1. Chemical Building Blocks:
Atomic Theories, Models, and
Electronic Structure

2. Outline
1. A Brief History
 Pre-atomic view of matter
 The Greek Idea
 Lavoisier: Conservation of Mass
 Proust: Definite Proportions
 Dalton: Atomic Theory
2. Electricity and the Atom
3. X-rays and Radioactivity
4. Atomic Models
 Rutherford’s nuclear model
 Structure of the nucleus
 The Bohr model and electronic arrangement
5. The Quantum Model
 Electronic configuration
 Quantum numbers

3. Pre-Atomic View of Matter
Matter was seen as
continuous, as the
four Greek elements
were.
No distinct divisions
between fire, water,
earth, air
Mixtures of the four
gave the properties of
being hot or cold,
moist, or dry

4. EARLY THOUGHTS ON MATTER
 LEUCIPPUS: “There must  DEMOCRITUS: Referred to
be tiny particles of water these particles as atomos;
that could not be Each atom was distinct in
subdivided.” Observe the size and shape (eg. Water
SAND. as round balls, Fire as
sharp)

5. EARLY THOUGHTS ON MATTER
 ARISTOTLE: Matter was
continuous, not atomistic
 Five centuries after, the
Roman LUCRETIUS wrote a
poem, “On the Nature of
Things,” where strong
arguments for the atomic
nature were presented
 Late 1600’s, ROBERT
BOYLE publishes The
Sceptical Chymist. Proper
experimentation can show
if theory is valid.
Introduction of ELEMENT
and COMPOUND

6. LAVOISIER: Conservation of Mass
 Antoine Laurent Lavoisier
 “When a chemical reaction is
carried out in a closed system, the
total mass of the system is not
changed.”
 Red mercuric oxide  Mercury +
OXYGEN
 1st to use systematic names; 1st
chem bk. ; “father”
 Experiments with burning coal
(combustion), and breathing guinea
pigs (respiration). We cannot create
 LAW: Matter is neither created nor from nothing.
destroyed in a chemical change.
The total mass of the reaction Chemistry is
products is always equal to the total about
mass of the reactants transformation.

7. Proust: Definite Proportions
 Joseph Louis Proust: Copper carbonate always had
the same composition
 LAW OF DEFINITE PROPORTIONS or CONSTANT
COMPOSITION: A compound always contains the
same elements in certain definite proportions and in
no other combinations.
 J.J.Berzelius: Prepared an extensive list of atomic
weights; Lead sulfide experiments
 Henry Cavendish: 1783;
 Hydrogen gas + Oxygen gas  Water
 1800: Volta designed a powerful battery W.Nicholson
and A.Carlisle would use to separate water into its
elements.

8. Dalton: Law of Multiple
Proportions
 Elements could
combine in in more
than one set of
proportions.
 If elements A and B
react to form two
different compounds,
the masses of B
combined with a fixed
mass of A, can be
expressed as a ratio of
small whole numbers

9. Dalton’s Atomic Theory
 Matter is composed of extremely small particles
called atoms.
p All atoms of a given element are identical,
having the same size, mass and chemical
properties. The atoms of one element are
different from the atoms of all other elements.
 Compounds are composed of atoms of different
elements combined in fixed proportions.
p Chemical reactions only involve the
rearrangement of atoms. Atoms are not created
or destroyed in chemical reactions.

10. 2

11. 16 X + 8Y 8 X2Y

12. ELECTRICITY and the ATOM
 ELECTROLYSIS
 CATHODE RAY TUBES
 William Crookes passed an electric current
through a tube with air at low pressure. The
tube has metal electrodes and the beam of
current is seen as green fluorescence when it
strikes the Zinc fluoride coated screen

13. J.J. Thomson, measured mass/charge of e-
(1906 Nobel Prize in Physics) 2.2

14. Thomson’s Experiment: M/C
 Were the rays beams of particles or did it
consist of energy, much like light?
 1897: Joseph John Thomson
 Cathode rays are deflected in an electric
field. Attracted by POSITIVE, deflected by
NEGATIVE.
 Particles were the same regardless of the
electrodes or gas.
 What would he conclude, then?

15. Foundations of Atomic Theory
Law of Conservation of Mass
Mass is neither destroyed nor created during ordinary chemical
reactions.
Law of Definite Proportions
The fact that a chemical compound contains the same elements
in exactly the same proportions by mass regardless of the size
of the sample or source of the compound.
Law of Multiple Proportions
If two or more different compounds are composed of the
same two elements, then the ratio of the masses of the
second element combined with a certain mass of the first
elements is always a ratio of small whole numbers.

16. Law of Definite Proportions
Whether synthesized in the laboratory or obtained from
various natural sources, copper carbonate always has
the same composition.
Analysis of this compound led Proust to formulate
the law of definite proportions.
+ +
103 g of 53 g of 40 g of oxygen 10 g of carbon
copper carbonate copper

17. Law of Multiple Proportions
John Dalton (1766 – 1844)
If two elements form more than one
compound, the ratio of the second
element that combines with 1 gram of the
first element in each is a simple whole
number.
e.g. H2O & H2O2
water hydrogen peroxide
Ratio of oxygen is 1:2 (an exact ratio)

18. Cathode Ray Tube
2.2

19. Plum-Pudding Model
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56

20. Thomson’s Calculations Thomson
PAPER
Cathode Ray Experiment
 Thomson used magnetic and electric fields to measure and calculate
the ratio of the cathode ray’s mass to its charge.
charge of electric length of length of
Electric ray particle x field x deflection region x drift region
=
deflection mass of ray velocity of 2
x
particle ray particle
charge of magnetic length of length of
Magnetic ray particle x field x deflection region x drift region
=
deflection mass of ray velocity of
x
particle ray particle
magnetic deflection magnetic field
= x velocity
electric deflection electric field

21. Measured mass of e-
(1923 Nobel Prize in Physics)
e- charge = -1.60 x 10-19 C
Thomson’s charge/mass of e- = -1.76 x 108 C/g
e- mass = 9.10 x 10-28 g

22. ROENTGEN RAYS
 WILHELM CONRAD ROENTGEN
 1895
 Working in a darkroom on the glow
produced in certain substances by cathode
rays
 He noticed the glow also occurring on a
chemically-treated piece of paper (even in
next room). The “ray” could travel through
walls!
 When he waved his hand between the
radiation source and the glowing paper, he
could see the bones of his own hand on the
paper.
X-RAYS

23. Discovery of Radioactivity
 ANTOINE HENRI BECQUEREL
 1895
 Studied fluorescence by
1. Wrap photographic film in black paper
2. Place fluorescing crystals on top
3. Place paper in strong sunlight
i. Like ordinary light: would not pass through
paper
ii. Like X-rays: would pass and fog the film
4. Work with URANIUM: always fogging the film!
5. Perhaps, radiation coming from uranium was
unrelated to fluorescence but a characteristic of
uranium.
 MARIE SKLODOWSKA: coined the term
“Radioactivity”: the spontaneous emission of
radiation from certain unstable elements

24. The CURIES
 MARIE SKLODOWSKA
married PIERRE
CURIE, a French
Physicist and discover
radioactive polonium
and radium
 1903 Nobel in Physics
(Becquerel, Curie,
Curie)
 Marie Curie: 2nd Nobel
prize in 1911

25. TYPES OF RADIOACTIVITY
 Three types of radiation emanated
from these radioactive elements,
 ERNEST RUTHERFORD, a New
Zealander, chose the names
alpha, beta, and gamma for
these.
 ALPHA: beams of positive particles
(identical to He2+ ions
 BETA: negatively-charged; identical to
cathode rays; therefore, an electron
 GAMMA: not deflected; penetrating; a
form of energy

26. (Uranium compound)

27. Health Effects of Ionizing Radiation

28. REM – roentgen equivalent man
Degree of exposure How much time?
5-25:genetic damage. genetic damage in 18-90s
50:alter white blood cells.
75-125: radiation alteration of white blood
sickness. cells in 3 minutes
400: kill 50% of exposed
people radiation sickness in 5-8
500-600 :will kill almost minutes
all exposed people.
death in 35-60 minutes.