Atomic Structure and the Periodic Table Adapted from Addison Wesley Chemistry

History of the atom 4th Century B.C. : Democritus suggested that matter was made up of very small particles called atoms. 1766-1844: John Dalton studied atoms based on experiments and ratios. He created an Atomic Theory.

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

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.

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

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.

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

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

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

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

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

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.

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

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

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

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

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

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

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

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

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.