Lecture Summaries

7/31/2019 Lecture Summaries

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SOLUTIONS AND COLLOIDS

We more frequently encounter systems having more than one component. Such systemsfor

example, air, smoke, sea water, milk, blood, and rocks

aremixtures.

If a mixture is uniform throughout at the molecular level, we call it ahomogeneous mixture or,

more commonly, a solution. Air and sea water, for example, are bothsolutions. They are clear

and transparent.

In contrast, in most rocks, we can see distinct regions separated from each other by well-defined

boundaries. Such rocks areheterogeneous mixtures. Another example is a mixture of sand andsugar. We can easily distinguish between the two components; the mixing does not occur at the

molecular level. Thus mixtures are classified on the basis of how they look to the unaided eye.

Some systems, however, fall between homogeneous and heterogeneous mixtures. Cigarettesmoke, milk, and blood plasma may look homogeneous, but they do not have the transparency of

air or sea water. These mixtures are classified ascolloidal dispersions (suspensions).

Although mixtures can contain many components, we will generally restrict our discussion here

to two-component systems, with the understanding that everything we say can be extended to

multicomponent systems.

Systems containing more than one component are mixtures.o Homogeneous mixtures are uniform throughout.o Heterogeneous mixtures exhibit well-defined boundaries between phases.o When a solution consists of a solid or gas dissolved in a liquid, the liquid acts as

the solvent, and the solid or gas is the solute.o When one liquid is dissolved in another, the liquid present in greater amount is

considered to be the solvent. The solubility of a substance is the maximum amount of the substance that dissolves in a given

amount of solvent at a given temperature.


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o "Like dissolves like" means that polar molecules are soluble in polar solvents and thatnonpolar molecules are soluble in nonpolar solvents.

o The solubility of solids and liquids in liquids usually increases with increasingtemperature; the solubility of gases in liquids usually decreases with increasingtemperature.

Percent concentration is given in either weight per unit volume of solution (w/v), weight per unitweight of solvent (w/w), or volume of solute per unit volume of solution (v/v).o Molarity is the number of moles of solute in 1 L of solution.

Water is the most important solvent because it dissolves polar compounds and ions throughhydrogen bonding and dipole-dipole interactions.

o Hydrated ions are surrounded by water molecules (solvation layer) that move togetherwith the ion, cushioning it from collisions with other ions.

o Aqueous solutions of ions and molten salts are electrolytes and conduct electricity. Colloids exhibit a scattering of light called the Tyndall effect).

o They also exhibit a chaotic random motion, called Brownian motion.o Colloids are stable mixtures despite the relatively large size of the colloidal particles (1 to

1000 nm).o The stability results from the solvation layer that cushions the colloids from direct

collisions and from the electric charge on the surface of colloidal particles. Colligative properties are properties that depend only on the number of solute particles

present.o Freezing-point depression and osmotic pressure are two examples of colligative

properties.o Osmotic pressure operates across an osmotic semipermeable membrane that allows

only solvent molecules to pass but screens out all other molecules. In osmotic pressurecalculations, concentration is measured in osmolarity, which is the molarity of thesolution multiplied by the number of particles produced by dissociation of the solute.

Red blood cells in hypotonic solution swell and burst, a process called hemolysis.o Red blood cells in hypertonic solution shrink, a process called crenation.o Some semipermeable membranes allow small solute molecules to pass through along

with solvent molecules. In dialysis, such membranes are used to separate largemolecules from smaller ones

REACTION RATES AND CHEMICAL EQUILIBRIA

Some chemical reactions take place rapidly; others are very slow. For example, glucose and oxygen gas

react with each other to form water and carbon dioxide:

This reaction is extremely slow, however. A box of glucose exposed to in the air shows no measurable

change even after many years.

In contrast, consider what happens when you take one or two aspirin tablets for a slight headache. Very

often, the pain disappears in half an hour or so. Thus the aspirin must have reacted with compounds in

the body within that time.


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Many reactions occur even faster. For example, if we add a solution of silver nitrate (AgNO3) to a solution

of sodium chloride (NaCl), a precipitate of silver chloride (AgCl) forms almost instantaneously. The net

ionic equation is as follows:

This reaction is essentially complete in considerably less than 1 s.

The study of reaction rates is calledchemical kinetics. The rate of a reaction is the change in

concentration of a reactant (or product) per unit time. Every reaction has its own rate, which must be

measured in the laboratory.

Consider the following reaction carried out in the solvent acetone:

To determine the reaction rate, we can measure the concentration of the product, iodomethane (CH3I), in

the acetone at periodic time intervalssay, every 10 min. For example, the concentration might increase

from 0 to 0.12 mol/L over a period of 30 min. The rate of the reaction is the change in the concentration

of iodomethane divided by the time interval:

This unit is read "0.0040 mole per liter per minute." During each minute of the reaction, an average of

0.0040 mol of chloromethane, (CH3CI), is converted to iodomethane for each liter of solution.

The rate of a reaction is not constant over a long period of time. At the beginning, in most reactions, the

change in concentration is directly proportional to time. This period is shown as the linear portion of thecurve in Figure 7.1. The rate calculated during this period, called the initial rate, is constant during this

time interval. Later, as the reactant is used up, the rate of reaction decreases.


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The rates of chemical reactionsboth the ones that we carry out in the laboratory and the ones that

take place inside our bodiesare very important. A reaction that goes more slowly than we need may be

useless, whereas a reaction that goes too fast may be dangerous. Ideally, we would like to know what

causes the enormous variety in reaction rates. In the next three sections, we examine this question.

Some reactions are fast; others are slow.o The rate of a reaction is the change in concentration of a reactant or a

product per unit time.

o Reactions take place when molecules or ions collide.o The rate of a reaction increases with an increasing number ofeffective

collisionsthat is, collisions that lead to a reaction.

o The energy necessary for a reaction to begin is the activation energy.o Effective collisions are those that have

(1) more than the activation energy required for the reaction toproceed forward and

(2) the proper orientation in space. The lower the activation energy, the faster the reaction.

o An energy diagram shows the progress of a reaction.o The position at the top of the curve is called the transition state.

Reaction rates generally increase with increasing concentration andtemperature; they also depend on the nature of the reactants.

o The rate of some reactions can be increased by adding a catalyst, whichprovides a pathway with a lower activation energy.

The rate constant that gives the relationship between the rate of the reaction andthe concentration of the reactants remains constant at a constant temperature.

Many reactions are reversible and eventually reach equilibriumsome slowly,others quickly.

o At equilibrium, the forward and reverse reactions take place at equalrates, and concentrations do not change.


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o Every equilibrium has an equilibrium expression and an equilibriumconstant, K, which does not change when concentrations change but

does change when temperature changes.

o The equilibrium constant is large for some reactions and small forothers. Kand the rate of a reaction are not related.

Le Chatelier's principle tells us what happens when we put stress on a systemin equilibrium.

o The addition of a component causes the equilibrium to shift to theopposite side.

o The removal of a component causes the equilibrium to shift to the sidefrom which the component is taken.

o Increasing the temperature drives an exothermic equilibrium to the sideof the reactants.

o Addition of a catalyst has no effect on the position of equilibrium.

ACIDS AND BASES

We frequently encounter acids and bases in our daily lives. Oranges, lemons, and

vinegar are examples of acidic foods, and sulfuric acid is in our automobile batteries.

As for bases, we take antacid tablets for heartburn and use household ammonia as a

cleaning agent. What do these substances have in common? Why are acids and bases

usually discussed together?

In 1884, a young Swedish chemist named Svante Arrhenius (1859-1927) answered thefirst question by proposing what was then a new definition of acids and bases.

According to the Arrhenius definition, anacidis a substance that produces H3O+ ions

in aqueous solution, and abase is a substance that produces OH- ions in aqueous

solution.

This definition of acid is a slight modification of the original Arrhenius definition,

which was that an acid produces hydrogen ions, H+. Today we know that ions cannot

exist in water. A H+ ion is a bare proton, and a charge of +1 is too concentrated to

exist on such a tiny particle (Section 3.2). Therefore, a H+ ion in water immediately

combines with an H2O molecule to give ahydronium ion, H3O+.


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Apart from this modification, the Arrhenius definitions of acid and base are still valid

and useful today, as long as we are talking about aqueous solutions.

When an acid dissolves in water, it reacts with the water to produce H3O+. For

example, hydrogen chloride, HCl, in its pure state is a poisonous, choking gas. When

HCl dissolves in water, it reacts with a water molecule to give hydronium ion andchloride ion:

Thus a bottle labeled aqueous "HCl" is actually not HCl at all, but rather an aqueous

solution of and ions in water.

We can show the transfer of a proton from an acid to a base by using a symbol calledacurved arrow. First we write the Lewis structure of each reactant and product. Then

we use curved arrows to show the change in position of electron pairs during the

reaction. The tail of the curved arrow is located at the electron pair. The head of the

curved arrow shows the new position of the electron pair.

In this equation, the curved arrow on the left shows that an unshared pair of electrons

on oxygen forms a new covalent bond with hydrogen. The curved arrow on the right

shows that the pair of electrons of the H - C1 bond is given entirely to chlorine to

form a chloride ion. Thus, in the reaction of HCl with H2O, a proton is transferred

from HCl to H2O and, in the process, an bond forms and an bond is broken.

With bases, the situation is slightly different. Many bases are metal hydroxides, such

as KOH, NaOH, Mg(OH)2 and Ca(OH)2. These compounds are ionic solids and, when

they dissolve in water, their ions merely separate, and each ion is solvated by watermolecules (Section 6.6). For example,


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Other bases are not hydroxides. Instead, they produce ions in water by reacting with

water molecules. The most important example of this kind of base is ammonia, a

poisonous, choking gas. When ammonia dissolves in water, it reacts with water to

produce ammonium ions and hydroxide ions.

Ammonia is a weak base and the position of the equilibrium for its reaction with

water lies considerably toward the left. In a 1.0 M solution of NH3 in water, for

example, only about 4 molecules of NH3 out of every 1000 react with water to form

NH4+ and OH-. Thus, when ammonia is dissolved in water, it exists primarily as

NH3 molecules. Nevertheless, some OH- ions are produced and, therefore, NH3 is a

base.

Bottles of NH3 in water are sometimes labeled "ammonium hydroxide" or "NH4OH "

but this gives a false impression of what is really in the bottle. Most of the

NH3 molecules have not reacted with the water, so the bottle contains mostly NH3 and

H2O and only a little NH4+ and OH-.

We indicate how the reaction of ammonia with water takes place by using curved

arrows to show the transfer of a proton from a water molecule to an ammonia

molecule. Here, the curved arrow on the left shows that the unshared pair of electrons

on nitrogen forms a new covalent bond with a hydrogen of a water molecule. At the

same time as the new N

H bond forms, an O

H bond of a water molecule breaksand the pair of electrons forming the HO bond moves entirely to oxygen forming

OH-.

Thus ammonia produces an ion by taking from a water molecule and leaving behind.

By the Arrhenius definitions, acids are substances that produce H3O+

ions in

aqueous solution, and bases are substances that produce OH-ions in aqueous

solution.o A strong acid reacts completely or almost completely with water to form

H3O+

ions.


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o A strong base reacts completely or almost completely with water to formOH- ions.

The Bronsted-Lowry definitions expand the definitions of acid and base tobeyond water: An acid is a proton donor, and a base is a proton acceptor.

o Every acid has a conjugate base, and every base has a conjugate acid.o The stronger the acid, the weaker its conjugate base.o Conversely, the stronger the base, the weaker its conjugate acid.o An amphiprotic substance, such as water, can act as either an acid or a

base.

o In an acid-base reaction, the position of equilibrium favors the reactionof the stronger acid and the stronger base to form the weaker acid and

the weaker base.

The strength of a weak acid is expressed by its ionization constant, Ka.o The larger the value ofKa, the stronger the acid pKa = - log [Ka].o Acids react with metals, metal hydroxides, and metal oxides to

give salts, which are ionic compounds made up of cations from the base

and anions from the acid.

o Acids also react with carbonates, bicarbonates, ammonia, and amines togive salts.

In pure water, a small percentage of molecules undergo self-ionization:

As a result, pure water has a concentration of 10-7Mfor H3O+ and 10-7Mfor

OH-.

o The ion product of water, Kw, is equal to 1.0 X 10-14. pKw = 14. Hydronium ion concentrations are generally expressed in pH units, with pH = -

log [H3O+]. pOH = - log [OH-].

o Solutions with pH less than 7 are acidic; those with pH greater than 7 arebasic.

o A neutral solution has a pH of 7.o The pH of an aqueous solution is measured with an acid-base indicator

or with a pH meter.o Aqueous solutions of salts of strong acids and weak bases are acidic;

those of weak acids and strong bases are basic; and those of strong acids

and strong bases are neutral.

We can measure the concentration of aqueous solutions of acids and bases byusing a titration.


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o In an acid-base titration, a base of known concentration is added to anacid of unknown concentration (or vice versa) until an end point is

reached, at which point the acid or base being titrated is completely

neutralized.

A buffer does not change its pH very much when either hydronium ions orhydroxide ions are added to it.

o Buffer solutions consist of approximately equal concentrations of a weakacid and its conjugate base.

o The buffer capacity depends on both its pH and its concentration. The most effective buffer solutions have a pH equal to the pKa of

the weak acid.

The greater the concentration of the weak acid and its conjugatebase, the greater the buffer capacity.

The most important buffers for blood are carbonate andphosphate.

The Henderson-Hasselbalch equation is a mathematical relationship betweenpH, the pKa of a weak acid, and the concentrations of the weak acid and its

conjugate base:

ORGANIC CHEMISTRY (REVIEW)

Organic chemistryis the chemistry of the compounds of carbon. As you study Chapters 10-18 (organic

chemistry) and 19-31 (biochemistry), you will see that organic compounds are everywhere around us.

They are in our foods, flavors, and fragrances; in our medicines, toiletries, and cosmetics; in our plastics,

films, fibers, and resins; in our paints, varnishes, and glues; and, of course, in our bodies and the bodies

of all other living organisms.

Perhaps the most remarkable feature of organic compounds is that they involve the chemistry of carbonand only a few other elementschiefly, hydrogen, oxygen, and nitrogen. While the majority of organic

compounds contain carbon and just these three elements, many also contain sulfur, a halogen (fluorine,

chlorine, bromine, or iodine), or phosphorus.

Organic chemistry concentrates on carbon. The chemistry of the other elements comes under the field of

inorganic chemistry. Carbon is far from being among the most abundant elements in Earth's crust. In


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1896) put forth a new definition-organic compounds are those containing carbon-and his definition has

been accepted ever since.

A second reason for the study of carbon compounds as a separate discipline is the sheer number of

organic compounds. Chemists have discovered or made more than 10 million of them, and an estimated

10,000 new ones are discovered or prepared in the laboratory each year. By comparison, chemists havediscovered or made an estimated 1.7 million inorganic compounds. Thus approximately 85% of all known

compounds are organic compounds

A third reasonand one particularly important for those of you who plan to study biochemistryis that

carbohydrates, lipids, proteins, enzymes, nucleic acids (DNA and RNA), hormones, vitamins, and almost

all other important chemicals in living systems are organic compounds. Furthermore, their reactions are

often strikingly similar to those observed occurring in test tubes. For this reason, knowledge of organic

chemistry is essential for an understanding of biochemistry.

One final point about organic compounds. They generally differ from inorganic compounds in many of

their properties, some of which are shown in Table 10.1. Most of these differences stem from the factthat the bonding in organic compounds is almost entirely covalent, while most inorganic compounds

have ionic bonds.

Of course, there are many exceptions to the generalizations in Table 10.1.Some organic compounds

behave like inorganic compounds, and vice versa, but the generalizations remain largely true for the vast

majority of compounds of both types.


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Organic chemistry is the study of compounds containing carbon.

o Chemists obtain organic compounds either by isolation from plant and animal sources

or by synthesis in the laboratory.

Carbon normally forms four bonds and has no unshared pairs of electrons.

o Its four bonds may be four single bonds, two single bonds and one double bond, or one

triple bond and one single bond.

o Nitrogen normally forms three bonds and has one unshared pair of electrons.

o Its bonds may be three single bonds, one single bond and one double bond, or one triple

bond.

o Oxygen normally forms two bonds and has two unshared pairs of electrons.

o Its bonds may be two single bonds or one double bond.

Functional groups are sites of chemical reactivity; a particular functional group, in whatever

compound it is found, always undergoes the same types of reactions. In addition, functional

groups are the characteristic structural units by which we both classify and name organic

compounds.

o Important functional groups include the hydroxyl group of 1o, 2o, and 3o alcohols;

o the amino group of 1o, 2o, and 3o amines;

o the carbonyl group of aldehydes and ketones;

o and thecarboxyl group of carboxylic acids.

---- End of Lecture Summaries Series (#1)----

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