Download Unit 3 Notebook Notes

Unit 3: History of the Periodic Table
 400 BC – Democritus thought matter could not be divided
indefinitely. This led to the idea of atoms in a void.
o “Atomos” Greek word for indivisible
o Could not be proven scientifically
 350 BC - Greek philosopher Aristotle said all matter is made of
Earth, Wind, Fire and Water.
o He was wrong, but his theory persisted for 2000 years
 1808 John Dalton developed the “Atomic Theory” based on
Experiments
1. All matter is made of small particles called atoms.
2. Atoms of an element are identical in size, mass and other
properties. Atoms of different elements differ in size,
mass, and other properties.
3. Atoms cannot be subdivided, created or destroyed.
4. Atoms of different elements combine in simple whole
number ratios to form compounds
5. In chemical reactions, atoms combine, separate
o Dalton’s theory explains the law of conservation of mass and
the law of constant composition.
o Later some parts of his theory were proven wrong. (by the
development of new technology)
o Dalton’s atom was compared to the Billiard ball, solid and
indivisible.
o
 1871 Dmitri Mendeleev, a Russian Chemist arranged the elements
in order based on increasing atomic mass
o He noticed a repeating pattern of chemical and physical
properties
 Hence the name “Periodic Table”
o Mendeleev left 3 blank spaced and predicted the discovery of
those elements.
 1897 J.J. Thompson
o He used a Cathode Ray Tube to discover the electron and its
charge.
o He could not find the positive charge but new it had to exist.
o Plum Pudding model. Electrons in a sea of positive material.
o
 1911 Ernest Rutherford
o Gold foil experiment using alpha particles to discover the
dense positive core of an atom, and that an atom is made of
mostly empty space.
 1911 Henry Mosley rearranged the periodic table based on
increasing atomic number
o He established Periodic Law with chemical and physical
properties as a function of their atomic number.
 1913 Niels Bohr
o Electrons in fixed paths, and why they did not attract to
positive nucleus
o Electron location cannot be predicted at a given moment
o Electrons travel so fast the form an electron cloud.
 1932 Sir James Chadwick proved/discovered Neutrons
Periodic Families and Properties
 Group 1: Alkali Metals
o In the pure state, they have a silvery appearance and are soft
enough to cut with a knife.
o They react violently with water and must be stored in oil.
o A very reactive family on the P. table
 Group 2: Alkali Earth Metals
o Less reactive than Alkali Metals, but are still too reactive to
be found in the pure form naturally.
 Group 17: Halogens
o Most reactive nonmetals
o They react with metals to form salts
 Group 18: Noble Gases
o Generally unreactive because they have a full valence shell of
electrons.
Periodic Trends
 Atomic Radius: ½ of the distance between the nuclei of identical
atoms that are bonded together
o Increases going down because in an increase in the size of the
electron cloud
 Ionization Energy:
o An Ion is an atom or group of bonded atoms that has a
positive or negative charge
 Cation is Positive charge
 Anion is Negative charge
o Any process that results in the formation of an ion is called
Ionization.
o Ionization Energy is the energy required to remove 1 electron
from a neutral atom of an element.
 Electron Affinity is the energy change that occurs when an
electron is acquired by a neutral atom.
 Electronegativity is a measure of an atom’s ability to attract
electrons from another element
Protons, Neutrons, and Electrons
 An Atom has 3 Parts
o Proton – Positive
o Neutron – No Charge
o Electron = Negative
 The Proton and Neutron are found in the nucleus
 Electrons Orbit the nucleus
 The Bohr Model
Lewis Dot Structure




Atomic Mass = Protons + Neutrons
Atomic Mass – Protons = Neutrons
Protons = Electrons (in a neutral Atom)
Protons = Atomic Number
 Isotopes – Atoms of the same element with a different mass
o Isotopes have the same number of protons, but a different
number of Neutrons
o Hyphen Notation Uranium-235
o Nuclear Symbol
 Calculating Average Atomic Mass
(Mass x Abundance) + (Mass x Abundance)……= Avg Atomic Mass
Example 1
Mass
15.995
16.995
17.999
Oxygen 16
Oxygen 17
Oxygen 18
% Abundance
99.759%
.037%
.204%
(15.995 x .99759) + (16.995 x .00037) + (17.999 x .00204) =
15.956
+ .00629 + .0367
= 15.999 Avg Atomic Mass