of plus two. For Cl-1 the oxidation number is minus or negative one.3. Polyatomic ions, such as the hydroxyl radical, and the dichromate radical have, as a radical, their assigned charges-thus for the hydroxide it is a minus one.4. a. Group I A and II A are + 1 and +2. b. Hydrogen can be either +1 or -1 depending on what it is combined with. If it is

with a metal such as in NaH it is a -1 and for non-metals like in H Cl it is a +1.c. Oxygen is generally a -2 except in peroxides where it is -1.d. Halogens are normally -1 unless combined with another non-metal where it

may be something else. 5. The sum of the oxidation numbers in a neutral compound is zero. For a charged radical or ion, the sum of the oxidation numbers will be the charge associated with the radical (as in 3 above)

Examples-

HClO3 H = +1, O = -2/each; therefore Cl = 6-1 = 5

KMnO4 MnO2

K2CrO4 CrO CrO2 CrO3 K2Cr2O7

H3PO4 H3PO3 P4O10 P4O6

H2SO4 H2SO3 H2S SO2 SO3

NO NO2 N2O4 HNO3 HNO2

FeCl2 FeCl3

CuCl2 Cu2Cl2

CaNaPO4 NaHCO3 SO4 -2 ClO4

-1

CH3Cl CH2Cl2CHCl3 CCl4 CO2 CO C

Following is brief summary of rules for determining Oxidation Number.1. Oxidation number of an element in the elementary state is zero. Fe, O2, Al, Xe, Cl2

2. Oxidation number of a charged cation or anion is the charge on the cation or anion. Na+=1, Ba+2=2, Cl-=-1, Fe+2=2, H+=13. In almost all metal-containing compounds Halides=-1, Group IA=1, Group IIA=2, H=1, O=-24. Sum of all oxidation numbers is the charge on the species-if neutral then zero, or charged species it is the net charge of the species.

Now we return to our reaction of iron with oxygen to form iorn (II) oxide, FeO.

2Fe + O2 2FeO

The oxidation number for Fe is zero, for O2 is zero and for Fe in FeO it is +2, and for O in FeO it is -2. We see that the oxidation number for iron has increased from 0 to +2 as it changed from metallic Fe on the left side of the equation to FeO on the product side of the equation. Iron has been oxidized, reacted with oxygen. So in oxidation there is an increase in oxidation number. Conversely, the oxidation of O has decreased from 0 to -2. Reduction is then a decrease in oxidation number. This is our first opposite.

Now, let us divide the equation into two half cells, one dealing with oxidation and one dealing with reduction.

Fe Fe+2

To balance this half cell we need to have the charge be the same on each side of the arrow. To do this we add two electrons to the product side giving

Fe Fe+2 + 2e-

Thus, iron as lost two electrons to become Fe+2. Overall, for each increase or loss of oxidation number there will be a loss or gain in one electron. Thus, in oxidation there is a loss in electrons.

For oxygen we have the reduction half cell.

O2 2O-2

Or a decrease in two oxidation numbers for each oxygen atom for a total change of 4 oxidation number change. Again, for the reactant and product side to be balanced we need the charges to be the same, not necessarily zero. For this to occur we can add four electron to the reactant side giving

4 e- + O2 2 O-2

Thus, in reduction there is a gain in electrons.

Another set of opposites. For oxidation there is an increase in oxidation number but a decrease in electrons. For reduction there is a decrease in oxidation number but an increase in electrons.

Finally, our last set of opposites. An oxidizing agent is the material that has been reduced, here oxygen, while the reducing agent is the material that has been oxidized, here iron.

SummaryOXIDATION REDUCTION

Increased oxidation number Decreased oxidation numberLoss in electrons Gain in electronsCalled reducing agent Called oxidizing agent

But ButIt is oxidized It is reduced

For the reaction of calcium with chlorine gas we have

2Ca + Cl2 2CaCl2

OXIDATION-is defined as a gain in oxidation or loss in electrons. Ca -----> Ca++ + 2e-

Here calcium increases its oxidation number from zero to plus two and this is caused because of the loss of two electrons. The reaction of calcium metal going to calcium plus two ion is called OXIDATION and the chemical that causes or brings about this loss in electrons is called an OXIDIZING AGENT because it is the "agent" or cause of the elemental calcium losing its electrons.

REDUCTION-is the opposite of OXIDATION and is defined as the LOSS in oxidation number or the GAIN in electrons. Cl-Cl + 2e- ------> 2Cl-

Thus, diatomic chlorine decreases, lowers its oxidation number from zero to minus one per chlorine atom. Diatomic chlorine is said to be REDUCED and it has been reduced by an agent or chemical that supplies the needed electrons. The agent that supplies these electrons is called an REDUCING AGENT.

If the two reactions cited above were brought together, Ca would be said to be OXIDIZED by Cl-Cl which would be called the OXIDIZING AGENT. Similarly, Cl-Cl would be said to be REDUCED and the agent or cause of this reduction or gain in electrons (that is the chemical that supplied the necessary electrons) is Ca and it is then called the REDUCING AGENT.

Of interest, when asked what is the reducing/oxidizing agent, who loses/gains electrons and who gains/loses electrons the answer always is on the product side, never on the product side.

In combustion reactions with hydrocarbons and materials that contain only C, H, and O the main products are carbon dioxide and water.

Thus, the product of methane and oxygen is carbon dioxide and water.

CH4 + O2 CO2 + H2O

Practice balancing combustion reactions such as these.

For methane the balanced reaction is

CH4 + 2O2 CO2 + 2H2O

WHY REACTIONS ARE DRIVEN TO OCCUR

ELEMENTARY MY DEAR DR. WATSON (all in an aqueous environment).

A. FORMATION OF A PRECIPITATE-To be able to correctly "predict" or identify which set of compounds will produce a precipitate, you need to know the general solubilities of common ionic compounds. The solubility table gives a brief description of such compounds and should be memorized.

INSOLUBLE and INSOLUBLE are relative terms that also require designating the LIQUID or SOLVENT. Here we will be talking about water-but it could be acidic or basic or neutral. Ag+ + Cl- -----> AgCl (s) but some reactions "will not go" or do not occur.NaCl + KBr -----> simply gives Na+, Cl-, K+, Br- ions in solution-so that not only are NaCl and KBr soluble in water, but so also are NaBr and KCl.

B. NEUTRALIZATION REACTIONS-reaction of an acid and a base gives a salt and water.Acid and base can be strong or weak-same-salt and water. HCl + KOH -------> KCl + HOH ACID + BASE -------> SALT + WATERWe can write the acid portion as H+ or as a hydronium, H3O+, ion.Some neutralization reactions are a bit more complex and occur in stages. Thus reaction of phosphoric acid with sodium hydroxide occurs giving a complex of products-always a salt and water-the proportion of products are dependent on the relative amounts of sodium hydroxide and phosphoric acid.

NaOH + H3 PO4 ----> NaH2PO4 ----> Na2HPO4 ----> Na3PO4

C. EXCHANGE OF ELECTRONS-REDOX REACTIONS

D. GAS FORMATION

REACTION RATES

Reaction rates tell about the speed of reactions; they tell us how fast reactants are used up and products are formed. Reaction rates increase with increase in temperature. For a reaction to occur there must be sufficient energy. The energy barrier needed to be overcome, the barrier to getting the reactants to change into products, is called the activation energy or energy of activation. Thus, to get H2 and O2 to react forming H2O sufficient energy must be available to push them together allowing the H-H and O-O bonds to break and reform to give H-O-H bonds. Reaction rate increases as temperature increases since increased temperature means that more of the reactants will have sufficient energy to get over the activation energy barrier. Conversely, reaction rate decreases as temperature decreases. This is another example of a direct relationship.

To react the molecules must come into contact, often 1010 times or more and they must be in the right geometry. Catalysts generally place molecules in the right geometry so that this hurdle is lowered acting to lower the overall activation energy. Since molecules must come into contact to react, reaction rate is increased as the concentration of reactants increases.

Reactions are either net exothermic, giving off energy like the burning of natural gases and petroleum products, or endothermic, taking in a net amount of energy (cold packs). Below are two “reaction profiles.” The one on the left is for an endothermic reaction and the one on the right is for an exothermic reaction. The heat of reaction which tells if it is exothermic or endothermic is the potential energy difference between the reactants and products. If the potential energy for the products is above that of the reactants, then the reaction is endothermic; conversely, if the potential energy for the products is below that of the reactants, then the reaction is exothermic and a net amount of energy is given off.

Figure 9.1 Reaction profiles showing activation energies and heats of reaction.

The energy difference between the reactants and the activation state is called the activation energy or energy of activation. The reaction rate or speed of a reaction is related to the activation energy. In general, if the activation energy is relatively low or small, the reaction will be fast or rapid; conversely, if the activation energy is relatively high or great, the reaction will be slow. The relationship is called an exponential relationship so that small activation energy differences are reflected in large changes in the rate.

The rate of a reaction is related to the amount of energy or heat available. As the energy available to run a reaction becomes greater, more molecules will have the necessary energy to overcome the activation energy and the faster the reaction. Again, the relationship between reaction rate and temperature is exponential so that relatively small temperature variations will greatly influence the rate.

Enzymes are proteins that are natural catalysts. They lower the energy of activation through assuring the correct geometry (orientation) as well as holding the reactants in the general vicinity of one another so that they do not have to “find” one another. These catalysts can have a phenomenal affect on the reaction rate. One reaction takes about a billion years to occur but in the presence of the correct catalysts occurs in picoseconds, an increase in about 1017.