1. Atomic Structure and the Periodic Table
Adapted from Addison Wesley Chemistry

2. History of the atom
4th Century B.C. : Democritus suggested that matter was made up of very small particles called atoms.
: John Dalton studied atoms based on experiments and ratios. He created an Atomic Theory.

3. Dalton’s Atomic Theory
1. All elements are composed of tiny indivisible particles called atoms
2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
3. Atoms of different elements can physically mix, or chemically combine in simple whole number ratios to form compounds.
4. Chemical reactions occur when atoms separate, join or are rearranged. Atoms of one element may not be changed into atoms of another element in chemical reactions.
Dalton’s Atomic Theory

4. Electrons Electrons are negatively charged subatomic particles
1897: J.J. Thomson discovered the electron
He discovered that particles that were attracted to the metal plates in a cathode ray had a positive charge and those that were repelled had a negative charge, electrons.
He determined that these particles moved at a high speed and were 1/2000 the size of the mass of the hydrogen atom.
1916: Robert A. Millikan calculated an accurate mass of the electron.
Millikan found that electrons carries one unit of a negative charge and its mass is more precisely 1/1840 the mass of a hydrogen atom.

5. Atoms are electrically neutral or negatively charged particles will combine in whole number ratios with positively charged particles to achieve neutrality.
1886: E. Goldstein observed the proton with a positive charge opposite of the negatively charged electron. The mass of the proton was 1840 x the mass of the electron. This is represented as relative mass as 1.
1932: James Chadwick confirmed the existence of the neutron. This subatomic particle has no charge with the same relative mass as a proton, represented as 1.
1911: Ernest Rutherford performed the gold foil test directing a narrow beam of alpha particles on to a gold sheet of foil. Some particles were deflected back, He proposed the atom is mostly empty space with a concentrated nucleus, which is tiny compared to the atom, made up of neutrons and protons.
Protons and Neutrons

6. Perspective Rutherford’s experiment to the right
If an atom were the size of a football field, the nucleus would be the size of a marble in the center
The nucleus is made up of protons and neutrons in the center, with electrons moving in the rest of the space of the football field.

7. Atomic number
Atomic number = # of protons = # electrons in a neutral atom.
Atomic number identifies the element.
These can be found on the periodic table

8. Mass number/Atomic Weight
Mass number = Protons + Neutrons
Mass number/Atomic Weight/Atomic Mass Unit can be found on the Periodic table
Neutrons = Mass number – Protons
Neutrons are always calculated and are not found directly on the Periodic Table.
Mass number/Atomic Weight

9. Isotopes
Isotopes have the same number of protons, these never change, but different number of neutrons.
If neutrons change, the Mass changes.

10. How Mass is calculated due to Isotopes
Atomic Mass takes into consideration abundance and mass of Isotope
Use the number of masses and abundances given, this sample has 3.
Mass 1 x percent abundance 1 =__________+
Mass 2 x percent abundance 2 =__________ +
Mass 3 X percent abundance 3 =___________
Add answers and divide by 100, this takes you from percent to decimal
This equals Average Atomic Mass which is on the Periodic Table
How Mass is calculated due to Isotopes

11. Isotopes Isotopes, not just a Baseball team in Springfield
Protons never change, Neutrons do, so Mass does also
If you see an unfamiliar Mass, you may have an Isotope

12. The Periodic Table Development
Mid 1800’s : Dimitri Mendeleev listed elements in columns by increasing atomic mass and thus created the first Periodic Table
He left appropriate blank spaces for unknown elements and predicted the physical and chemical properties of the missing elements.
1913 : Henry Moseley determined the atomic number of elements. He arranged the Periodic Table by order of atomic number instead of atomic mass.
This is how the Periodic Table is arranged today.

13. Modern Periodic Table Horizontal Rows are called periods
Periodic Law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Because of Periodic Law, elements with similar properties end up in the same column called groups or families.
Group A are representative or main group elements
Group B are Transition and Inner Transition Metals

14. Metals, Non-Metals and Metalloids or Semi-Metals
Metals are on the left side, Non- Metals to the right and Metalloids down the step near the right side of the Periodic Table.
You may add At and Po to your Metalloids

15. Alkali Metals 1A: Alkali Metals less dense than other metals
one loosely bound valence electron
highly reactive, with reactivity increasing moving down the group
largest atomic radius of elements in their period
low ionization energy
low electronegativity

16. Alkaline Earth Metals 2A: Alkaline Earth Metals
two electrons in the valence shell
readily form divalent cations
low electron affinity
low electronegativity

17. Transition Metals Group B: Transition Metals
very hard, usually shiny, ductile, and malleable
high melting and boiling points
high thermal and electrical conductivity
form cations (positive oxidation states)
tend to exhibit more than one oxidation state
low ionization energy

18. Metalloids or Semi-Metals
Group 3A through 7A: Metalloids or Semi-Metals
electronegativity and ionization energy intermediate between that of metals and nonmetals
may possess a metallic luster
variable density, hardness, conductivity, and other properties
often make good semiconductors
reactivity depends on nature of other elements in the reaction

19. Non-Metals 4A through 8A: Non-Metals high ionization energy
high electronegativity
poor electrical and thermal conductors
form brittle solids
little if any metallic luster
readily gain electrons

20. Halogens 7A : Halogens extremely high electronegativity very reactive
seven valence electrons, so elements from this group typically exhibit a -1 oxidation state

21. Noble Gases 8A : Noble Gases
The noble gasses have complete valence electron shells, so they act differently. Unlike other groups, noble gasses are unreactive and have very low electronegativity or electron affinity.