Handouts Chem

8/6/2019 Handouts Chem

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Chemistry: Matter and its Classifications

You may have done some or all of the

following tasks: eaten some meals, drunk coffee orcola, taken a shower with soap, gone to the library

to research a paper, taken notes in a class, checkedemail on a computer, watched some television,

ridden a bike or car to a part-time job, taken an

aspirin to relieve a headache, and spent some of the

evening having snacks and refreshments with

friends. Perhaps, without your awareness, your life

was touched by chemistry in each of theseactivities. What, then, is this discipline we call

chemistry? What is a matter?

Chemistry is the study of matterits

composition, properties, and transformations.

Matteris anything that has mass and takes up

volume.

In other words, chemistry studies anything that

we touch, feel, see, smell, or taste, from simple

substances like water or salt, to complex

substances like proteins and carbohydrates that

combine to form the human body.

Some matter is naturally occurringand some aresynthetically made.

*****************************************

PROBLEM 1.1

Look around you and identify five objects.

Decide if they are composed of natural or

synthetic materials.*****************************************

Matter exists in three common statessolid,

liquid, and gas.

A solid has a definite volume, and maintains its

shape regardless of the container in which it is

placed. The particles of a solid lie close together,and are arranged in a regular three-dimensional

array.

A liquid has a definite volume, but takes on the

shape of the container it occupies. The particles of

a liquid are close together, but they can randomly

move around, sliding past one another.

A gas has no definite shape or volume. The

particles of a gas move randomly and are separated by a distance much larger than their size. The

particles of a gas expand to fill the volume and

assume the shape of whatever container they are

put in.

All matter can be classified as either a puresubstanceor a mixture. These can either undergo

physical change or chemical change.

Matter can undergo chemical change orphysical

change.

A chemical change, (1) one or more substances are

used up (at least partially), (2) one or more new

substances are formed, and (3) energy is absorbed

or released. As substances undergo chemical

changes, they demonstrate their chemicalproperties.

Aphysical change, on the other hand, occurs with

no change in chemical composition. Physical

properties are usually altered significantly as

matter undergoes physical changes

*****************************************

PROBLEM 1.2

Which of the following are pure

substances? Which of the following are

homogeneous mixtures? Explain your

answers.

(a) sugar dissolved in water;

(b) tea and ice;(c) french onion soup;

(d) mud;(e) gasoline;

(f ) carbon dioxide;

(g) a chocolate-chip cookie.

*****************************************


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Measurements

2.1 UNIT MEASUREMENTS

Units of Measurement

The metric system of weights and measures is used by scientists of all fields, including chemists. This System

uses the base 10 for measurements; for conversions, measurements may be multiplied or divided by 10.The measures of length, volume, mass, energy, and temperature are used to evaluate our physical and chemical

environment. The laboratory equipment associated with obtaining these measures is also listed.

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PROBLEM 2.1.1

What term is used for each of the following units:

(a) a million liters;

(b) a thousandth of a second;

(c) a hundredth of a gram;(d) a tenth of a liter?

**********************************************************************


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2.2 Conversion of Units

In the scientific world, the most applicable system of measurement is the Metric System. However, another

system measurement exists, the English System. So in the case that an English Unit is given, we need to

convert it into Metric Unit. The table below shows the relationship of the two system measurement.

SAMPLE PROBLEM 2.2.1

Express 1.47 miles in inches.

2.3 SCIENTIFIC NOTATION

Numbers uses lots and lots of application especially in the daily lives. During a laboratory or research,

there is s probability of using large numbers or very small value. Such numbers can be simplified and expressed

intoScientific notation.

In using such large and small numbers, it is inconvenient to write down all the zeroes. In scientific (exponential)

notation, we place one nonzero digit to the left of the decimal.



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2.4 SIGNIFICANT FIGURES

There are two kinds of numbers. Numbers

obtained by counting or from definitions are exact

numbers. They are known to be absolutelyaccurate. For example, the exact number of people

in a closed room can be counted, and there is nodoubt about the number of people. A dozen eggs is

defined as exactly 12 eggs, no more, no fewer.

The second kind is Numbers obtained from

measurements are not exact. Every measurement

involves an estimate. For example, suppose you areasked to measure the length of this page to the

nearest 0.1 mm.

Significant figures are digits believed to be correct

by the person who makes a measurement.

Here are the following rules of finding significant

figures.

1.Nonzero digits are always significant.

For example, 38.57 mL has four

significant figures; 288 g has three significant

figures.

2. Zeroes are sometimes significant, and sometimes

they are not.

a. Zeroes at the beginning of a number

(used just to position the decimal point)

are never significant.

For example, 0.052 g has two significant

figures; 0.00364 m has three significant

figures. These could also be reported in

scientific notation as 5.2 x 10-2 g and3.64 x 10-3 m, respectively.

b. Zeroes between nonzero digits are

always significant.

For example, 2007 g has four significant

figures; 6.08 km has three significant

figures.c. Zeroes at the end of a number that

contains a decimal point are alwayssignificant.

For example, 38.0 cm has three significant

figures; 440.0 m has four significant

figures. These could also be reported as

3.80 x 101 cm and 4.400 x 102 m,respectively.

d. Zeroes at the endof a number that does

not contain a decimal point may or

may not be significant.

For example, the quantity 24,300 kmcould represent three, four, or five

significant figures. We are given

insufficient information to answer the

question. If both of the zeroes are used

just to place the decimal point, the number

should appear as 2.43 x 104 km

(three significant figures).

If only one of the zeroes is used to place

the decimal point (i.e., the number wasmeasured +/-10), the number is 2.430 x

104 km (four significant

figures).

If the number is actually known to be

24,300 +/- 1, it should be written as 2.4300 x 10 4

km (five significant figures).

Operations of Significant Figures

1. Addition and Subtractionthe number that hasthe highest significant figures should be the same

significant figures in the result of the operation

performed.

12.1 + 13.45 = 25.55 No. of ( 3 ) ( 4 ) ( 4 )

significant figure

2. Multiplication and Division the number that

has the least number of significant figures should

be the same in the result of the operation

performed.

12.1 x 13.45 = 283.795 284

No. of ( 3 ) ( 4 ) ( 5 ) (3)significantfigure

***************************************************************************

PROBLEM 2.4.1

Perform the operation and report the correct significant figure.

a. 24.003 x 5.28 + 3.921 b. 37.10 2.23 x 0.3156

c. 791.5789 - 32.11225 x 500.11 + 361.12

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Other Measurements: Temperature, Density and Specific gravity

3.1 TEMPERATURE

Temperature is a measure of how hot or cold an object is. Three temperature scales are used: Fahrenheit

(most common in the United States), Celsius (most commonly used by scientists and countries other than theUnited States), and Kelvin.

The Fahrenheit and Celsius scales are both divided into degrees. On the Fahrenheit scale, water freezes at 32 F

and boils at 212 F. On the Celsius scale, water freezes at 0 C and boils at 100 C.

To convert temperature values from one scale to another, we use two equations, where Cis the Celsius

temperature and Fis the Fahrenheit temperature.

To convert from Celsius to Fahrenheit: To convert from Fahrenheit to Celsius:

F = 1.8(C) + 32 C =F 32

1.8

The Kelvin scale is divided into kelvins (K), not degrees. The only difference between the Kelvin scale and the

Celsius scale is the zero point. A temperature of 273 C corresponds to 0 K. The zero point on the Kelvin scale

is called absolute zero, the lowest temperature possible. To convert temperature values from Celsius to Kelvin,

or vice versa, use two equations.

To convert from Celsius to Kelvin: To convert from Kelvin to Celsius:

K= C + 273 C = K 273


An infant had a temperature of 104 F. Convert this temperature to both C and K.

First convert the Fahrenheit temperature to degrees Celsius using the equation

C = (F 32)/1.8.

Then convert the Celsius temperature to kelvins by adding 273.

3.2 DENSITY

Density is a physical property that relates the mass of a substance to its volume. Density is reported in

grams per milliliter (g/mL) or grams per cubic centimeter (g/cc).

density = . mass (g) .

volume (mL or cc)


Calculate the mass in grams of 15.0 mL of a saline solution that has a density 1.05 g/mL.

Use density (g/mL) to interconvert the mass and volume of a liquid.


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***********************************************************************

PROBLEM 3.2.1

Calculate the mass in grams of 10.0 mL of diethyl ether, an anesthetic that has a density of 0.713 g/mL.

***********************************************************************

3.3 SPECIFIC GRAVITY

Specific gravity is a quantity that compares the density of a substance with the density of water at the same

temperature.

***********************************************************************PROBLEM 3.2.1

(a) If the density of a liquid is 0.80 g/mL, what is its specific gravity?

(b) If the specific gravity of a substance is 2.3, what is its density?

***********************************************************************

Elements

4.1 ELEMENT

An elementis a pure substance that cannot be broken down into simpler substances by a chemical reaction.

Of the 114 elements currently known, 90 are naturally occurring and the remaining 24 have been

prepared by scientists in the laboratory. Some elements, like oxygen in the air we breathe and aluminum in asoft drink can, are familiar to you, while others, like samarium and seaborgium, are probably not. An

alphabetical list of all elements appears on the inside front cover.

Each element is identified by a one- or two-letter symbol. The element carbon is symbolized by the

single letter C, while the element chlorine is symbolized by Cl. When two letters are used in the element

symbol, the first is upper case while the second is lower case. Thus, Co refers to the element cobalt, but CO is

carbon monoxide, which is composed of the elements carbon (C) and oxygen (O). While most element symbols

are derived from the first one or two letters of the element name, 11 elements have symbols derived from the

Latin names for them.

4.2 ELEMENTS IN THE PERIODIC TABLE

The elements in the periodic table are divided into three groupsmetals, nonmetals, and metalloids.


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The solid line that begins with boron (B) and angles in steps down to astatine (At) marks the three regions

corresponding to these groups. All metals are located to the leftof the line. All nonmetals except hydrogen are

located to the right. Metalloids are located along the steps.

Metals are shiny solids that are good conductors of heat and electricity. All metals are solids at room

temperature except for mercury, which is a liquid.

Nonmetals do not have a shiny appearance, and they are generally poor Conductors of heat and

electricity. Nonmetals like sulfur and carbon are solids at room temperature; bromine is a liquid; andnitrogen, oxygen, and nine other elements are gases.

Metalloids have properties intermediate between metals and nonmetals. Only seven elements are

categorized as metalloids: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb),

tellurium (Te), and astatine (At).

4.3 ELEMENTS IN THE HUMAN BODY

Because living organisms selectively take up elements from their surroundings, the abundance of

elements in the human body is very different from the distribution of elements in the earths crust.

Four nonmetalsoxygen, carbon, hydrogen, and nitrogencomprise 96% of the mass of the

human body, and are called the building-block elements (Figure 2.2). Hydrogen and oxygen are the elements

that form water, the most prevalent substance in the body. Carbon, hydrogen, and oxygen are found in the fourmain types of biological moleculesproteins, carbohydrates, lipids, and nucleic acids. Proteins and nucleic

acids contain the element nitrogen as well. These are biological molecules.

Seven other elements, called the major minerals ormacronutrients, are also present in the body in

much smaller amounts (0.12% by mass). Sodium, potassium, and chlorine are present in body fluids.

Magnesium and sulfur occur in proteins, and calcium and phosphorus are present in teeth and bones.

Phosphorus is also contained in all nucleic acids, such as the DNA that transfers genetic information from one

generation to another. At least 100 mg of each macronutrient is needed in the daily diet.

Many other elements occur in very small amounts in the body, but are essential to good health. Thesetrace elements ormicronutrients are required in the daily diet in small quantities, usually less than 15 mg.

Each trace element has a specialized function that is important for proper cellular function. For example, iron is

needed for hemoglobin, the protein that carries oxygen in red blood cells, and myoglobin, the protein that stores

oxygen in muscle. Zinc is needed for the proper functioning of many enzymes in the liver and kidneys, and

iodine is needed for proper thyroid function. Although most of the trace elements are metals, nonmetals like

fluorine and selenium are micronutrients as well.


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Atoms and its Structure

All matter is composed of the same basic building blocks called atoms. An atom is much too small

to be seen even by the most powerful light microscopes. The period at the end of this sentence holds about 1

108 atoms, and a human cheek cell contains about 1 10 16 atoms. An atom is composed of three subatomic

particles. A proton, symbolized byp, has a positive (+) charge.

An electron, symbolized by e, has a negative () charge.

A neutron, symbolized by n, has no charge.

Protons and neutrons have approximately the same, exceedingly small mass. The mass of an electron is much

less, 1/1,836 the mass of a proton. These subatomic particles are not evenly distributed in the volume of an

atom. There are two main components of an atom.

The nucleus is a dense core that contains the protons and neutrons. Most of the mass of an atom resides

in the nucleus.

The electron cloudis composed of electrons that move rapidly in the almost Empty space surrounding

the nucleus. The electron cloud comprises most of the volume of an atom.

Every atom of a given type of element always has thesame number of protons in the nucleus, a value

called theatomic number, symbolized byZ. Conversely, two differentelements have differentatomic numbers.

Thus, the element hydrogen has one proton in its nucleus, so its atomic number is one. Lithium has

three protons in its nucleus, so its atomic number is three. The periodic table is arranged in order of increasing

atomic number beginning at the upper left-hand corner. The atomic number appears just above the element

symbol for each entry in the table.

**********************************************************************PROBLEM 5.1

Identify the element with each atomic number, and give the number of protons and electrons in the neutralatom:

(a) 2;

(b) 11;

(c) 20;

(d) 47;

(e) 78.

**********************************************************************

Both protons and neutrons contribute to the mass of an atom. The mass number, symbolized byA, is the sum

of the number of protons and neutrons.

Mass number (A) = the number of protons (Z) + the number of neutrons.

***************************************************************************PROBLEM 5.2

How many protons, neutrons, and electrons are contained in each atom with the given atomicnumber and mass number?

a.Z= 17,A = 35

b.Z= 14,A = 28

c.Z= 92,A = 238

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Isotopes and Atomic Weight

6.1 ISOTOPES

Isotopes are atoms of the same element having a

different number of neutrons. Shown below, an

example of the isotopes of the element chlorine.

The mass number differ since the number of the

neutrons present for each isotope is different.


In the given isotopes shown above, find

the number of neutrons for each isotope

Analysis: Since the mass number is equal

to the sum of number of protons and neutrons, we

have

mass number = no. of n + no. of proton,Z

Rearranging the equation,

No. ofn = mass number no. of proton ,Z

Solution:

For the isotope 3517Cl :

Z= 17 , mass number= 35

No. of n = 35 17 = 18 neutrons

For the isotope 3717Cl :Z= 17 , mass number= 37

No. of n = 37 17 = 20 neutrons

6.2 ATOMIC WEIGHT

Some elements like fluorine occur

naturally as a single isotope. More commonly, an

element is a mixture of isotopes, and it is useful to

know the average mass, called the atomic weight

(oratomic mass), of the atoms in a sample.

Theatomic weight is the weighted average of the

mass of the naturally occurring isotopes of aparticular element reported in atomic mass units.

As shown below, element carbon presented the

atomic weight 12.01. This value is the average

mass of the naturally occurring isotopes of carbon.

SAMPLE PROBLEM 6.2.1:

What is the atomic weight of the element

chlorine?

Mass Isotopic Abundance

Cl-35 34.97 75.78% = 0.7578

Cl-37 36.97 24.22% = 0.2422

Solution:

Step 1: Multiply the isotopic abundance in nature

to its mass.

As for Cl-35,

34.97 amu x 0.7578 = 26.5003 amu

As for Cl-3736.97 amu x 0.2422 = 8.9541 amu

Step 2: Add the resulting mass of the two isotopes.

As for Cl-35,

34.97 amu x 0.7578 = 26.5003 amu

As for Cl-3736.97 amu x 0.2422 = 8.9541 amu

____________________

Atomic Weight, Cl =35.4544 amu 35.45 amu

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PROBLEM 6.2.1

Calculate the atomic weight of each element given the mass and natural occurrence of each

isotope.

a. Magnesium Mass (amu) Isotopic Abundance

Mg-24 23.99 78.99%

Mg-25 24.99 10.00%

Mg-26 25.98 11.01%

b. Vanadium Mass (amu) Abundance

V-50 49.95 0.250%

V-51 50.94 99.750%


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LESSON 7- The Periodic Table

The periodic table evolved over many years, and it resulted from the careful observations and experiments of

many brilliant scientists in the nineteenth century. Most prominent was Russian chemist Dmitri Mendeleev,whose arrangement in 1869 of the 60 known elements into groups having similar properties in order of

increasing atomic number became the precursor of the modern periodic table.

5.1 Basic Features of the Periodic Table

The periodic table is arranged into seven horizontal rows and 18 vertical columns. The particular rowand column tell us much about the properties of an element.

A row in the periodic table is called a period. Elements in the same row are similar in size.

A column in the periodic table is called a group. Elements in the same group have similarelectronic

and chemical properties.

The rows in the periodic table are numbered 17. The number of elements in each row varies. The first

period has just two elements, hydrogen and helium. The second and third rows have eight elements each, and

the fourth and fifth rows have 18 elements. Also note that two groups of fourteen elements appear at the bottomof the periodic table. The lanthanides,beginning with the element cerium (Z= 58), immediately follow the

element lanthanum (La). The actinides, beginning with thorium (Z = 90), immediately follow the element

actinium (Ac).

Each column in the periodic table is assigned a group number. Groups are numbered in two ways. In

one system, the 18 columns of the periodic table are assigned the numbers 118, beginning with the column

farthest to the left. An older but still widely used system numbers the groups 18, followed by the letter A or B.

The main group elements consist of the two columns on the far left and the six columns on the far right

of the table. These groups are numbered 1A8A.

The transition metal elements are contained in the 10 short columns in the middle of the table,

numbered 1B8B.

The inner transition elements consist of the lanthanides and actinides, and they are not assigned group

numbers.

5.2 Groups in the Periodic Table

The elements in the Periodic Table are classified by the following groups:

Alkali the elements in Group-I starting from the element lithium (Li) down to the element

francium (Fr). It has the following characteristics:

They are soft and shiny and have low melting points.

They are good conductors of heat and electricity.

They react readily with water to form basic solutions.

The Alkali does not exist in nature as pure element, but instead existed as compounds combining withother kind of element. An example of a compound from group 1A elements includes sodium chloride

(NaCl), table salt;

Alkaline Earth elements, located in group 2A (group 2), include beryllium (Be), down to

radium (Ra). Similar to alkali group, it also exists as compound such as barium sulfate (BaSO4), which

is used to improve the quality of X-ray images of the gastrointestinal tract.

Halogens, located in group 7A (group 17), include fl uorine (F), chlorine (Cl), bromine (Br),

iodine(I), and the rare radioactive element astatine (At). In their elemental form, halogens contain two

atoms joined togetherF2, Cl2, Br2, and I2. Fluorine and chlorine are gases at room temperature,

bromine is a liquid, and iodine is a solid. Halogens are very reactive and combine with many other

elements to form compounds.


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Noble gases, located in group 8A (group 18), include helium (He), neon (Ne), argon (Ar),

krypton (Kr), xenon (Xe), and radon (Rn). Unlike other elements, the noble gases are especially stable

as atoms, and so they rarely combine with other elements to form compounds. The noble gas radon

has received attention in recent years. Radon is a radioactive gas, and generally its concentration in the

air is low and therefore its presence harmless. High radon levels are linked to an increased risk of lungcancer.

Electronic Configurations of Elements

The chemistry of an element is determined by the number of electrons in an atom. These electrons occupy space

in the atomic structure. Thus, interaction to another atom is based on the orientation of electrons in the region of

space of an atom. The electronic structure is described by the following:

Electrons do not move freely in space; rather, an electron is confined to a specific region, giving it a

particular energy.

Electrons occupy discrete energy levels. The energy of electrons is quantized; that is, the energy is

restricted to specific values.

The electrons that surround a nucleus are confined to regions called the principal energy levels, orshells.

The shells are numbered, n = 1, 2, 3, 4, and so forth, beginning closest to the nucleus.

Electrons closer to the nucleus are held more tightly and are lower in energy.

Electrons farther from the nucleus are held less tightly and are higher in energy.

The number of electrons that can occupy a given shell is determined by the value ofn. The farther a shell is

from the nucleus, the larger its volume becomes, and the more electrons it

can hold.

To determine these configurations, we use the Aufbau Principle, Paulli exclusion Principleand Hunds Rule

as a guide:

Rule [1] (Aufbau Principle)

Electrons are placed in the lowest energy orbitals beginning with the 1s orbital.

In comparing similar types of orbitals from one shell to another (e.g., 2s and3s), an orbital closer to the

nucleus is lower in energy. Thus, the energy of a 2s orbital is lower than a 3s orbital. Within a shell,orbital energies increase in the following order:s, p, d, f.


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Rule [2] (Paulli Exclusion Principle)

Each orbital holds a maximum of two electrons.

Rule [3] (Hunds Rule)When orbitals are equal in energy, one electron is added to each orbital until the

orbitals are half-filled, before any orbital is completely filled. For example, one electron is added toeach of the threep orbitals before filling any p orbital with two electrons. Because like charges repel

each other, adding electrons to different p orbitals keeps them farther away from each other, which is

energetically favorable.

LESSON 10- Trends in Periodic Table

Many properties of atoms exhibit periodic trends; that is, they change in a regular way across a row or down acolumn of the periodic table. Two properties that illustrate this phenomenon are atomic size and ionization

energy.

10.1 ATOMIC SIZE

The size of atoms increases down a column of the periodic table, as the valence electrons are farther from the

nucleus.


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The size of atoms decreases across a row of the periodic table as the number of protons in the nucleus increases.

An increasing number of protons pulls the electrons closer to the nucleus, so the atom gets smaller.

**********************************************************************PROBLEM 10.1.1

Rank the atoms in each group in order of increasing size.

a. boron, carbon, neon d. krypton, neon, xenon

b. calcium, magnesium, beryllium e. sulfur, oxygen, silicon

c. silicon, sulfur, magnesium f. fl uorine, sulfur, aluminum

**********************************************************************

10.2 IONIZATION ENERGY

Since a negatively charged electron is attracted to a positively charged nucleus, energy is required to remove an

electron from a neutral atom. The more tightly the electron is held, the greater the energy required to remove it.Removing an electron from a neutral atom forms a cation.

The ionization energy is the energy needed to remove an electron from a neutral atom.

A cation is positively charged, and has fewer electrons than the neutral atom.

Two periodic trends characterize ionization energy.:

Ionization energies decrease down a column of the periodic table as the valence electrons get farther

from the positively charged nucleus.

Ionization energies generally increase across a row of the periodic table as the number of protons inthe nucleus increases.

***************************************************************************PROBLEM 10.2.1

Arrange the elements in each group in order of increasing ionization energy.

a. phosphorus, silicon, sulfurb. magnesium, calcium, beryllium

c. carbon, fluorine, beryllium

d. neon, krypton, argon

e. tin, silicon, sulfur


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f. calcium, aluminum, nitrogen

***************************************************************************LESSON 11- Chemical Bonding: Ionic Bonding

Matter mostly exists as compounds. These compounds are formed due to their chemical nature. This

nature is the ability to attract one atom or another. This interaction is called as Chemical Bonding.Bondingis

the joining of two atoms in a stable arrangement.

Not all elements can readily form a bond. These elements are called as Noble gases. Their electronic

configuration is especially stable such that it doesnt readily bond to form a compound. In bonding, elements

gain, lose, or share electrons to attain the electronic configuration of the noble gas closest to them in the

periodic table. Bonding is classified into two major types, the ionic bonding and covalent bonding.

IONIC BONDING

Ionic bondsresult from the transfer of electrons from one element to another. Ionic bonds form between a metal

on the left side of the periodic table and a nonmetal on the right side. For example, when the metal sodium (Na)

bonds to the nonmetal chlorine (Cl2), the ionic compound sodium chloride (NaCl) forms. Since ionic

compounds are composed of ionscharged species in which the number of protons and electrons in anatom is notequal.

Ionic compounds consist of oppositely charged ions that have a strong electrostatic attraction for

each other. There are two types of ions called cations and anions.Cations are positively charged ions. A cation

has fewer electrons than protons.Anions are negatively charged ions. An anion has more electrons than protons.

Each of these ions formed from a main group element has the s and threep orbitals filled with eight electrons.

This results in the octet rule.

Example 1:

Sodium (group 1A) has an atomic number of 11, giving it 11 protons and 11 electrons in the neutral atom.

Electronic configuration of Na : 1s22s22p63s1

This gives sodium one more electron than neon, the noble gas closest to it in the periodic table.

Electronic configuration of Ne : 1s22s22p6

In losing one electron, sodium forms a cation with a +1 charge, which still has 11 protons, but now has only 10

electrons in its electron cloud. Electronic configuration of Na : 1s22s22p6 = Na+(Sodium Cation)

in losing electron

Example 2:

For chlorine (Cl), it has 17 electrons and 17 protons.

Electronic configuration of Cl: 1s22s22p63s23p5= Cl (neutral atom) =

It must gain one electron to attain same configuration of the noble gas Argon.

Electronic configuration of Ar: 1s2

2s2

2p6

3s2

3p6

Electronic configuration of Cl: 1s22s22p63s23p6 = Cl- (Chloride anion)=

(After gaining electron)


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IONIC COMPOUNDS

Ionic compounds are composed of cations and anions. The ions in ionic compounds are arranged to maximize

the attractive forces between molecules.

Example:A sodium ion and chloride ion can form a compound sodium chloride.

Na+ + Cl- NaCl

LESSON 12- Compounds: Naming Ionic Compounds

In naming Ionic Compounds, one must be familiar with the common cations and anions including the

polyatomic anions. As shown in the table below, the following are elements and its ion form and names. Use the

names as described by the table.

Table 12.1 Common Cations Table 12.2 Common Anions

Table 12.3 Common name of Polyatomic anions


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EXAMPLE 12.1What is the name of the ionic compound formed from the ions of Mg2+ and OH-?

LESSON 13- Compounds: Formation through Covalent Bonding

Covalent Bonds result from the sharing of electrons between two atoms. Recall the valence electrons, to

achieve sharing of electrons, valence electron of an atom should meet the electronic configuration of the closest

noble gas as to octet rule or duet rule in the case of hydrogen atom to reach helium electronic configuration.


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We use a solid line between two element symbols to represent a two-electron bond.

In covalent bonding, atoms share electrons to attain the electronic configuration of the noble gas closest to them

in the periodic table.

Covalent bonds are formed when two nonmetals combine.Nonmetals do not easily lose electrons, and as a

result, one nonmetal does not readily transfer an electron to another nonmetal.

The F atom shares two electrons in one covalent bond, and it also contains three pairs of electrons that it doesnot share with hydrogen. These unshared electron pairs are called nonbonded electron pairs orlone pairs.

The presentation above of each atom or the formed compound is what we called the Lewis-Dot Structure. It

shows the bonding and the lone pairs to meet the electronic configuration of the Noble gases as to octet rule or

duet rule.

LESSON 14- Compounds: Naming Covalent Compounds

Although some covalent compounds are always referred to by their common namesH 2O (water) and NH3(ammonia)these names tell us nothing about the atoms that the molecule contains. Other covalent compounds

with two elements are named to indicate the identity and

number of elements they contain.

STEPS IN NAMING COVALENT COMPOUNDS

For the compound NO2 ,

STEP 1: Name the first nonmetal by its element name and the second using the suffix

ide.

Nitrogen + Oxide(First Nonmetal element) (Second Nonmetal element)

STEP 2: Add prefixes to show the number of atoms of each element. (Prefixes are shown

below.)

NO2 contains one N atom, so the prefix mono- is understood. Since NO2 contains two O atoms, use the

prefix di- dioxide.

Thus, NO2 is nitrogen dioxide.


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LESSON 15- Chemical Reactions and Balancing Chemical

Equations

When ice melted into liquid water, there is no chemical change happened but only physical change. The

changes are only to its state from solid to liquid. In chemical change, there is chemical reaction involved, thus

changing the nature of the matter.

Chemical reactions involve breaking bonds in the starting materials, called reactants, and

forming new bonds in theproducts.

We have different types of Chemical reaction:

Single Displacement Reaction - AB + C A + CBDouble Displacement Reaction - AB + CD AD + CBDecomposition - AB A + BOxidation/ Combustion (With Oxygen)- AB + O2 AO2 + B

NOTE : A, B, C and D are not elements, this is only representation of all elements.

What is shown above in the types of chemical reaction are the Chemical Equations. Chemical equations are

used to describe chemical reactions, and they show

(1) the substances that react, called reactants;

(2) the substances formed, called products; and(3) therelative amounts of the substances involved.

We write the reactants to the leftof an arrow and the products to the rightof the arrow.

EXAMPLE 15.1:

The reactants are CH4 and O2, and the products are CO2 and H2O.

You may happen to observe that there is numerical value 2 on the left side of O2. This is what we called thecoefficientwhich gives 2 moles O2. In the case of CH4, There is no numerical value but this is understood as

1mole CH4. This is also similar to the products 1 mole CO2. Thus, there is also 2moles H2O.

BALANCING CHEMICAL EQUATIONS


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In any chemical reaction, there is no detectable change in the quantity of matter during the course of reaction.

This guiding principle, the Law of Conservation of Matter, provides the basis for balancing chemical

equations and for calculations based on those equations.

An equation must be balanced by adding coefficients in front of some formulas so that the number of atoms of

each element is equal on both sides of the equation.

EXAMPLE 15.2:

Write a balanced chemical equation for the reaction of propane (C3H8) with oxygen (O2) to form carbon dioxide

(CO2) and water (H2O).

STEP 1. Write the equation with the correct formulas.

STEP 2. Balance the equation with coefficients one element at a time.

For C,

For H,

For O,

STEP 3. Check to make sure that the smallest set of whole numbers is used.

LESSON 16- The Mole Conversion

A mole defines a quantity, much like a dozen items means 12, and a case of soda means 24 cans. The only

difference is that a mole is much larger.

Amole is a quantity that contains 6.02 1023 itemsusually atoms, molecules, or ions.

The definition of a mole is based on the number of atoms contained in exactly 12 g of the carbon-12 isotope.This number is called Avogadros number, after the Italian scientist Amadeo Avogadro, who first proposed the

concept of a mole in the nineteenth century. One mole, abbreviated as mol, always contains an Avogadros

number of particles which is 6.021023.

So to say,

1 mole of C atoms = 6.02 1023 C atoms

1 mole of CO2 molecules = 6.02 1023 CO2 molecules

1 mole of H2O molecules = 6.02 1023 H2O molecules

1 mole of vitamin C molecules 6.02 1023 vitamin C molecules

SAMPLE PROBLEM 15.1: (Moles to Avogadros number)


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How many molecules are contained in 5.0 moles of carbon dioxide (CO2)?

Step 1. Identify the original quantity and the desired quantity.

Step 2. Write out the conversion factors.

Step 3. Set up and solve the problem.

When reactions are carried out in the laboratory, single atoms and molecules are much too small to measure out.

Instead, substances are weighed on a balance and amounts are typically reported in grams, not atomic mass

units. To determine how many atoms or molecules are contained in a given mass, we use its molar mass.

The molar massis the mass of one mole of any substance, reported in grams per mole. The value of the molar

mass of an element in the periodic table (in grams per mole) is the same as the value of its atomic weight (in

amu).

SAMPLE PROBLEM 15.2:

What is the molar mass of nicotine (C10H14N2), the toxic and addictive stimulant in tobacco?

ANALYSIS Determine the number of atoms of each element from the subscripts in

the chemical formula,multiply the number of atoms of each element by the atomic weight,

and add up the results.

SOLUTION10 C atoms 12.01 amu = 120.1 amu

14 H atoms 1.01 amu = 14.14 amu

2 N atoms 14.01 amu = 28.02 amu

Formula weight of nicotine = 162.26 amu rounded to 162.3 amu

Answer: Since the formula weight of nicotine is 162.3 amu, the molar mass of

nicotine is162.3 g/mol.

The molar mass is a very useful quantity because it relates the number ofmoles to the number ofgrams of asubstance. In this way, the molar mass can be used as a conversion factor. For example, since the molar mass

of H2O is 18.0 g/mol, two conversion factors can be written.

Using these conversion factors, we can convert a given number of moles of water to grams, or a specifi c

number of grams of water to moles.

SAMPLE PROBLEM 15.3: (Moles to Mass)

How many grams does 0.25 moles of water weigh?

Step 1. Identify the original quantity and the desired quantity.

Step 2.Write out the conversion factors

Step 3.Set up and solve the problem.


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LESSON 17- Stoichiometry

In every chemical reaction, The Law of conservation of Mass is always observed. From this idea, thedetermination of amount reactants or products is possible through the chemical equation. The process is called

stoichiometry.

The stoichiometry process requires the following, molecular weight/ atomic weight, moles, mass, andstoichiometric ratio .

SAMPLE PROBLEM 17.1

Phosphorus, P4, burns with excess oxygen to form tetraphosphorus decoxide, P4O10. In thisreaction, what mass of P4 reacts with 1.50 moles of O2?

PlanThe balanced equation tells us that one mole of P4 reacts with five moles of O2.

The stoichiometric ratios of the chemical equation shown above are as follows:

1mol P4 , 1mol P4 , 5mol O2

5 mol O2 1mol P4O10 1mol P4O10

We will use the ratio 1mol P4

5 mol O2

LESSON 18- Gases and its Laws

18.1 KINETIC MOLECULAR THEORY AND PRESSURE

Simple gases in the atmosphereoxygen (O2), carbon dioxide (CO2), and ozone (O3)are vital to life. Oxygen,which constitutes 21% of the earths atmosphere, is needed for metabolic processes that convert carbohydrates

to energy. Green plants use carbon dioxide, a minor component of the atmosphere, to store the energy of the sun

in the bonds of carbohydrate molecules during photosynthesis. Ozone forms a protective shield in the upper

atmosphere to filter out harmful radiation from the sun, thus keeping it from the surface of the earth.

The properties of all gases, can be explained by the kinetic-molecular theory of gases, a set of principles based

on the following assumptions:

A gas consists of particlesatoms or moleculesthat move randomly and rapidly.

The size of gas particles is small compared to the space between the particles.

Because the space between gas particles is large, gas particles exert no attractive forces on each other.

The kinetic energy of gas particles increases with increasing temperature.

When gas particles collide with each other, they rebound and travel in new directions.

When gas particles collide with the walls of a container, they exert a pressure.


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Because gas particles move rapidly, two gases mix together quickly. Moreover, when a gas is added to a

container, the particles rapidly move to fill the entire container.

When many gas molecules strike a surface, they exert a measurable pressure. Pressure (P) is the force (F)

exerted per unit area (A).

Shown below are the units of Pressure.

18.2 GAS LAWS

Gases behave depending on its condition. The volume, temperature, pressure and number of particles are

involved for its behavior. We have different laws that provide description of behavior of gases.

Boyles law: For a fixed amount of gas at constant temperature, the pressure and volume of a

gas are inversely related.

The general formula is,

SAMPLE PROBLEM 18.1

If a 4.0-L container of helium gas has a pressure of 10.0 atm, what pressure does the gas exert if the

volume is increased to 6.0 L?

Given:

Solution:

Charless law

All gases expand when they are heated and contract when they are cooled. Charless law describes how the

volume of a gas changes as the Kelvin temperature is changed.For a fixed amount of gas at constant pressure,the volume of a gas is proportional to its Kelvin temperature.

Volume and temperature areproportional; that is, as one quantity increases, the otherincreases as well.

Thus, dividing volume by temperature is a constant (k).


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Charles Law Equation

SAMPLE PROBLEM 18.2

A balloon that contains 0.50 L of air at 25 C is cooled to 196 C. What volume does the balloon now

occupy?

Given:

Solution:

(continue solving using Charles Law)

GayLussacs law:

For a fixed amount of gas at constant volume, the pressure of a gas is proportional to its Kelvin

temperature. GayLussacs law describes how the pressure of a gas changes as the Kelvin temperature is

changed. GayLussacs law describes how the pressure of a gas changes as the Kelvin temperature is changed.

Pressure and temperature areproportional; that is, as one quantity increases, the otherincreases. Thus, dividing

the pressure by the temperature is a constant (k).

Combined Gas lawAll three gas lawsBoyles, Charless, and GayLussacs lawscan be combined in a single

equation, the combined gas law, that relates pressure, volume, and temperature.

Avogadros law

It describes the relationship between the number of moles of a gas and its volume. When the pressure

and temperature are held constant, the volume of a gas is proportional to the number of moles present. As the

number of moles of a gas increases, its volume increases as well.


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Ideal Gas Law

An ideal gas is one that exactly obeys these gas laws. Many real gases show slight deviations from

ideality, but at normal temperatures and pressures the deviations are usually small enough to be ignored.

The product of pressure and volume divided by the product of moles and Kelvin temperature is a constant,

called the universal gas constant denoted by R.

Daltons Law of Partial Pressure

Since gas particles are very far apart compared to the size of an individual particle, gas particles

behave independently. As a result, the identity of the components of a gas mixture does not matter, and a

mixture of gases behaves like a pure gas. Each component of a gas mixture is said to exert a pressure called

its partial pressure. Daltons law describes the relationship between the partial pressures of the components

and the total pressure of a gas mixture. Daltons law: The total pressure (Ptotal) of a gas mixture is the sum of

the partial pressures of its component gases.

LESSON 18- Solutions

Most matter with which we come into contact, however, is a mixture composed of two or more pure

substances. The air we breathe is composed of nitrogen and oxygen, together with small amounts of argon,

water vapor, carbon dioxide, and other gases. Seawater is composed largely of sodium chloride and water. A

mixture may be heterogeneous orhomogeneous.

A heterogeneous mixture does not have a uniform composition throughout a sample.Example: Peperoni Pizza, halo-halo, salt and pepper.

A homogeneous mixture has a uniform composition throughout a sample.

There are two types of Homogeneous mixture:

A solution is a homogeneous mixture that contains small particles. Liquid solutions are transparent.

Example: Saltwater, coffee, vinegar

A colloid is a homogeneous mixture with larger particles, often having an opaque appearance.

Example: Milk and whipped cream

In Solutions; there are two things involved,

Solute- the dissolved particle. Smaller quantity

Solvent- the dissolving medium. Larger quantity

PROPERTY OF SOLUTION

Solubility is a common property that is required for the preparation of a solution. It is the ability of the matter to

be dissolved in a medium. There several factors that may increase or decrease its solubility;


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Polarity-Polarsubstances can have the ability to dissolve polar substances. Much like salt

dissolves in water since both are polar. Non-polar substances can dissolve non-polar. An

example here is the oil, it does not dissolve in water and vice versa.

Saturation- This involves in the amount of the solute to be dissolved in the solvent. We have

three saturation points

o Unsaturated- Means that little amount of solute is dissolved in the solution, addition

of solute may have the tendency to be dissolved in the solution

o Saturated- At this point, the capacity of the solvent to dissolve a solute is reached.

Beyond this, the solute may not be dissolved anymore by the solvent.

o Supersaturated- The state may achieved if there is conditions is met such as high

temperature.

Temperature- Higher temperature may increase the solubility of the substance.

Pressure- Pressure changes do not affect the solubility of liquids and solids, but pressure

affects the solubility of gases a great deal. Henrys law describes the effect of pressure on gas

solubility. ( Henrys law: The solubility of a gas in a liquid is proportional to the partial

pressure of the gas above the liquid.)

Thus, the higher the pressure, the higher the solubility of a gas in a solvent.

LESSON 19- Solutions: Percent Concentration

Medicines are given only by a capsule and or syrup. In every dosage of the medicine, there is already a specific

amount of active substance in the mixture of a medicine. To express this amount, the term concentration is

applied. There are several concentration unit that is used.

19.1 Weight per volume Percent

One of the most common measures of concentration is weight/volume percent concentration,

(w/v)%that is, the number of grams of solute dissolved in 100 mL of solution.



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A vinegar contains 5 g of acetic acid dissolved in 100 mL of solution, so the acetic

acid concentration is 5% (w/v).

19.2 Volume per Volume Percent

Volume/volume percent concentration, (v/v)%that is, the number of milliliters of solute dissolved in 100 mL

of solution.

SAMPLE PROBLEM 19.2

A bottle of rubbing alcohol that contains 70 mL of 2-propanol in 100 mL of solution

has a 70% (v/v) concentration of 2-propanol.

19.3 Parts per Million

When a solution contains a very small concentration of solute, concentration is often expressed in parts per

million (ppm). The parts may be expressed in either mass or volume units as long as thesame unit is used forboth the numerator and denominator.

A sample of seawater that contains 1.3 g of magnesium ions in 106 g of solution contains 1.3 ppm of

magnesium.


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LESSON 20- Solutions: Molarity and Dilution

20.1 MOLARITY

The most common measure of concentration in the laboratory is molaritythe number of

moles of solute per liter of solution, abbreviated as M.

SAMPLE PROBLEM 20.1

A solution that is formed from 1.00 mol (58.4 g) of NaCl in enough water to give 1.00 L ofsolution has a molarity of 1.00 M. A solution that is formed from 2.50 mol (146 g) of NaCl in

enough water to give 2.50 L of solution is also a 1.00 M solution. Both solutions contain thesame number of moles per unit volume.

20.2 DILUTION

Sometimes a solution has a higher concentration than is needed. Dilution is the addition of solvent to

decrease the concentration of solute. For example, a stock solution of a drug is often supplied in a

concentrated form to take up less space on a pharmacy shelf, and then it is diluted so that it can be administered

in a reasonable volume and lower concentration that allows for more accurate dosing. A key fact to keep in

mind is that the amount of solute is constant. Only the volume of the solution is changed by adding solvent.

If we have,

We get the number of moles, thus, if we have initial values for the molarity and volume (M 1 and V1), we can

calculate a new value for the molarity or volume (M 2 or V2), since the product of the molarity and volumeequals the number of moles, a constant.

SAMPLE PROBLEM 20.2

What is the concentration of a solution formed by diluting 5.0 mL of a 3.2 M glucose solution to

40.0 mL?

Given:

Solution:


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LESSON 21- Solutions: Colligative Properties and Osmotic

Pressure

21.1 COLLIGATIVE PROPERTIES

Colligative properties are properties of a solution that depend on the concentration of the solute but not its

identity. We have two properties,Boiling point elevation and Freezing point depression

Boiling point elevation

A solute in a solution can be volatile ornonvolatile. A volatile solute readily escapes into

the vapor phase. A nonvolatile solute does not readily escape into the vapor phase, and thus it has a

negligible vapor pressure at a given temperature.

If there are fewer solvent molecules in the solution than there are in the pure liquid,

there are fewer molecules in the gas phase as well. As a result, the vapor pressure above the

solution is lowerthan the vapor pressure of the pure solvent.

The boiling point is the temperature at which the vapor pressure equals the atmospheric

pressure. A lower vapor pressure means that the solution must be heated to a higher temperature to get

the vapor pressure to equal the atmospheric pressure. This results in boiling point elevation.

Therefore, the amount that the boiling point increases depends only on the number of dissolved

particles.



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FREEZING POINT DEPRESSION

In a similar manner, a dissolved solute lowers the freezing point of a solvent. Thepresence of solute molecules makes it harder for solvent molecules to form an organized crystalline

solid, thus lowering the temperature at which the liquid phase becomes solid. This results in freezing point

depression.

A liquid solution that contains a nonvolatile solute has a lower freezing point than the solvent alone.

The amount of freezing point depression depends only on the number of dissolved particles.


20.3 OSMOTIC PRESSURE

Osmosis is the passage of water and small molecules across a semipermeable membrane from a

solution of low solute concentration to a solution of higher solute concentration. Osmotic pressure is the

pressure that prevents the fl ow of additional solvent into a solution on one side of a semipermeable membrane.

The greater the number of dissolved particles, the greater the osmotic pressure.

SAMPLE PROBLEM 20.3

A 0.1 M glucose solution is separated from a 0.2 M glucose solution by a semipermeablemembrane.

(a) Which solution exerts the greater osmotic pressure?

(b) In which direction will water flow between the two solutions?

(c) Describe the level of the two solutions when equilibrium is reached.

Solution:

a. The greater the number of dissolved particles, the higher the osmotic pressure, so the

0.2M glucose solution exerts the greater pressure.b. Water will fl ow from the less concentrated solution (0.1 M) to the more

concentrated solution (0.2 M).

c. Since water fl ows into the 0.2 M solution, its height will increase, and the height ofthe 0.1 M glucose solution will decrease.LESSON 22- Acids and Bases: Nature and its Measure

ACIDS and BASES

When a covalent acid dissolves in water, proton transfer forms H3O+ and an anion. This process is

called dissociation. Acids differ in their tendency to donate a proton; that is, acids differ in the extent to which


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they dissociate in water. A strong acid readily donates a proton. When a strong acid dissolves in water,

essentially 100% of the acid dissociates into ions. A weak acid less readily donates a proton. When a weak acid

dissolves in water, only a small fraction of the acid dissociates into ions. Common strong acids include HI, HBr,

HCl, H2SO4, and HNO3. It has common properties of being corrosive, sour smell and taste, and low pH.

Bases also differ in their ability to accept a proton. A strong base readily accepts a proton. When a

strong base dissolves in water, essentially 100% of the base dissociates into ions. A weak base less readilyaccepts a proton. When a weak base dissolves in water, only a small fraction of the base forms ions. The most

common strong base is hydroxide, OH, used as a variety of metal salts, including NaOH and KOH.

Commonly, bases areslippery, bittertaste and high pH.

A strong acid readily donates a proton, forming a weakconjugate base.

A strong base readily accepts a proton, forming a weakconjugate acid.

Table 22.1 Acids Strength and its Conjugate base

THE pH SCALE

The pH or function of hydrogen measures the presence of H3O+ or H+ ions in a solution. This will determine if

the substance is acidic or basic by nature.

IfpHis from 1 to 6, it is acidic.

And 8-14 is basic.

The pH value 7 is considered to be neutral.

Since values for the hydronium ion concentration are very small, with negative powers of ten, the pH scale isused to more conveniently report [H3O

+]. The pH of a solution is a number generally between 0 and 14, defined

in terms of the logarithm (log) of the H3O+ concentration.

A logarithm is an exponent of a power of ten.

SAMPLE PROBLEM:

The value of [H3O+] in apple juice is about 1 104, or 104 written without the coefficient.

The pH of this solution is?


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The relationships of concentration of H3O+ are as follows:

The lowerthe pH, the higherthe concentration of H3O+.

LESSON 23- Energy and its changes

Energy is the capacity to do work. Whenever you throw a ball, ride a bike, or read a newspaper, you

use energy to do work. There are two types of energy. Potential energy is stored energy. Kinetic energy is

the energy of motion.

Energy can be converted from one form to another, one rule, the law of conservation of energy,

governs the process. It is the total energy in a system does not change. Energy cannot be created or destroyed.

The energy stored in chemical bondsboth ionic and covalentis a form of potential energy. In

chemical reactions, potential energy may be released and converted to heat, the kinetic energy of the movingparticles of the product. Reactions that form products having lowerpotential energy than the reactants are

favored.

Units of Energy

Energy can be measured using two different units, calories (cal) and joules (J). A calorie is the

amount of energy needed to raise the temperature of 1 g of water 1 C.

SAMPLE PROBLEM:

A reaction releases 421 kJ of energy. How many kilocalories does this correspond to?

In a chemical reaction, Energy is also involved. The energy is utilized either for breaking bonds orformation of bonds. The breaking of bonds requires high energy while formation of bonds is releasing ofenergy. The transferred from or to a system maybe described by Enthalpy. Enthalpy is the quantity of heat

transferred into or out of a system as it undergoes a chemical or physical change at constant pressure, qp, and is

defined by enthalpy change, H. It is presented by positive or negative value.

When energy is absorbed, the reaction is said to be endothermic and His positive (+).

When energy is released, the reaction is said to be exothermic and His negative ().

SAMPLE PROBLEM:


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The combustion of propane (C3H8) with O2 according to the given balanced chemical equation

releases 531 kcal/mol. How many kilocalories of energy are released when 0.750 mol of propane is

burned?

The reaction in the problem is an exothermic reaction. A molecule for each substance creates collision, only

collisions that has sufficient energy and proper orientation leads to reaction. This can be illustrated by the

energy diagram shown below.

(EXOTHERMIC REACTION)

The Ea is the activation energy or required energy to reach the transition state. The transition state refers to the

breaking or starting to form bonds of the molecules of the substance, thus, reaction is possible.

(ENDOTHERMIC REACTION)

LESSON 24- Reaction Rates

Reaction rates refer to the how fast do the reaction occurs. There are several factors that affect reaction rates:

Concentration -Increasing the concentration of the reactants increases the number of

collisions, so the reaction rate increases.

Temperature- Increasing the temperature increases the reaction rate.

Catalyst - A catalystis a substance that speeds up the rate of a reaction. A

catalyst is recovered unchanged in a reaction, and it does not appear in theproduct.

- Catalysts accelerate a reaction by lowering the energy of activation

Metals and Enzymes are one example of a catalyst.

- Lactase is the enzyme that binds lactose, the principal carbohydrate in

milk. Once bound, lactose is converted into two simpler sugars,

glucose and galactose. When individuals lack adequate amounts of this


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enzyme, they are unable to digest lactose, and this causes abdominal

cramping and diarrhea.

LESSON 25- Nuclear Chemistry

Radioactivity is the nuclear radiation emitted by a radioactive isotope. A radioactive isotope, called a

radioisotope, is unstable and spontaneously emits energy to form a more stable nucleus. Radioactive decay is

the process, by which an unstable radioactive nucleus emits radiation.

There are different forms of radiation are emitted when a radioactive nucleus is converted to a more stablenucleus, including alpha particles, beta particles, positrons, and gamma radiation.

An alpha particleis a high-energy particle that contains two protons and two neutrons.

A beta particleis a high-energy electron

o

Apositronis called an antiparticleof a particle, since their charges are different but their masses are

the same.

Gamma raysare high-energy radiation released from a radioactive nucleus.


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An example of radioactive substance is Cobalt-60. It is used in external radiation treatment for cancer.

Radiation generated by cobalt-60 decay is focused on a specific c site in the body that contains cancerous cells.

By directing the radiation on the tumor, damage to surrounding healthy tissues is minimized.

The nuclear reactions used in nuclear power plants occur by a process called nuclear fission, whereas the

nuclear reactions that take place in the sun occur by a process called nuclear fusion.

Nuclear fissionis the splitting apart of a heavy nucleus into lighter nuclei and neutrons.

The produced high energy neutron can bombard to another uranium molecule that causes another

reaction and so on. Thus, a Chain reaction will occur and releases massive energy.

Nuclear fusionis the joining together of two light nuclei to form a larger nucleus.

HALF-LIFE

Radionuclides have different stabilities and decay at different rates. Some decay half of its amount by nearly

completely in a fraction of a second and others only after millions of years. The isotope cobalt-60 has a half life

of 5.27 years. Suppose we have 20.0g cobalt-60, after 5.27 years, the amount of cobalt-60 left was only 10.0g.

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