Chemistry – Dr. May Notes
Atomic Structure
Early Atomic Models
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Democritus – empty space and tiny particles called atoms. This was the
earliest record of atomic theory.
Aristotle – continuous material and not small particles. Accepted until the
17th century.
Isaac Newton – atomic nature of elements. Newton had a theory but did no
experiments and had no proof.
Robert Boyle – atomic nature of elements. Boyle agreed with Newton.
John Dalton – offered hypothesis and studied the experimental data generated
by others.
Antoine Lavoiser – matter can not be created or destroyed. He discovered
that in a closed system the mass before a chemical
reaction was always equal to the mass after the reaction.
Joseph Proust – specific substances contain elements in the same ratio by
mass (The Law of Definite Proportions).
Dalton’s Hypothesis
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Matter has small particles called atoms
Atoms can not be broken down into smaller particles
All atoms of an element are alike
All atoms of a different element are different
Atoms combine in simple ratios to form compounds
Atoms are not destroyed but rearranged during reaction
Mass before a reaction must be equal to the mass after
All of an element have the same mass
Ratio of masses of one element that combine with the constant mass of
another element can be expressed in small whole numbers (Law of multiple
proportions)
Gay-Lussac – gases react in the ratio of small whole numbers
Avogadro – equal volume of gases have the same number of molecules
Avogadro’s number is 6.02 x 1023. That is the number of molecules or atoms (in
the case of elements) in one mole of anything.
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Early Research on Atomic Particles
Anode – Positive electrode
Cathode – Negative electrode
Cathode Rays – generated at the cathode and move toward the anode
Isotopes and Atomic Number
Isotopes – atoms of the same element with a different mass
Atomic Number – number of protons in an atom
Mass number – total of protons and neutrons
The Nuclear Atom
Rutherford (1912-1913) – gold foil experiment helped develop the modern
concept of atomic structure
Radioactivity – rays emitted by an unstable atomic nuclei
Parts of the Atom
Radiation:
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Alpha particles () – a helium nucleus of two protons and two neutrons
Beta particles () – a high energy electron
Gamma rays () – high energy X-rays
The Rutherford-Bohr Atom
The atom as a planetary model making the hydrogen atom similar to a solar
system consisting of a sun and one planet.
Plank’s Hypothesis – amount of energy given off in the electromagnetic spectrum is
directly proportional to the frequency of the light emitted.
The Hydrogen Atom and Quantum Theory
Activated electrons jump out to higher levels and give off characteristic
frequencies of energy (light) when they fall back to the ground state.
Atomic Mass - Protons and neutrons have essentially the same mass.
Electron = 9.10953 x 1028 grams = 0.000549 u
Proton = 1.67265 x 1024 grams = 1.0073 u
Neutron = 1.67495 x 1024 grams = 1.0087 u
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Average Atomic Mass
The atomic mass of an element as given on the periodic table is an average of all
the naturally occurring isotopes. The most abundant (98.89%) isotope of carbon (carbon
– 12 ) has 6 protons (atomic number) and 6 neutrons. The atomic mass is 6 + 6 = 12
(atomic mass). The least abundant (1.11%) isotope (carbon – 13) has 6 protons (atomic
number) and 7 protons. The atomic mass is 6 + 7 = 13 (atomic mass).
The periodic table gives the average atomic mass of carbon as 12.011. You can
easily calculate this if you know the relative abundance of the isotopes (12.0000 x .9889
+ 13.0034 x .0111) = 12.011.
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