In 1869, Russia's Dmitri Mendeleev and Germany's Lothar Meyer published nearly identical element classificaiton schemes, recognizing similar chemical and physical properties recur periodically when elements are arranged in order of increasing atomic weight.
Although their observations were nearly identical, Mendeleev is given more credit because his first published table predicted the existance of undiscovered elements, and left spaces in his table for them - even predicting their properties. Meyer left blanks in his second published table in 1870, not his first. Because of this, most of the world (except Germany) considers Mendeleev the Father of the Periodic Table.
In 1913, Henry Moseley developed the concept of atomic numbers - correctly stating that the atomic number was equal to the number of protons in the nucleus of an atom and the number of electrons surrounding the nucleus.
Arranging atoms by their increasing atomic number brought the elements in line with todays recognized periodic trends.
In 1945, Glenn Seaborg grouped the "transuranium" elements into the Lanthanide and Actinide series and proposed pulling them out of the main body of elements on the table. This arrangement does a better job grouping atoms according to the quantum numbers.
A Periodic Table is based on the atomic theory. The history of the Periodic Table traces our understanding of the atom.
Make notes on the front and back and use this "Paper Periodic Table" on any class test. If it is lost, you may not make a photocopy of another student's table. You may print another copy and re-write your notes onto it. Students must have their OWN Paper Periodic Table, or use the one on the classroom wall. No other notes or computers may be used on "closed-computer" tests.
The class Periodic Table is online. The icon at the top of the left column of any chemistry class page will bring it up in a new window on your computer.
Developing The Atomic Theory of Matter
The word "atom" dates to the time of Democritus (about 400 BC), who thought that the material world must be made up of tiny indivisible particles he called "atomos", meaning indivisible or uncuttable.
Aristotle (about 340 BC), proposed that there can be no ultimately indivisible particles and the "atomic" view of matter faded for centuries.
Sir Isaac Newton, and others, (about 1700 AD) favored the idea of atoms when trying to explain the observed behavior of gases.
Understanding atoms as the fundamental building blocks of all elements begins with John Dalton (about 1800).
Dalton's atomic theory proposed the following:
Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical to one another in mass and other properties, but atoms of one element are different from the atoms of all other elements.
Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.
In 1897, J. J. Thomson published a paper describing cathode rays are a stream of negatively charged particles. This paper is generally accepted as the "discovery" of electrons. Thomson's cathode ray experiments made it possible to calculate the charge-to-mass ratio for an electron to be 1.76 X 108 Coulombs per gram.
In 1909, the "oil drop" experiment of Robert Millikan determined the actual charge on an electron to be 1.603 X 10−19 Coulombs. He then calculated the mass of an electron to be 9.10 X 10−28 gram.
In the early 1900s, Thomson proposed that the atom consisted of a uniform positive sphere of matter in which the electrons are embedded - the "plum-pudding" model.
Based on his "alpha-scattering" experiment in 1911, Ernest Rutherford proposed that most of the mass of a gold atom and all of its positive charge reside in a very small, extremely dense region, which he called the nucleus. He further suggested that most of the total volume of an atom is empty space in which electrons move around the nucleus.
In 1913, the Danish scientist Niels Bohr proposed the planetary model, in which electrons more in definite orbits around the nucleus - like planets moving around the Sun. Bohr suggested that each electron can only move around the nucleus at specific distances, or energy levels.
Protons (positive particles in the nucleus) were discovered in 1919 by Rutherford.
Neutrons (neutral particles in the nucleus) were discovered in 1932 by James Chadwick.
The first experimental evidence of antimater can in a particle track observed in a cloud chamber at Caltech by physicist Carl Anderson in 1932. The track curved in the opposite direction of the track of an electron, showing that the particle had a positive charge - the positron.
Today's quantum model explains the structure of atoms using 6 types of quarks, 6 types of leptons, and 4 types of force carriers. The model even supports the theory of antimatter - in which case, each of the 12 particles has a corresponding antimatter particle.
The Standard Model of Particle Physics is left to the physists and their supermachines. The accelerators below, and the extreme amounts of energy they produce, are needed to study the particles and forces in the atomic nucleus. The nuclear shell model represents the current understanding of the atomic nucleus.
Accelerators can be used to create new elements. The atomic theory predicts the characteristics of these elements as they are arranged into an "extended" periodic table.
Einstein's equation, E = mc2, involves the binding energy holding the atomic nucleus together and expresses the concept of mass-energy equivalence.
Using this equation to express mass as energy,
E / m = c2 = (299,792,458 m/s)2 ≈ 9 X 1016 joules per kilogram
So one gram of mass − approximately the mass of a U.S. dollar bill − is equivalent to the following amounts of energy:
24.9 million kilowatt-hours (≈25 GW⋅h) The electrical energy produced by Grand Coulee Dam's turbines every 3.7 hours.
21.5 billion kilocalories (≈21 Tcal)
21.5 kilotons of TNT
85.2 billion BTUs
Chemists are concerned with only three particles - the proton, neutron, and electron - because they explain the basic chemical behavior of atoms.
Normal chemical reactions only involve valence electrons, those farthest from the nucleus. Chemical reactions produce a small fraction of the energy produced by nuclear reactions.
Atoms have extremely small masses. Because of this, the mass of atoms is generally expressed in atomic mass units, not grams or kilograms.
Atomic mass units may be abbreviated amu.
The formal SI abbreviation for the atomic mass unit is u.
1 u = 1.660538782 X 10−27 kg
A proton has a mass of 1.0073 u.
A neutron has a mass of 1.0087 u.
An electron has a mass of 5.486 X 10−4u. It would take 1836 electrons to equal the mass of 1 proton.
Most atoms have diameters between 100-500 picometers (pm). Although not an SI unit of length, the angstrom, , is often used to express atomic dimensions. One angstrom equals 10−10 (1 ten-billionth) meter. In everyday terms, a sheet of paper is approximately 1,000,000 angstroms thick. The diameter of a hydrogen atom is about 1 angstrom.
The atomic number of an element is the number of protons in the nucleus of its atoms.
The mass number of an atom is equal to the total number of protons plus neutrons in the atom.
Atoms with identical atomic numbers but different mass numbers are called isotopes. Isotopes have the same number of protons but different numbers of neutrons.
Isotopes are represented by symbols such as this: . The symbol represents "carbon twelve", carbon-12. The atomic number is shown by the subscript and the mass number is shown by the superscript.
Most elements occur in nature as mixtures of isotopes. The average atomic mass, also called atomic weight, of an element is determined by using the masses of its various isotopes and their relative abundances.
Calculating Average Atomic Mass
Natural carbon is composed of 98.93% carbon-12 (exactly 12 u) and 1.07% carbon-13 (13.00335 u).
In a normal atom # of protons = # of electrons, making the number of positively charged particles and the number of negatively charged particles equal.
The mass number is equal to the total of the protons and neutrons.
neutrons = mass # - atomic #
Since an atom will not have fractions of neutrons, we will round off the mass number to the nearest whole number in this class.
Protons and neutrons are made up of even smaller particles called quarks.
The nucleus of an atom is unchanged during a chemical reaction - only nuclear reactions can do that. But atoms readily gain or lose electrons during chemical reactions.
When electrons are added to or removed from an atom, a charged particle called an ion is formed.
An ion with a positive charge, (electrons removed),
is called a cation.
An ion with a negative charge, (electrons added),
is called an anion.
The chemical properties of ions are very different from the chemical properties of the atoms from which they form.
A naming controversy arose with the man-made elements, beginning with #104. The International Union of Pure and Applied Chemistry (IUPAC) now assigns a temporary systematic name and a 3-letter symbol to new or theoretical elements.
Chemical symbols are shown in the following colors on the class Periodic Table:
Many, but not all, Periodic Tables have a line dividing these two. The dividing line begins between boron and aluminum and stair-steps down and to the right, one square at a time. Metals are left of the line, nonmetals are to the right.
When moving from left to right across the Periodic Table, metal characteristics decrease and nonmetal characteristics increase. In the area of the dividing line, some elements have both metallic and nonmetallic properties. These elements are often called metalloids.
Almost 3/4 of the elements are metals.
Electron Energy Levels
The number of electron energy levels in an atom is indicated by the horizontal row on which the element is found on the Periodic Table. The horizontal rows are numbered 1 to 7 at the extreme left of the table.
Electron Energy Sublevels
The shape of the Periodic Table makes the four sublevels easy to see.
s sublevel - the two tall columns on the left.
Helium is at the far right because it is inert (like the other Nobel Gases). For the purpose of reading sublevels, it should be thought of as sitting in the second column beside hydrogen.
p sublevel - the six tall columns on the right, without helium.
d sublevel - the ten short columns in the middle of the table.
f sublevel - the fourteen columns of two below the main body of the table.
These "areas" on the Periodic Table are often refered to as "blocks":