Basic Chemistry I (Print)

8/3/2019 Basic Chemistry I (Print)

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Basic Chemistry I


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MATTER AND ENERGY

Matter

All the materials that interest

chemistsin fact, all the things we

can see or touch or feelare

examples ofmatter, whether they bebooks, pencils, telephones,

hamburgers, or people.

Matter is defined as anything that

takes up space and has mass. In

setting down this definition, we are

very careful to specify the term mass

rather than weight, even though we

often use the terms as if they were

interchangeable. Mass and weight

are not really the same.

A Ping-Pong ball moving at 30 km h -

1 (30 kilometers per hour), for

example, is easily deflected by a softbreeze, but a cement truck moving at

the same speed is not.

Quite clearly, the mass of the cement

truck is considerably greater than

that of the Ping-Pong ball.


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The term weight refers to the force

with which an object of a certain

mass is attracted by gravity to the

earth or to some other body that it

may be near, such as the moon.

Force and mass are related to each

other by Newton's equation

(Newton's law).

F = ma

whereF= force, m = mass, and

a = acceleration. In order to accelerate a

body, a force must be applied to it. When an

object is dropped, it accelerates because of

the gravitational attraction of the earth. An

object resting on the earth or moon exerts a

force (its weight, W) that is equal to its

mass, m, multiplied by the acceleration due

to gravity,g, that is,

W= mg

For example, at the earth's surface,g

= 9.81 m s-2.

Thus a one-kilogram mass would

have a weight (or experience a force

downward at the earth's surface) of

For convenience, the derived SI unit

of force or weight, possessing units 1

kg m s

-2

, is defined as the newton(N).

A one-kilogram mass has a we 9.81

N at the earth's surface. You can

experience the magnitude of a one

force by placing a large lemon (or

any object with slightly more than

100 g mass in your hand; the "push"

downward represents a one-newton

force..' moon the gravitational

acceleration is only about one-sixth

of that on earth, so the same object

weighs only one-sixth as much on

the moon as on the earth, even

though its mass is the same in both

locations.

Even on the earth the value ofghas

been found to vary slightly from

place to place. This means that if

very precise measurements are made,

an object's weight also varies slightly

according to its location on the earth.

Because of this, we specify an

object's mass instead of its weight

when we wish to report the quantity

of matter in the object.

The measurement of mass (a process,

oddly enough, called weighing) is

actually performed by comparing the

weights of two objects, one of

known mass, the other of unknown

mass.

The apparatus used for this is called

a balance. Figure 1.7 is a drawing of

a traditional two-pan balance. The

object to be weighed is placed on the

left pan of the balance and objects of

known mass are added to the other

until the pointer comes to the center

of the scale.

At this point the contents of both

pans weigh the same, and since they

both experience the same

gravitational acceleration, both pans

contain equal masses.


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In chemistry, as mentioned

previously, we generally measure

mass in grams.

Most balances found in labs today

function on the same principle. Themost modern electronic balances,

operate on a somewhat different

principle. They are rugged, fast, and

convenient, and make the

measurement of mass a routine

laboratory operation.

Energy

When chemical changes occur, theyare almost always accompanied by

either an absorption or release of

energy.

These energy changes tell us a great

deal about the nature of the

chemicals that are reacting, and

equally important to us, certain

chemical reactions provide the

energy that our bodies and oursociety need in order to function.

Therefore, an understanding of what

energy is and how it can be

transferred from one object to

another is as important to those who

study chemistry as an understanding

of matter itself.

Energy itself is a difficult concept tocomprehend, because energy is so

different from matter. You can't put

energy in a bottle to examine it. All

you can do is examine the effects of

energy on objects.

Energy is usually defined as the

capacity to do work or to make

things happen. When an object has

energy, it can affect other objects by

doing work on them.

A moving car possesses energy

because it can do work on another

car by moving it some distance in a

collision.

Coal and oil possess energy because

they can be burned and the heat that

is liberated can be harnessed to do

work.

Because energy can be transferred

from one object to another as work,

the units of energy and work are the

same.

An object can possess energy in two

ways: as kinetic energy and as

potential energy. Kinetic energy

(K.E.) is associated with motion and

is equal to one-half of an object'smass, m, multiplied by its velocity,

v, squared.

Thus, we see that the amount of

work a moving body can do depends

on both its mass and its velocity.

For example, a truck moving at 30 km h-1

can do more work on the rear end of a car

than a bicycle moving at the same speed.We also know that a truck moving at 80 km

h-1 can do more work on a car than one

traveling at only 5 km h-1.

Potential energy (P.E.) represents

energy a body possesses because of

the attractive or repulsive forces it


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experiences with other objects. If

there are no attractive or repulsive

forces, the body does not possess

potential energy.

As we have indicated, chemicalsubstances contain energy that can

be released through chemical

reactions. For example, when wood

burns, it reacts with oxygen from the

air. As the products of combustion

are formed, rather sizable quantities

of heat are released. In this case, the

wood and the oxygen are sources of

potential energy, sometimes called

"chemical energy," and the reaction

causes this energy to be released into

the surroundings as heat.

The quantity of energy released or

absorbed in a chemical reaction

depends on the amounts of materials

that react.

The burning of a match, for example,

releases only a very small quantity of

energy, but a large bonfire produces

much more.

When a change (either chemical or

physical) releases energy to the

surroundings, it is called an

exothermic change. Changes that

absorb energy are termed

endothermic.

It is usually easy to recognize

whether a process is exothermic or

endothermic; an exothermic change

raises the temperature of its

surroundings, and an endothermic

change causes its surroundings to

become cool

The total quantity of energy that an object

possesses is equal to the sum of its kinetic

energy and its potential energy. In anisolated system, such as we believe our

universe to be, the total quantity of energy is

constant. This leads to a very important

physical law, called the law of conservation

of energy, which states that energy can be

neither created nor destroyed. It only can be

changed from one kind of energy to another.

Units of energy

Energy is transmitted from one body

to another in a variety of ways, for

example, as light, sound, electricity,

and/or heat. These various forms of

energy can be converted from one to

another and therefore are ultimately

equivalent.

The SI (System International) unit

for energy is the joule (J), which isdefined as 1 kg m2/s2 in terms of the

SI base units. It represents the kinetic

energy possessed by an object with a

mass of 2.0 kg traveling at 1 m/s.

In referring to energies involved in

chemical reactions, it is useful to use

the term kilo joule (kj). One kilo

joule is equal to 1000 joules (1 kj =

1000 J).

All forms of energy ultimately are

transformed into heat, and when

chemists measure energy, it is

usually in the form of heat.


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We know that heat always flows by

itself from a region of high

temperature to one of low

temperature, as when a hot cup of

coffee becomes cool.

We also know that the rate of heat

transfer depends on the difference in

temperature between two objects. If

we want to warm something slowly,

we place it in contact with an object

just slightly warmer; if we wish to

heat it quickly, we place it in contact

with something very hot.

The original definition of the unit of

heat energy called the calorie

(symbolized cal) relied on the

concept of temperature. The calorie

was defined as the quantity of heat

needed to raise the temperature of 1

gram of water, initially at 15C, by

1C.

The kilocalorie (kcal), like the

kilojoule, is a more appropriately

sized unit for dealing with energy

changes in chemical reactions.

For example, we might measure the

heat liberated during the combustion

of natural gas, which is composed

principally of the compound

methane. The measurement of heat

involves the concept of temperature,

which we discussed earlier.

It is important to remember that heat

is not the same as temperature. Heat

is energy, while temperature is a

measure of the intensity of heat or

hotness. Another useful definition of

temperature is that it defines the

direction and rate of heat flow.

It has also been the unit used in

nutrition for reporting the energy

content of foodsthe Calorie(written with a capital C) is the same

as a kilocalorie. Thus, if we read that

a serving of mashed potatoes

contains 230 Cal, we are being told

that 230 kcal of energy are liberated

when the body metabolizes this food.

Until recently, nearly all the

chemical scientific literature used the

calorie (or kilocalorie) in reportingenergy changes. With the

introduction of SI units, the joule (or

kilojoule) is now preferred and the

calorie has been redefined in terms

of this SI unit. The calorie and

kilocalorie are defined exactly by the

relationships

1 cal = 4.184 J

1 kcal = 4.184 kJ

Thus in terms of SI, one Calorie of

food energy is more properly

expressed as 4.184 kJ.

PROPERTIES OF MATTER

Objects are recognized by their

characteristics, or properties. For

example, Iron has shiny surfaces,

pull by magnet, react with oxygen

produce rust. If we burn magnesium,

bluish white color is emitted. Boiling

point of alcohol is lower than water.

Luster, color, magnetism, the

tendency to corrode, boiling point


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are just some of the many properties

that we use to recognize and classify

different samples of matter.

These properties themselves can be

divided into two broad categories:extensive properties, which depend

on the size of a sample of matter; and

intensive properties, which are

independent of the sample size. Of

the two, intensive properties are the

more useful, because a substance

will exhibit the same intensive

property regardless of how much of

it we examine.

Examples of extensive properties are

mass and volumeas the amount of

a substance increases, its mass and

volume also increase. Some

examples of intensive properties are

melting point and boiling point.

Another intensive property is

density, which is defined as the ratio

of an object's mass to its volume.

Liquid water, for instance, has a

density of 1 g/ cm3. This means that

if we had 1 g of water, it would

occupy a volume of 1 cm3. If we had

20 g of water, it would occupy a

volume of 20 cm3, but the ratio of

the water's mass to its volume is the

same: 20 g/20 cm3 = 1 g/1 cm3 = 1 g

cm-3. Notice that we have created an

intensive property by taking the ratio

of two extensive ones. Later in our

discussion of chemistry, we will

encounter quite a few other intensive

properties defined in a similar way.

Normally, when a substance is

heated or cooled, its volume expands

or contracts. This means that the

object's mass is packed into either a

larger or smaller volume, so the

density also changes with

temperature.

Therefore, for very accurate work,

the temperature corresponding to a

reported density must be specified.

For example, at 25.0C (room

temperature) the density of water is

0.9970 g cm-3, while at 35.0C its

density is 0.9956 g cm-3. (For most

purposes, we can take the density of

water to be 1.00 g cm-3.)

Density provides a relationship

between an object's mass and its

volume. It tells us, in this instance,

that if we have 1.00 cm3 of copper,

its mass is 8.96 g. It also tells us that

if we have 8.96 g of copper, its

volume is 1.00 cm3. We call/use the

density as a conversion factor in two

ways

Since we are given the mass of

copper, we must multiply by the

second factor to eliminate the units

grams.

A property closely related to density

is specific gravity (often abbreviated

sp.gr.), or relative density, which is

defined as the ratio of the density of

a substance to the density of water.

In this way two tablesone

containing the specific gravities of

substances and the other containing


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the density of water in a variety of

unitstake the place of the many

tables that would otherwise be

needed to express the densities of

substances in all those different

units.

EXAMPLE 1.9 PROBLEM

FINDING SPECIFIC GRAVITY

Hexane, a solvent used for rubber

cement, has a specific gravity of

0.668. What is the density of hexane

in g cm-3 and in kg m-3. Water has a

density of 1.00 g cm-3 or 1.00 x 103

kg m-3.

SOLUTION

By definition:

dhexansp. gr. hexane = ------------

dwater

Therefore

(sp. gr. hexane) x dwater= dhexan

In units of grams per cubic centimetre

dhexan = (0.668) x (1.00 g/cm3) = 0.668 g

cm-3

Notice that in units of g cm -3 the

numerical values of density and

specific gravity are the same.

In units of kilograms per cubic

metre,

dhexan =

(0.668) x (1.00 x 103 kg/m3) = 6.68 x 103

kg m3

Specific heat capacity

An intensive property of matter associatedwith energy is specific heat capacity

(sometimes termed specific heat), the

quantity of heat required to raise the

temperature of 1 g of a substance by 1C.

The specific heat capacity of water is 4.184

J g-1C-1. Most other substances have

smaller specific heats. Iron, for example, has

a specific heat of only 0.452 J g-1oC-1. This

means that it takes less heat to raise thetemperature of 1 g of iron by 1C than it

does to cause the same temperature change

for a gram of water. It also means that a

given quantity of heat will raise the

temperature of 1 g of iron more than it will

raise the temperature of 1 g of water.

The large specific heat capacity of

water is responsible for the

moderating effect the oceans have onweather, since they cool very slowly

in winter and warm up slowly in the

summer.

Air moving over the oceans in winter

never gets very cold, and in the

summer the air never gets very hot.

EXAMPLE 1.10 CALCULATIONS

INVOLVING SPECIFIC HEAT

PROBLEM

How many joules are required to raise the

temperature of a 7.5-cm iron nail with a

mass of 7.05 g from room temperature


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(25C) to 100C? The specific heat capacity

of iron is 0.452 J g-1 oC-1.

SOLUTION

To solve this problem, we must multiplythe specific heat capacity by mass (g) and

the temperature change (C) to eliminate

these units and obtain joules as the units of

the answer.

specific heat capacity x mass x

temperature change = heat energy

Notice that the answer has been rounded to

two significant figures. Do you know why?

If not, review the significant figure rules for

multiplication.

Physical and chemical properties

In speaking of the properties of

substances, we also distinguish

between physical properties and

chemical properties. A physical

property can be specified without

reference to any other substance.Density, specific heat, color,

magnetism, mass, and volume are all

examples of physical properties.

A chemical property, on the other

hand, states some interaction

between chemical substances. When

iron is exposed to oxygen and water,

it corrodes and produces a new

substance called iron oxide rust.

This is a chemical property of iron.

We also say, for instance, that

sodium is very reactive toward

water. Reactivity is a chemical

property that refers to the tendency

of a substance to undergo a particular

chemical reaction.

However, to say simply that a

substance is very reactive, without

specifying "with what" or underwhat conditions, is not particularly

helpful. Sodium, for example, is very

reactive with water but quite

unreactive toward the gas helium.

ELEMENTS, COMPOUNDS, AND

MIXTURES

In nature, there are very few material

that even approach being pure compoundsnearly everything is a mixture.

The three words that form the title to this

section lie very close to the heart of

chemistry, because we work with elements,

compounds, and mixtures in the laboratory.

We must therefore understand what they are

and how to distinguish among them.

Elements are the simplest forms of

matter that can exist underconditions that we find in a chemical

laboratory; they thus are the simplest

forms of matter with which the

chemist deals directly. Elements

serve as the building blocks for all of

the more complex substances that we

encounter, from common table salt to

extremely complex proteins.

All are composed of a limited set ofelements. At present, there are 108

known elements, but only a much

smaller number will be of real

interest to us.


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Elements combine with each other to

form compounds. A compound is

characterized by having its

constituent elements always present

in the same proportions. For

example, you probably know that

water is composed of two elements:

hydrogen and oxygen. All samples of

pure water contain these two

elements combined in the proportion

of one part by mass hydrogen to

eight parts by mass oxygen (for

example, 1.0 g of hydrogen to 8.0 g

of oxygen). Also, when hydrogen is

allowed to react with oxygen toproduce water, the relative quantities

of hydrogen and oxygen that

combine are always the same.

Thus, whenever 1.0 g of hydrogen

reacts, it is always observed that only

8.0 g of oxygen are consumed, even

if more than that quantity of oxygen

is available.

Mixtures differ from elements and

compounds in that they may be of

variable composition. (As a result,

they are not considered to be pure

substances.) A solution of sodium

chloride (table salt) in water is a

mixture of two substances. We know

that by dissolving varying quantities

of salt in water (or a bowl of soup),

we can obtain solutions with a widerange of compositions. Most materi-

als found in nature or prepared in the

laboratory are not pure but instead

are mixtures.

Mixtures can be described as being

either homogeneous or

heterogeneous. A homogeneous

mixture is called a solution and has

uniform properties throughout. If we

were to sample any portion of a

sodium chloride solution, we would

find that it has the same properties

(e.g., composition) as any other

portion of the solution; we say that it

consists of a single phase. Thus, we

define a phase as any part of a

system that has uniform properties

and composition.

A heterogeneous mixture, such as oil

and water, is not uniform (Figure

1.11). If we were to sample one

portion of the mixture, it would have

the properties of water, while some

other part of the mixture would have

the properties of oil. This mixture

consists of two phases: the oil and

the water.

If we shook the mixture so that the

oil was dispersed throughout the

water as small droplets (as in a salad

dressing), all the oil droplets taken

together would still constitute only a

single phase, since the oil in one

droplet has the same properties as the

oil in another. If we added an icecube to this "brew," we would then

have three phases: the ice (a solid),

the water (a liquid), and the oil

(another liquid). In all these

examples, we can detect the presence


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of two or more phases because a

boundary exists between them.

A useful feature of pure substances is

that they undergo phase changes

(e.g., solid to liquid or liquid to gas)at constant temperature. Ice, for

instance, melts at a temperature of

0C, a temperature that remains

constant while the water undergoes

the change from solid to liquid.

When mixtures undergo phase

changes, they generally do so over a

range of temperatures. This

phenomenon provides us with oneexperimental test to determine when

we have obtained a pure substance.

There is another way that mixtures

differ from compounds and

elements. When a mixture is

prepared, the chemical properties

(and often, the physical properties)

of the components do not change,

but when elements are combined to

form a compound, very profound

changes occur in both chemical and

physical properties. For example,

copper and sulfur are two elements.

Copper, of course, is a reddish-

colored metal, a good conductor of

electricity, and is relatively resistant

to corrosion. Sulfur is a yellow

nonmetallic substance

A mixture of sulfur and copper is

easily prepared, but in the mixture

we can still see traces of the

properties of copper and the

properties of sulfur. The formation of

the mixture has involved a physical

processa process that has not

altered the chemical characteristics

of the components.

If the mixture of copper (Cu) and

sulfur (S) is heated, a chemicalreaction, or chemical change, takes

place. The copper and sulfur

combine to form a compound and

this is accompanied by dramatic

changes in the properties of the

substances. After the reaction is

over, we can't find anything that has

the properties copper, and we can't

find anything that has the properties

of sulfur. We find a new substance,

called copper(II) sulfide, that has

new properties.

Figure 1.12

(a)Alongside a crucible and its

cover we see a coil of red-colored

copper wire and yellow powdered

sulfur, (b) When mixed in the

crucible, the copper and sulfurretain their individual properties, (c)

When the mixture of copper and

sulfur is heated, a reaction takes

place and a new substance called

copper(ll) sulfide is formed. The

copper(ll) sulfide has the same

shape as the coiled copper wire from

which it was formed. Notice that its

properties differ from both the

copper and sulfur.

It doesn't conduct electricity; it

doesn't have the color of either

copper or sulfur; it has a density that

is different from both sulfur and

copper; and its chemical properties


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are entirely different, too. Such

changes are what characterize

chemical reactions.

Separating elements, compounds, and

mixtures

Since elements form compounds by

chemical reactions, decomposing

compounds into their elements also involves

chemical reactions, and they also are

accompanied by drastic changes in

properties.

Mixtures can be separated by physical

processes in which the chemical propertiesof the components are not altered. Even so,

separating mixtures is not always an easy

job, and a variety of methods have been

devised by chemists for this purpose.

For example, a mixture of salt in water can

be left to evaporate, and the departure of the

water leaves the salt behind as a solid

If we wished to recover the water as

well, we could boil the mixture in an

apparatus similar to the one in Figure

1.14 and collect the water after it has

condensed from the steam. This

process is called distillation. It is one

method that is used to obtain

drinking water from sea water (the

desalination of sea water).

Another method of separating

mixtures, called chromatography,

makes use of the different tendencies

that substances have for being

adsorbed onto the surface of certain

solids. For example, in thin-layer

chromatography (Figure 1.15), a

small spot of a solution containing

several components is placed onto

one end of a glass plate that is thinly

coated with a material such as silica

gel.

A suitable solvent is then allowed tocreep up the coating from a

reservoir. As the solvent flows past

the spot, the different components

tend to be lifted from the silica gel

surface with different degrees of

ease.

This causes the components of the

mixture to move through the silica

gel at different rates, with the morestrongly adsorbed components

moving more slowly.

The result is a separation of the

original spot into a set of spots, each

containing (we hope) one compo-

nent. This technique is widely used

today by chemists who synthesize

new compounds.

Ilustrasion of separating process

Plogiston theory

The early history of chemistry was marked,

not surprisingly, by incorrect theory about

what occurred during chemical reactions. It

had long been observed for example, that

when the combustion of wood took place,

was very light and fluffy. Metals alsochanged their appearance when heated in

air. The resulting material was less dense

than the original metal and thus appeared

lighter.

These observations led to the conclusion

something which the German chemists


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philosophers Leucippus and Democritus

suggested, as early as 400 to 500 b.c., that

matter cannot be forever divided into

smaller and smaller parts and that ultimately

particles would be encountered that would

be indivisible. These early proposals,

however, were not based on the results of

experiments and were little more than

exercises in thought.

Dalton's theory was different because it was

based on the laws of conservation of mass

and definite proportions, laws that had been

derived from many direct observations.

The theory Dalton proposed can beexpressed by the following

postulates:

1. Matter is composed of indivisible

particles called atoms.

2. All atoms of a given element have the

same properties (e.g., size, shape, and mass),

which differ from the properties of all other

elements.

3. A chemical reaction merely consists of a

reshuffling of atoms from one set of

combinations to another. The individual

atoms themselves, however, remain intact.

SOME PROPERTIES OF THE

ELEMENTS

The range of properties shown by the

elements is tremendous. At roomtemperature, some are gases, two are

liquids, and the rest are solids.

Some are metallic and some are not,

and some have properties in

between. Some are hard and some

are soft; some are very dense and

others have very low densities. With

such variety, we have to look for

ways to classify these properties so

that it is possible to make some sense

out of them.

One of the simplest methods of

classification is to divide the

elements into three categories:

metals, nonmetals, and metalloids.

The elements in each of these

categories have certain distinctive

characteristics.

Metals

Everyone has seen metals of one

kind or anotheran iron nail,

aluminum foil, copper wire, or a

"chrome-plated" bumper on a car, for

example. And you are no doubt

familiar with some of the properties

that characterize metals, even if you

haven't thought about them very

much. One of these, for instance, is

the distinctive appearance that metals

have.

They shine with a luster that is so

characteristic that it's called a

metallic luster.

Metals are also similar in their

abilities to deform without breaking

when hit with a hammer and tostretch when pulled.

All metals have both these properties

to some degree. The ability to

deform when hammered is called

malleability, and some metals, such


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as gold, can be hammered or

squeezed into extremely thin sheets.

Gold leaf(Figure 3.1), for example,

consists of gold with a small amount

of silver and copper that has beenbeaten into sheets that are so thin

(about 90 nm) that they are

translucent; some light can actually

be seen passing through them.

Malleability is also a property that a

blacksmith relies on when forging a

horse-shoe, and a silversmith uses

the malleability of silver in

hammering a design into a fine silver

tray.

The ability of a metal to stretch when

pulled from opposite directions is

called ductility. This property is used

in the manufacture of wire, which is

illustrated in Figure 3.2. The metal to

be made into wire, which might be

steel, copper, or brass, is first formed

into a rod. One end is tapered, fed

through a

die, and attached to a pulling device

on the other side. The metal is then

drawn through the die where it

undergoes a reduction in size and an

increase in length.

Everyone knows that metals are

good conductors of electricity. They

are also good conductors of heat.

If you have ever touched a metal

object that has been lying in the sun

for a while, you know how very hot

it feels. In fact, it feels much hotter

than other objects alongside that are

not metallic. The reason is that as

your hand absorbs heat from the

metal, heat travels quickly from the

neighboring parts of the object to

replace it.

Nonmetallic objects don't feel as hotbecause when your hand removes

heat, it can't be replaced rapidly, and

the part of the object in contact with

your hand becomes cooler.

More than 70% of the elements are

metals, and although there are some

similarities among them, there are

many differences, too.

Some metals are quite common and

we encounter them nearly every day

in their elemental forms, that is,

uncombined with other elements.

The metals mentioned previously

(iron, aluminum, copper, and

chromium) are only a few examples.

There are other metals that are so

reactive only chemists, or chemistrystudents, ever have an opportunity to

see them. For example, in Chapter 1

we mentioned the reaction between

sodium and chlorine to form sodium

chloride, common table salt.

Sodium is a very reactive metal that

combines not only with chlorine, but

avidly with both oxygen and

moisture. In Figure 3.3, we see a barof sodium, heavily tarnished on the

outside, but the element's bright

metallic luster is revealed by the

freshly cut slice of the metal.


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The range of chemical reactivity of

metals is very broad. Sodium is

typical of one extreme, while gold is

typical of the other. Jewelry is made

from gold for several reasons, one of

which is the fact that it doesn't

tarnish when exposed to air and

moisture.

This very low reactivity, combined

with gold's high electrical

conductivity, accounts for one of this

metal's most important commercial

uses: the plating of electrical

contacts in computers and other

electronic devices.

Besides chemical reactivity, metals

differ in certain physical properties

such as hardness and melting point.

Some metals are very hard, and some

are very soft. Chromium and iron are

examples of hard metals; gold and

lead are examples of soft ones.

Sodium is also a soft metal; in Figure

3.3, we see it being cut with a knife.

The extremes of melting point are

even more impressive.

Tungsten has the highest melting

point of any element, 3400C, which

accounts for its use as the filament in

electric light bulbs.

Mercury has the lowest melting point

of any metal, -38.9C, which means

that it is a liquid at room

temperature. As you know, mercury

is the fluid commonly used in

thermometers.

Nonmetals

Most of the nonmetallic elements

(the nonmetals) are rarely

encountered in our daily activities in

their elemental forms; instead, they

are usually found in compounds.

One nonmetal that most people have

seen is carbon, which occurs in

nature in two different forms (Figure

3.4). The more common variety is

called graphite. This is the form that

we find in charcoal briquets and the

lead in lead pencils. The less

common and more valuable form of

carbon is diamond. Graphite and

diamond have properties that are

quite different from those that we

associate with metals. Neither has

the luster of a metal, and neither is

malleable or ductile.

Other nonmetals that you have

encountered are oxygen and

nitrogen, which are the principal

components of the atmosphere.

Usually, you are not aware of their

existence, because they are colorless

gases and you can't see them.

Oxygen and nitrogen occur as

diatomic molecules, molecules

containing two atoms each. Other

nonmetallic elements form similar

molecules, and most are gases, also.

These are hydrogen (H2), fluorine(F2), and chlorine (Cl2). Bromine

(Br2) and iodine (I2) are also

diatomic, but bromine is a liquid and

iodine is a solid at room temperature.

Just as the properties of the metals

cover a broad range, so do the


17/18

properties of the nonmetals. As

we've seen, some are gases and one

(bromine) is a liquid. There are

others that are solids; carbon is just

one example. Besides differing in

these physical properties, nonmetals

differ from each other in their

chemical properties.

Fluorine, for example, is extremely

reactive while helium is inert(totally

unreactive). These differences,

which are very important, will be

explored in more detail later in the

book.

Metalloids

Metalloids are elements that have

properties that are intermediate

between those of metals and those of

nonmetals.

The best known example is the

element silicon. Two others are

arsenic (As) and antimony (Sb),which are shown in Figure 3.5. In

terms of outward appearances, these

elements have something of a

metallic look about them, but their

dark color gives them away. They

certainly differ in appearance from

typical metals such as iron or silver.

Metalloids are typically

semiconductorsthey conduct

electricity, but not nearly as well as

metals. These semiconductor

properties are especially valuable in

the electronics industry, because they

make possible all the microelectronic

devices found in hand-held

calculators and microcomputers.

Except for their electrical properties,

however, the metalloids are much

more like nonmetals than metals.

3.2 THE FIRST PERIODIC

TABLE

Chemical and physical properties

like those described in the last

section were discovered early on in

the history of chemistry.

Scientists, even as early as 1800, had

accumulated significant quantities of

information about the elements


18/18

known to them. This knowledge,

however, existed for the most part as

isolated and unrelated facts that

needed to be correlated in some

fashion before their total significance

could be grasped.

Early attempts at a classification of

the elements met with only limited

success, and it wasn't until 1869 that

the forerunner of our modern

periodic table was devised. This

resulted from the work of two

chemists, a Russian named Dmitri

Mendeleev and a German named

Julius Lothar Meyer. Both worked

independently and produced similar

tables at about the same time.

Mendeleev presented the results of

his work to the Russian Chemical

Society in the early part of 1869, but

Meyer's table didn't appear until

December of that same year.

Because Mendeleev had the good

fortune of publishing first, he is

usually given credit for the periodic

table.

Mendeleev was a chemistry teacher,

and while he was preparing a

textbook for his students, he

discovered that if he arranged the

elements in order of increasing

atomic weight, elements with similar

properties occurred at periodic

intervals. For example, he could pick

out the elements lithium (Li), sodium

(Na), potassium (K), and rubidium

(Rb). Each of these elements forms a

water-soluble compound with

chlorine that has the general formula,

MCI, where M stands for Li, Na, K,

and so on. Although this is an

interesting fact by itself, what is

especially significant is that if we

examine the elements that

immediately follow Li, Na, K, and

Rb in the list, they form another

group of similar elements.

Thus, Be follows Li, Mg follows Na,

Ca follows K, and Sr follows Rb.

These elements form the compounds

BeCl2, MgCl2, CaCl2, and SrCl2.

Recognizing this, Mendeleev dividedthe list into a series of rows and

stacked them so that those elements

having similar properties are

arranged in vertical columns. The

result was the first periodic table,

which is illustrated in Figure 3.6

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Basic Chemistry I (Print)

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