This document provides an overview of the history and development of atomic theory from Dalton's model to the modern electron cloud model. It discusses key scientists and experiments that led to important discoveries, such as Thomson discovering electrons, Rutherford determining the structure of the nucleus, and Bohr proposing electron energy levels. The document also defines important atomic concepts like atomic number, mass number, isotopes, and how elements are distinguished by their number of protons.

1. Introduction to Atoms
CHAPTER 4
SECTION 1

2. History of Atom
 All atoms share the same basic structure
 During past 200 years, scientists have proposed
different models.
 An atom is the smallest particle of an element.
 Atomic theory grew as a series of models that
developed from experimental evidence. As more
evidence was collected, the theory and models were
revised.

3. Dalton’s Model
 Based on experiments, Dalton developed a theory
of structure of matter
 4 main concepts:
 All matter is composed of tiny, indivisible particles called
atoms
 Atoms of each element are exactly alike
 Atoms of different elements have different masses
 Atoms of different elements can join to form compounds

4. Dalton’s Model

5. Thomson’s Model
 End of 1800s
 Thomson discovered that atoms were not
simple, solid spheres
 Atoms contained subatomic particles
 Very small, negatively charged
 Called them electrons

6. Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray
tube to deduce the presence of a negatively
charged particle: the electron

7. Modern Cathode Ray Tubes
Television Computer Monitor
Cathode ray tubes pass electricity
through a gas that is contained at a
very low pressure.

8. Thomson’s Model
 Also knew that atoms were electrically neutral
 Must contain enough positive charge to balance negative
charge of electrons
 Developed model where electrons were stuck into a
positively charged sphere
 Like chocolate chips in cookie dough

9. Thomson’s Model

10. Rutherford’s Model
 By early 1900s, scientists knew that positive charge
of atom comes from subatomic particles called
protons
 A proton is a positive charged particle in the
nucleus of an atom.
 1911—Rutherford begins to test theory
 His experiments led him to believe that protons are
concentrated in a small area at center of atom
 Called this area the nucleus

11. Rutherford’s Model
 Rutherford’s model describes an atom as mostly
empty space, with a center nucleus that contains
nearly all the mass
 Like the pit in a peach

12. Bohr’s Model
 Modified Rutherford’s model in 1913
 Proposed that each electron has a certain amount of
energy
 Helped electron move around nucleus
 Electrons move around nucleus in region called
energy levels
 The energy level is the specific amount of energy
an electron has.
 Energy levels surround nucleus in rings, like layers of
onion

13. Bohr’s Model
 Has been called planetary model
 Energy levels occupied by electrons are like orbits of planets at
different distances from the sun (nucleus)

14. Electron Cloud Model
 Model accepted today
 Electrons dart around in an energy level
 Rapid, random motion creates a “cloud” of negative
charge around nucleus
 Electron cloud gives atom its size and shape

15. Electron Cloud Model

16. Findings
 Eugen Goldstein in 1886 observed
what is now called the “proton” -
particles with a positive charge, and
a relative mass of 1 (or 1840 times
that of an electron)
 1932 – James Chadwick confirmed
the existence of the “neutron” – a
particle with no charge, but a mass
nearly equal to a proton

17. Atomic Number
 Atoms are composed of identical
protons, neutrons, and electrons
 How then are atoms of one element different
from another element?
 Elements are different because they
contain different numbers of PROTONS
 The “atomic number” of an element is
the number of protons in the nucleus
 # protons in an atom = # electrons

18. Atomic Number
Atomic number (Z) of an element is
the number of protons in the nucleus
of each atom of that element.
Element # of protons Atomic # (Z)
Carbon 6 6
Phosphorus 15 15
Gold 79 79

19. Mass Number
Mass number is the number of
protons and neutrons in the nucleus
of an isotope: Mass # = p+ + n0
Nuclide p+ n0 e- Mass #
Oxygen - 18 8 10 8 18
Arsenic - 75 33 42 33 75
Phosphorus - 31 15 16 15 31

20. The Complete Set-UP
 Contain the symbol of the
element, the mass number and the
atomic number.
Superscript →
Mass
number
Subscript →
Atomic
number
X

21. Isotopes
Dalton was wrong about all
elements of the same type being
identical
Atoms of the same element can
have different numbers of
neutrons.
Thus, different mass numbers.
These are called isotopes.

22. Isotopes
 Frederick Soddy (1877-1956)
proposed the idea of isotopes in
1912
 Isotopes are atoms of the same element having
different masses, due to varying numbers of
neutrons.
 Soddy won the Nobel Prize in Chemistry
in 1921 for his work with isotopes and
radioactive materials.

23. Naming Isotopes
We can also put the mass
number after the name of the
element:
carbon-12
carbon-14
uranium-235

24. Isotopes are atoms of the same element having
different masses, due to varying numbers of
neutrons.
Isotope Protons Electrons Neutrons Nucleus
Hydrogen–1
(protium) 1 1 0
Hydrogen-2
(deuterium) 1 1 1
Hydrogen-3 1 1 2
(tritium)

25. Isotopes
Elements
occur in
nature as
mixtures of
isotopes.
Isotopes are
atoms of the
same element
that differ in
the number of
neutrons.