1. Atomic Structure
Chapter 4

2. Early Theories Ideas of the Philosophers
Democritus ( B.C.) : matter is composed of tiny particles called atoms
Atom: smallest particle of an element that retains its identity in a chemical reaction
Believed atoms were indivisible and indestructible
Lacked experimental support and couldn’t explain chemical behavior
Aristotle ( 384 – 322 B.C.) : disagreed with the idea of atoms. Thought everything was made from Earth, Wind, Water & Fire
John Dalton (1803) : marks the beginning of modern atomic theory. Based statements on scientific research

3. Dalton’s Atomic Theory
Using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory
All matter is composed of tiny, indivisible particles called atoms
Atoms of the same element are identical in size, mass, and chemical properties
Atoms of different elements are different
Atoms of different elements can physically mix or chemically combine
Chemical reactions occur when atoms are separated, joined, or rearranged.

4. Subatomic Particles The building blocks of atoms
Protons: positively charged subatomic particles
Discovered by Eugen Goldstein in 1886
Found in the nucleus
Relative electrical charge: 1+
Relative mass: 1
Symbol: p+

5. Subatomic Particles
2. Electron: negatively charged subatomic particles
Discovered by J.J. Thomson in 1897
He used a Cathode Ray tube
Mass 1/1840th of a proton
Relative electrical charge: 1-
Relative mass: 0
(negligible compared
to the proton)
Symbol: e-
Robert Millikan (1916) calculated the mass and charge of an electron

6. Subatomic Particles
3. Neutrons: neutral (uncharged) subatomic particles
Discovered by James Chadwick in 1932
Found in the nucleus
Relative electrical charge: zero
Relative mass: 1
Symbol: n0

7. Developmental Steps of Atomic Theory
Dalton’s Model
Indivisible Model
The same throughout
2. Thomson’s Model
“Plum Pudding Model”
Diffuse, evenly positive charge
Negative electrons in fixed positions throughout
3. Rutherford’s Model
Nuclear atom
Positive charge concentrated in a central core called a nucleus (1911)
Negative electrons in motion in the empty space around the nucleus

9. Rutherford’s Gold Foil Experiment
An experiment designed to verify Thomson’s model
The results led to the definition of the nucleus
Observations of Experiments
Conclusions about Nucleus
Most of the alpha (α) particles went straight through
- Large spaces between nuclei
- Electrons are rapidly moving in the empty space between nuclei
Some of the (α) particles deflected as they passed through
Very small positively charged nucleus
A few of the particles deflected back at large angles, never making it through
Nucleus is very dense and contains all of the positive charge

12. Nucleus Nucleus: Central core of the atom (1/100,000th diameter)
Composed of the p+ and n0
Contains all of the atoms positive charge
Contains 99.97% of the atomic mass

13. Distinguishing Among Atoms
Atomic Number: Number of protons in the nucleus
Henry Moseley discovered that each element contains a unique positive charge
Determines the place for each element on the periodic table
Whole # on the periodic table
(# of protons = # of electrons when uncharged)
Mass Number: total number of protons & neutrons found in the nucleus
Remember: all the mass of the atom is found in the nucleus
# n0 = mass # - atomic #

14. Isotopes Isotopes: Atoms that have the same number of protons but
different number of neutrons
Because they have a different number of neutrons, that also have a different mass number
Each isotope has essentially the same chemical properties because the protons and electrons are the same
More changes to Dalton’s Atomic theory
Atoms of the same element (same # p+) can be different due to the different number of neutrons (Isotopes)

15. Expressing Mass Number
Two types of shorthand notation for mass # Au
Au – 197
Gold - 197
Mass Number
197
Chemical Symbol 79
Atomic #
Chemical Symbol
Mass Number

16. Atomic Mass
Atomic Mass: Mass of the average atom based on the relative abundance of each the isotopes
This is the mass number you find on the periodic table
Isotope
Atomic #
Mass #
#p+
# e-
# n0
C-14 6 14 8 Pb 82
206 Ag
206 82
107 47

17. Atomic Mass Unit
Because the masses of atoms and subatomic particles are so small, a convenient unit is used:
Atomic mass unit (amu): a unit of mass equal to one- twelfth the mass of a carbon-12 atom
used to express the mass of atomic and subatomic particles.
1 amu = x g
1 proton or 1 neutron relative mass: 1 amu

18. Atomic Mass
A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature
To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products.
(mass isotope 1 x abundance isotope 1) + (mass isotope 2 x abundance isotope 2)

19. For example, carbon has two stable isotopes:
Carbon-12, which has a natural abundance of 98.89%, and
Carbon-13, which has a natural abundance of 1.11%.