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.