1. P. Sci.
Unit 8
Chapter 4
Atoms

2. Atomic Structure – timeline
Ancient Greece - Democritus proposed the atom – a tiny solid particle that could not be subdivided.
1904 – J.J. Thomson – discovered that atoms contained small, negatively charged particles called electrons.

3. Dalton’s Atomic Theory (#9)
Every element is made of tiny, unique, particles called atoms that cannot be subdivided.
Atoms of the same element are exactly alike.
Atoms of different elements can join to form molecules or compounds

4. 1911 – Ernest Rutherford – proposed that the atom had two parts – the nucleus in the center (most of the mass) surrounded by the electrons.
1913 – Niels Bohr –
hypothesized that
electrons traveled in fixed
orbits around the atom’s nucleus.

5. 1913 – James Chadwick – concluded that the nucleus contained positive protons and neutral neutrons.
1926 – Erwin Schrodinger – developed the quantum mechanical model – which is based on the wavelike properties of the electron. (not a particle – leads to quantum physics)

6. 1927 – Werner Heisenberg – (the Heisenberg uncertainty Principle) described that it is impossible to know precisely both an electron’s position and path at a given time. Led to the electron cloud theory.

7. Atoms
The smallest particle that has the properties of an element.

8. Parts of an Atom
Nucleus – small, dense center of an atom made up of 2 subatomic particles that are identical in size and mass.
Protons – have a
positive charge
Neutrons – have no charge

9. Parts of an Atom cont.
Electrons – are tiny subatomic particles that have very little mass that moves around the outside of the nucleus. These particles are negatively charged and form a
“cloud” around
the nucleus.

10. The number of protons and electrons an atom has is unique for each element.

11. Atomic Charge
Atoms have no overall charge because the protons (+) cancel out the electrons (-).
Helium protons
2 neutrons
2 electrons
total charge 0

12. Protons positive (+) charge Found in the nucleus
# of protons = atomic #
The number of protons identify the element (atomic #)

13. Neutrons no charge Found in the nucleus
Along with protons makes up atomic mass
Atomic Mass – atomic number = # of neutrons
(rounded to whole #)

14. Electrons negative (–) charge
travel in orbitals (or energy levels) around the nucleus. (electron cloud)
Equals atomic number in neutral atoms
valence electrons - the # of electrons in the outer shell and relates to the oxidation #

15. Unit of measure for atomic particles is Atomic mass unit (amu) protons and neutrons = about 1 amu (electrons are about 1/2000 of the size of protons and neutrons))

16. Chemical symbols
The one or two letter abbreviation of the element name.
Some are based on Latin name
ALL 1st letter is upper case
ALL 2nd letter is lower case

17. Mass Number or Atomic Mass
the sum of the number of protons and the number of neutrons in the nucleus of an atom.
# of neutrons =
mass # - atomic #
Neutrons
Protons +

18. Atom Summary Atomic Number = protons = electrons
Atomic Mass = Protons + Neutrons
Neutrons = atomic mass – atomic number
Atomic symbols
First letter is ALWAYS upper case
Second letter is ALWAYS lower case
Example: Identify the Number of Protons, Neutrons and Electrons in Oxygen
Oxygen element 6 with mass 16
P = 8 N = 16-8 = 8
E = 8

19. Atom Summary
Example 2: Identify the Number of Protons, Neutrons and Electrons in Sulfur and Sodium

20. Isotopes
Atoms of the same element that have a different # of neutrons and a different atomic mass. (identified by the element name followed by the mass # )
ex. C-12, C-14, B-10, B-11)
Carbon 14 = 8 neutrons
6 protons
6 electrons
6 electrons
6 protons
Carbon 12 = 6 neutrons

21. Average atomic mass
the weighted - average mass of the mixture of all an atoms isotopes. The average atomic mass is close to the mass of its most abundant isotope.
This is the number found on the periodic table

22. Using Bohr’s Atomic Model
Bohr was the 1st person to propose the concept of electrons having specific energy levels
This explained how electrons could give off light ( by gain or lose energy)
Bohr’s Model can be used to show how the electrons will arrange themselves around a nucleus.

23. BOHR MODEL
Electrons are placed in energy levels surrounding the nucleus
You fill the lower energy levels first
1st energy level can hold max of 2 electrons
2nd energy level can hold max of 8 electrons
3rd energy level can hold max of 18 electrons
4th energy level can hold max of 32
18e-
8e-
Nucleus
(p+ & n0)
2e-

24. Example
Chlorine 17 protons 17 electrons 18 neutrons Valance Electrons: Electrons found in outer energy level. Chlorine has 7 electrons in outer energy level (ring 3)