Atomic Structure Lecture 1-2

Fundamentals of Inorganic Chemistry

Lectures 1 & 2

SCH 100

Dr. E. Changamu

[email protected]

Office: Biochemistry Rm 11

mailto:[email protected]

Evaluation

• 2 CAT 20 marks• Practical 10 marks (Rem. Lab coats)• End of term Exam 70

Syllabus• The early theories of atomic structure: the fundamental

particles of the atom; • Rutherford’s planetary model• The Bohr theory of the atom; failure of the Bohr theory. • Planck’s quantization of energy • The photoelectric effect; • Qualitative treatment of the atomic orbitals (s, p d and f).

• The Aufbau Principle and the periodic table. • Common oxidation states of the elements.

Syllabus• Naturally occurring and artificially made

isotopes, and their applications. • Atomic properties of the elements

– electronegativity – electron affinity– ionization energy

• Chemical bonding – ionic, – covalent, – metallic, – coordinate

Syllabus• Hybridization of atomic orbitals and shapes of simple

molecules and ions.

• The nature of ionic and covalent compounds as influenced by the above factors.

• The mole concept and its application. • General concepts of acids and bases - strong and weak acids

and bases; pH calculations. • Balancing of redox reactions.

Inorganic Chemistry

• Is the custodian of all the elements known.• Is concerned with the

– Occurrence of the elements in nature– Extraction from natural sources– The reactions of the elements with other

elements and compounds– Safety and application of the elements and/ or

their compounds

Inorganic Chemistry• Is related to the other divisions of chemistry

– Organometallic chemistry bridges inorganic and organic chemistry. It deals with compounds containing direct metal-carbon bonds.

The study of the composition, properties, and transformations of matter

Chemistry

Inorganic Chemistry

– Bioinorganic chemistry bridges biochemistry and inorganic chemistry

– Environmental chemistry includes the study of both inorganic and organic compounds.

Lecture Objectives At the end of this lecture you should be able to

– describe the historical development of atomic structure

– descried the nature of electrons, protons and neutrons

– explain the Thomson model of the atom– explain Rutherford model of the atom– describe the nature of electromagnetic radiation– discuss the Bohr model and the atomic hydrogen

spectrum

Atomic Structure

460 BC Democritus develops the idea of atoms

He pounded up materials in his pestle and

mortar until he had reduced them to smaller

and smaller particles which he called Atoma

ATOMA

(greek for indivisible)

HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY

1808 John Dalton1

Suggested that all matter was made up of tiny

spheres that were able to bounce around with

perfect elasticity and called them Atoms.

John Dalton, A New System of Chemical Philosophy, 1808; reprinted with an introduction by Alexander Joseph, Peter Owen Limited, London, 1965.


Dalton’s Atomic Theory • Elements are made of extremely small particles called atoms.• Atoms of a given element are identical in size, mass, and other

properties; atoms of different elements differ in size, mass, and other properties.

• Atoms cannot be subdivided, created, or destroyed.• Atoms of different elements combine in simple whole-number

ratios to form chemical compounds.• In chemical reactions, atoms are combined, separated, or

rearranged.


Access to materials /Notes

• http://soma.ku.ac.ke

• Log in using your user id and password provided by university.

E.g I20/xyzabc/2011

• Enrollment key = sch100

http://soma.ku.ac.ke/

1898 Joseph John Thompson

Found from his cathode ray tube experiments

that atoms could sometimes eject a far smaller

negative particle which he called an electron.

The electrons came from the atoms making up the cathode.


Joseph John Thompson

• Thomson found out that the electrons were the same regardless of

the metal he used for the cathode. Therefore he concluded they

were part of the structure of all atoms.

• He calculated the ratio of their charge to their mass

Charge/mass = 1.76 x 1011 C/kg

• Robert Millikan determined the charge of the electron

Charge of electron, e = 1.602 x 10-19 C

• Thomson calculated the mass of the electron to be

Mass of electron, me = 9.109 x 10-31 kg


Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge like plums surrounded by pudding.

PLUM PUDDING

MODEL


Ernest Rutherford, 1910

• Oversaw Geiger and Marsden carrying out his

famous experiment.

• They fired α-particles (a type of naturally-occurring

radiation consisting of positively charged helium

atoms) at a piece of gold foil which was only a few

atoms thick.

• They found that although most of them passed

through. About 1 in 10,000 hit something and

bounced right back.


gold foil

helium nuclei

They found that while most of the α-particles passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back.

helium nuclei


The Proton

• In 1918, while Rutherford was performing his various experiments in the field of radioactivity, bombarding nitrogen gas with alpha particles, he noticed that one of the experimental results was a surge of hydrogen.

• He correctly deduced that the hydrogen atoms must have come from within the nitrogen atoms themselves, which would mean that there was something within all of these atoms which was divisible, the amount of which would determine what element the atom represented.

The proton• The particle Rutherford isolated was the proton,

which by itself constitutes the nucleus of a single hydrogen atom, though in this case it was “ionized” (missing its electron, thereby giving it a net positive charge), which Rutherford determined by exposing the resulting hydrogen to magnetic fields.

• Rutherford’s new evidence

allowed him to propose a more

detailed model with a central

nucleus.

• He suggested that the positive

charge was all in a central

nucleus. With this holding the

electrons in place by electrical

attraction.


The Neutron

• By James Chadwick in 1932.• Last to be discovered due to its neutrality.• Now the structure was complete

(somewhat)• See details of discovery in

http://isaacmmcphee.suite101.com/the-discovery-of-the-neutron-a46060



Particle Charge (e) Mass (amu)Proton (p) +1 1Neutron (n) Neutral 1Electron (e-) -1 1/1836

Fundamental particles

1 atomic mass unit (amu) = 1.672621777(74)×10−27 kg

1 electric charge (e) = 1.602176565(35)×10−19 C

Nature of Electromagnetic radiation

Niels Bohr, 1913

• Studied under Rutherford at the Victoria

University in Manchester.

• Bohr refined Rutherford's idea by adding

that the electrons were in orbits. Rather like

planets orbiting the sun. With each orbit only

able to contain a set number of electrons.


Bohr’s Model

Nucleus

Electron

Orbit

Electron orbits26

HELIUM ATOM

+N

N

+-

-

proton

electron

neutron

Shell

What do these particles consist of?

ATOMIC STRUCTURE

the number of protons in an atom

the number of protons and neutrons in an atom

He2

4 Atomic mass = protons + neutrons

Atomic number = No of protons

number of electrons = number of protons

ATOMIC STRUCTURE

Electrons are arranged in Energy Levels

or Shells around the nucleus of an atom.

• first shell a maximum of 2 electrons

• second shell a maximum of 8

electrons

• third shell a maximum of 8

electrons

ATOMIC STRUCTURE

There are two ways to represent the atomic

structure of an element or compound;

1. Electronic Configuration

2. Dot & Cross Diagrams

ELECTRONIC CONFIGURATION

With electronic configuration elements are

represented numerically by the number of electrons

in their shells and number of shells. For example;

N

Nitrogen

7

14

2 in 1st shell

5 in 2nd shell

configuration = 2 , 5

2 + 5 = 7

ELECTRONIC CONFIGURATION

Write the electronic configuration for the following elements;

Ca O

Cl Si

Na20

40

11

23

8

17

16

35

14

28B

11

5

a) b) c)

d) e) f)

2,8,8,2 2,8,1

2,8,7 2,8,4 2,3

2,6

DOT & CROSS DIAGRAMS

With Dot & Cross diagrams elements and

compounds are represented by Dots or Crosses to

show electrons, and circles to show the shells. For

example;

Nitrogen N XX X

X

XX

X

N7

14

DOT & CROSS DIAGRAMS

Draw the Dot & Cross diagrams for the following elements;

O Cl8 17

16 35a) b)

O

X

XX

X

X

X

X

X

Cl

X

X

X

X X

X

XX

X

X

X

X

X

XX

X

X

X

SUMMARY

1. The Atomic Number of an atom = number of

protons in the nucleus.

2. The Atomic Mass of an atom = number of

Protons + Neutrons in the nucleus.

3. The number of Protons = Number of Electrons.

4. Electrons orbit the nucleus in shells.

5. Each shell can only carry a set number of electrons.

This powerpoint was kindly donated to www.worldofteaching.com

http://www.worldofteaching.com is home to over a thousand powerpoints submitted by teachers. This is a completely free site and requires no registration. Please visit and I hope it will help in your teaching.

http://www.worldofteaching.com/

http://www.worldofteaching.com/

Atomic Structure Lecture 1-2

Documents

atomic properties

atomic theory10

atomic orbitals s

idea of atoms

chemical compounds

covalent compounds

organic compounds

bohr model