The document discusses the history and development of atomic theory, including key contributors like Democritus, Dalton, Thomson, Rutherford, and Bohr. It describes the discovery of subatomic particles like electrons, protons, and neutrons through experiments. The structure of the atom is explained, with protons and neutrons located in the tiny nucleus at the center and electrons orbiting in empty space around the nucleus. Isotopes and how atomic mass is calculated are also covered.

1. Atomic Structure & the Periodic Table

2. Objectives
 Identify important developments in the history of atomic theory.
 Summarize Dalton’s atomic theory.
 Describe the size of an atom.
 Distinguish among protons, electrons, and neutrons in terms of
relative mass and change.
 Describe the structure of an atom, including the location of the
protons, electrons, and neutrons with respect to the nucleus.
 Explain how the atomic number identifies an element.
 Use the atomic number and mass number of an element to find the
number of protons, electrons, and neutrons.
 Explain how isotopes differ and why the atomic masses of elements
are not whole numbers.
 Calculate the average atomic mass of an element from isotope data.
TEKS:
2A, 2B, 2C, 2D, 2E, 3A, 3C, 3E, 4A, 4C,4D, 5A, 6A, 6B, 6C, 8A, 9B,
10A, 11A

3. Early Models of the Atom
400 B.C. – Democritus proposed the existence of
fundamental particles of matter that were indivisible and
indestructible - “atomos”.
Aristotle thought all matter was continuous; he did not
believe in atoms.
Neither idea was supported by any experimental
evidence – speculation only.

4. Foundations of Atomic Theory
 The late 1700’s –definitions and basic laws had
been discovered and accepted by chemists.
 Element – substance that cannot be broken down by
ordinary chemical means.
 Chemical Reaction – transformation of substance or
substances into one or more new substances.

5.  Law of Conservation of Mass – mass cannot be created or
destroyed just changed from one form to another. (Antoine
Lavosier)
 Law of Definite Proportions – a chemical compound contains
exactly the same elements in the same proportion regardless of
sample size. (Joseph Proust from work of Gay-Lussac &
Amadeo Avogadro – 1802/1804)
 Law of Multiple Proportions – If two or more different
compounds are composed of the same two elements, then the
ratio of the masses of those elements will always exist as a ratio
of small whole numbers. (John Dalton - 1808)

6. Dalton’s Atomic Theory
 All elements are composed of tiny indivisible particles
called atoms.
 Atoms of the same element are identical. The atoms of
one element are different from the atoms of another
element.
 Atoms combine in simple whole-number ratios.
 Atoms are separated, joined or rearranged in chemical
reactions. Atoms of one element are never changed into
atoms of another element as a result of a chemical
reaction.

7. Discovery of Electrons
 1897 – J.J. Thomson – “Cathode Ray Tube Experiment”
 Showed existence of first know sub-atomic particle
 Determined charge to mass ratio of the electron
 1909 – Robert Millikan found the charge of the electron –
“Millikan’s Oil Drop Experiment”

8. Cathode Ray Tube
High Voltage
Gas at very low
pressure
Metal disk
(anode)
Metal disk Cathode Ray
(cathode) (electrons)

9. Cathode Ray Tube
High Voltage
Gas at very low
Negative plate
pressure
Metal disk
(anode)
Metal disk Positive plate Cathode Ray
(cathode) (electrons)

10. Rutherford’s Gold Foil Experiment
 Rutherford, Geiger & Marsden (1912) -showed that
most of the atom was empty space, but that atoms
had a solid, positive core.
Alpha Particles
Lead
shield
Radioactive
source

11. Discovery of Protons
 1919 -J.J. Thomson & James Chadwick–
discovered particles traveling opposite of the
cathode rays.
 Determined existence, mass and charge of protons
 Idea had actually been previously proposed by
Goldstein in 1886.

12. Cathode Ray Tube
High Voltage
Gas at very
low pressure
Negative plate
protons
Metal disk
(anode)
Metal disk
Positive plate Cathode Ray
(cathode) (electrons)

13. Neutrons
 James Chadwick 1932 - confirmed the existence of the
neutron. Neutrons are subatomic particles with no
charge but with a mass nearly equal to that of a proton.
 Walter Bothe had first reasoned the existence of a third
subatomic particle in 1930.
 Bothe’s work was based in part on that of Henry Mosely
who showed by X-ray analysis that not all atoms of the
same element were identical. (Isotopes – 1907)

14. Counting Particles
 Atomic Number = number of protons
 Mass Number = number of protons and
neutrons
 Atomic Mass = average mass of the
isotopes
(also known as atomic weight)

15. Periodic Table
atomic number # of protons
mass number 8
-atomic number O round to 16 - mass
number ( # protons
# of neutrons & neutrons)
15.999
unrounded –mass
number (average
mass of the
isotopes)

16. Masses of Atoms
 A scale designed for atoms gives their small atomic
masses in atomic mass units (amu)
 An atom of 12C was assigned an exact mass of 12.00
amu
 Relative masses of all other atoms was determined
by comparing each to the mass of 12C
 An atom twice as heavy has a mass of 24.00 amu.
An atom half as heavy is 6.00 amu.

17. Atomic Mass
 Listed on the periodic table
 Gives the mass of “average” atom of each element
compared to 12C
 Average atom based on all the
isotopes and their abundance %.
 Atomic mass is not a whole number
Na
due to isotopes.
22.99

18. Isotopes
 Isotopes – atoms of the same element with different
numbers of neutrons.
Oxygen-16 Oxygen-17 Oxygen-18
16 17 18
8 8 8
p+ ‗‗‗‗ ‗‗‗‗ ‗‗‗‗
e- ‗‗‗‗ ‗‗‗‗ ‗‗‗‗
nº ‗‗‗‗ ‗‗‗‗ ‗‗‗‗

19. Calculating Average Atomic Mass
 Percent(%) abundance of isotopes
 Mass of each isotope of that element
 Weighted average =
mass isotope1(%) + mass isotope2(%) + …
100 100

20. Atomic Mass of Magnesium
Isotopes Mass of Isotope Abundance
24Mg = 24.0 amu 78.70%
25Mg = 25.0 amu 10.13%
26Mg = 26.0 amu 11.17%
Atomic mass (average mass) Mg = 24.3 amu
Mg
24.3

21. #16 The element copper has naturally occurring isotopes
with mass numbers of 63 and 65. The relative abundance
and atomic masses are 69.2% for mass = 63.0 amu, and
30.8% for mass = 65.0 amu. Calculate the average
atomic mass of copper.

22.  Finding An Isotopic Mass
Naturally occurring boron is 80.20% boron-
11 (atomic mass 11.0 amu) and 19.80% of a
different isotope of boron. What must the
mass of this isotope be if the average atomic
mass of boron is 10.81 amu?

23. Radioactivity
 Mosely’s X-ray analysis of atoms was an attempt to
explain radioactivity.
 1896 – Henri Becquerel – Uranium spontaneously emits
energy.
 1898 – Marie & Pierre Curie – first isolated a radioactive
element - Radium

24. Properties of Subatomic Particles
Particles Symbol Charge Relative Mass
Mass
Electron e- 1- 1/1840 amu 9.11 x 10-28 g
Proton p+ 1+ 1 amu 1.67 x 10-24 g
Neutron nº 0 1 amu 1.67 x 10-24 g

25. “Planetary” Model of the Atom
 Niels Bohr (1913) – developed the “planetary” model of
the atom based upon the following:
 Rutherford’s Gold Foil Experiment
 E = mc2 – Albert Einstein (1905)
 Quantum Theory – Max Planck (1910)

26. Atom
10-13 cm
electrons
protons
neutrons
nucleus
10-8 cm

27. Size of the Atom
Aluminum Atom
150 m
e-
1 mm Outside
edge of Al e- e-
atom stands
e-
e-
goal post
e-
nucleus - size
e- of a marble
e-
e-
e- Texas Memorial Stadium @ UT