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Chapter 2 Atom

1. Atom and History of Atom

2. Subatomic Particles

3. Isotopes

4. Ions

5. Atomic Terminology

Table of Contents

Chapter 2

Warm up

The History of The Atom

• Make a list of inferences about any properties of objects in the box.

• How could you learn more about the objects in the box without opening the box?

• Scientist face these same questions as they try to learn more about atoms.

Chapter 2 The History of The Atom

1. Atomic Models

2. Modern Atomic Model

3. Quantum Numbers

4. Electron Configuration

Table of Contents

Chapter 2

Warm up

The History of The Atom




Chapter 2 The History of The Atom

• The idea of an atomic theory is more than 2000 years old.

• Until recently, scientists had never seen evidence of atoms. • Democritus (460-400BC), a Greek philosopher, first thought

that the universe was made up of very small particle and

named them “atomos”, meaning indivisible.• Aristotle, however, thought that matter was continuous, that

there was no limit on how finely you could cut it up.

• In the twentieth century, the scientists studied about atom are

Dalton, Rutherford, Planck, Einstein, Bohr and Schrödinger.

Chapter 2 1. Atomic Models

1. Dalton’s Atomic ModelIn 1803 John Dalton proposed some ideas about atom;

1. Atom or atom groups have all the characteristic properties

of substances.

2. Atoms of an element are completely identical.

3. An atom is a filled sphere like a billiard ball.

4. Different types of atoms have different masses.

5. Atoms are the smallest unit of substances and cannot be

further divided.

6. Atoms form molecules in a definite numerical ratio.

Chapter 22. Thomson’s Atomic Model

In 1897 J. J Thomson discovered the electron.

1. Protons and electrons are charged particles.

2. In neutral atoms since the number of electron and number

of protons are equal, net charge is zero.

3. An atom has a shape of sphere with 10-8 cm radius. Protons

and electron are distributed in arbitrary positions.

4. The mass of an electron is so small that it can be neglected.

1. Atomic Models

Chapter 22. Thomson’s Atomic Model

1. Atomic Models

Chapter 23. Rutherford’s Atomic Model

In 1911 Ernest Rutherford discovered the nucleus and in 1919

proton . He proposed α-particle experiment. His ideas;

1. There is a small positively charged, dense region called

nucleus in an atom.

2. The mass of the atom approximately equal to the mass of

the protons and electrons.

3. Protons are in the nucleus and electrons are distributed

around nucleus.

1. Atomic Models

Chapter 23. Rutherford’s Atomic Model

1. Atomic Models

Chapter 24. Bohr’s Atomic ModelIn 1913 Niels Bohr proposed places of electrons around the

nucleus of atoms. He concluded his ideas as follows;

1. Electrons are found in certain places so called “energy

levels” or “shell” around nucleus with certain energies.

2. Electrons move in each stationary state in circular paths.

3. When an electron falls back to a lower energy level from

a higher one, it emits a quantum of light that is equal to the

energy difference between these two energy levels.

1. Atomic Models

Chapter 24. Bohr’s Atomic Model

4. The possible stationary energy levels of electrons are

named either by letters, K, L, M, N, O ... or by positive integer

numbers starting from the lowest energy level. These

numbers are generally denoted by “n” where n = 1, 2, 3, ...

Discovery of NeutronIn 1930s James Chadwick thought that since protons have

same charges in the nucleus and they found together, there

must have been some neutral particles which hold protons

together, then he called them “neutron”.

1. Atomic Models

Chapter 2

• Although Bohr’s model was valid for atoms of H or He+1 or

Li+2, it did not work for atoms having more than one electron.

But his ideas lead to a step forward in the development of

modern atomic theory.


• The pioneers to modern atomic theory are Lois de Broglie,

Heisenberg and Schrödinger.

• In 1924 Lois de Broglie proposed that small particles

sometimes show wave-like properties.

• In 1920 Werner Heisenberg stated his uncertainty principle

which explains position of electrons.

Chapter 2 2. Modern Atomic Model

Chapter 2

• Modern atomic theory explains the probability of finding

electrons around the nucleus by virtue of quantum numbers

and orbitals.


• The quantum numbers are the integer numbers designating the energy levels of the electrons in an atom, and the orbitalsare the probable regions in which the electrons might be found around the nucleus.


Comparison of Bohr’s Atomic Model and Modern Atomic Model

Chapter 2 1. Atom

• What is all matter made up? And to what do they look like?

Warm up




Chapter 2 2.1. Atom

• Scientist have accepted that the smallest parts of

substances are called atoms.

• Atom has basically two parts, nucleusand electrons. Nucleus is located in the center of atom and electrons are rotating around nucleus with high speed.

• Atom means indivisible derived from

atomos in Greek language because of

its very small size.

Chapter 2 2. Subatomic Particles

• Scientist believed that atoms were indivisible up to 20th

century. Today it is well known that atoms have

subatomic particles, called protons, neutrons and

electrons.

• Protons are positively charged particles found in the

nucleus of an atom, and denoted by “p”. Each element

has certain number of protons which differ the element

from others.


• Neutrons are neutral particles found in the nucleus of

an atom, and denoted by “n”.

• Electrons are negatively charged particles placed

around the nucleus of an atom, and shown by “e-”.

• Protons and neutrons almost have the same masses,

but electrons have negligible mass with respect to

protons and neutrons.


• Neutral atoms have the equal number of protons and

electrons.

• Electrons are rotating in certain places called orbit,energy level or shell. Energy levels are represented by letters, K, L, M, N, O…etc, or numbers, 1, 2, 3, 4, …etc.

• Each shell can hold a certain number of electrons calculated by the equation of “2n2” where n refers to number of shell.

• The electrons located in the outermost shell of atoms are

called valance electrons.


• In the 1st shell, No. of e- = 2x12 = 2e-

In the 2nd shell, No. of e- = 2x22 = 8e-

In the 3rd shell, No. of e- = 2x32 = 18e-

In the 4th shell, No. of e- = 2x42 = 32e-

Example 1Show the electron configuration of 6C and 13Al atoms.

6C: 2) 4)

13Al: 2) 8) 3)

Solution

Chapter 2 3. Isotopes• Isotope atoms have the same number of protons but

different number of neutrons.

• They have similar chemical properties but different

physical properties.

Example 2

C6

12C

6

13C

6

14atoms are isotopes.

All they have 6 protons but 6, 7 and 8 neutrons respectively.

Chapter 2 3. Isotopes

Chapter 2

• Isotones, atoms with the same number of neutrons, but

different numbers of protons.

Example 3

P31

15S

32

16atoms are isotones.

Each have 16 neutrons but 15 and 16 protons respectively.

• Isotone atoms are completely different atoms, they have

different chemical and physical characteristics.

3. Isotopes

Chapter 2 4. Ions

• Electrically charged atoms are called ions.

• When an atom loses electrons it becomes positively

charged ion, called cation.

• When an atom gains electrons it becomes negatively

charged ion, called anion.

• Charge of an atom, q , can be found with q = p - e.

Chapter 2 4. Ions

Chapter 2 4. Ions

Atom

Neurtal Atom Ion(Charged Atom)

Cation (+) Anion (-)(p>e) (p<e)

p=e

Chapter 2 4. Ions

Example 4Find the charge and ion type of atom.

Atom Proton Electron Charge IonFe 26 24Al 13 10O 8 10P 15 15Cl 17 18

Chapter 2 4. Ions

Solution

Atom Proton Electron Charge IonFe 26 24 +2 CationAl 13 10 +3 CationO 8 10 -2 AnionP 15 15 0 NeutralCl 17 18 -1 Anion

Chapter 2

Example 5

Li+1, Ca+2, Al+3, Pb+4 are cations.

F-1, O-2, P-3 are anions.

OH- NO3-

CO3-2

PO4-3NH4

+1are polyatomic ions

4. Ions

Chapter 2 5. The Atomic Terminology

• Atomic number = Number of Protons

• Each type of atom has different number of protons.

For a neutral atom,

• Atomic number = Number of protons = Number of electrons

Z = p = e

1. Atomic Number, Z


• Atomic mass number = Number of Protons + Number of Neutrons

A = p + n

2. Atomic Mass Number, A

Example 6Fill in the blanks in the table below.

Atom p n Z ATi 22 48Al 14 27S 16 16Br 45 35


Solution

Atom p n Z ATi 22 26 22 48Al 13 14 13 27S 16 16 16 31Br 35 45 35 80


XAtomic Mass Number, A

Atomic Number, ZProtons, p

Neutrons, n

Electron, e

Charge, q

q = p - e

A = p + n


Example 7

What is the number of protons and atomic mass

number of Zn.

Zn+2

28

?

?35


Solution

Zn+2

28

65

3035

q = p - e

A = p + n

+2 = p – 28

p = 30

A = 30 + 35

A = 65


Example 8

Cr+3 ion has 21 electrons and its atomic mass number is 52.

What is the number of neutrons for Cr ?

Solution

q = p - e

A = p + n3 = p – 21 p = 24

52 = 24 + n n = 28


• 12C is accepted as a standard atom. The mass of the 12C

isotope atom is accepted as 12 amu ( atomic mass unit) and

atomic masses of other atoms were calculated accordingly.

For example relative atomic masses of H atom is 1.008

(≈1 amu) and oxygen atom becomes 15.9994 (≈16 amu)

with respect to 12C.

3.Relative Atomic Mass and Relative Formula Mass


• Relative formula mass is sum of the relative atomic masses of

atoms found in a compound.

Example 9Relative formula mass of CH4 and NO2

CH4 = (1x12) + (4x1) = 16 amu

NO2 = (1x14) + (2x16) = 46 amu


• Most of the elements in nature are found as a mixture of

isotope atoms.

• The average atomic mass is the average masses of natural

isotopes of an element.

4.Average Atomic Mass

Average Atomic Mass =

A1x% of 1st isotope + A2x% of 2nd isotope + ……

A1 and A2 are atomic mass numbers of natural isotopes.


Example 10

Naturally occurring Ga consists of 60% 69Ga and 40% 71Ga.

What is the average mass of Ga?

SolutionAverage Atomic Mass =

A1x% of 1st isotope + A2x% of 2nd isotope + ……

Average Atomic Mass = 69x60 + 71x40100 100

= 69.80 amu


Example 11

Find the average atomic mass of Pb?

End of the chapter 2

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