Download Atomic Structure

Early Greek philosophers proposed that all matter consisted of four fundamental substances:
fire, earth, water, and air. Plato (427-347 B.C.) and Aristotle (384-322 B.C.) believed that matter
was infinitely divisible into smaller units. Others, including Democritus (460-370 B.C.),
believed that all materials were composed of small indivisible particles called atomos
(indivisible). The lack of experimental evidence prevented a clear resolution of this debate
regarding the true nature of substances.
Modern chemistry’s qualitative roots may be traced back to the work of Georg Bauer (German
metallurgist, 1494-1555) and Philippus Theophrastus Bombastus von Hohenheim (Swiss
alchemist/physician, 1493-1541). Robert Boyle’s (British chemist, 1627-1691) work with
pressure-volume relationships for gases is recognized as the first modern example of quantitative
experimentation.
Georg Stahl (German chemist, 1660-1734) proposed that “phlogiston” was released by burning
materials. He proposed that a substance burned until its supple of phlogiston was exhausted.
Karl W. Scheele (Swedish chemist, 1742-1768) first discovered oxygen, but Joseph Priestly
(English clergyman/scientist, 1733-1804) commonly receives the credit for the discovery since
he published his findings first.
Antoine Lavoisier’s (French chemist, (1743-1794) quantitative experimentation lead to the true
nature of combustion. Lavoisier’s work lead to development of the law of conservation of
mass. Joseph Proust (French chemist, 1754-1826) demonstrated that compounds have a constant
(definite) composition by mass – “Proust’s law” is also known as the law of definite
proportion. John Dalton (English school teacher, 1766-1844) discovered that elements can
combine in different ratios to form different substances – this is the law of multiple
proportions.
In 1808, Dalton published his atomic theory in A New System of Chemical Philosophy. The four
main points in his theory are:
1.
All forms of matter contain extremely small indivisible particles – atoms – that have their
own unique chemical properties. (We now know that atoms are divisible into subatomic
particles such as the proton, electron, etc.)
2.
An element consists of identical atoms possessing the same properties (including mass –
this part of Dalton’s atomic theory had to be revised due to the discovery of isotopes).
3.
A chemical reaction rearranges the atoms to produce a different substance without
creating or destroying the individual atoms.
4.
A compound is formed when atoms of multiple elements chemically combine in a fixed
ration by mass (we now know that it is an atom ratio) to form a new substance.
In 1897, J.J. Thomson (British physicist, 1856-1940) connected a high-voltage source to a sealed
evacuated glass tube. The charge flow from the cathode (negative electrode) toward the anode
(positive electrode) through the tube produced a green beam that was deflected by charged
electrical plates (from negative to positive). This electron or “cathode ray” thus consists of
negatively charged particles. In 1909, Robert Millikan (U.S. physicist, 1868-1953) determined
that the charge of an electron was 1.602 x 10-19 coulomb. Further, the electron’s mass-to-charge
yields an electron mass of 9.109 x 10-31 kg – about 1/1836th the mass of a hydrogen atom.
Thomson proposed that atoms consisted of positive and negative charges distributed like plums
in a pudding. This atomic model was later updated based on new experimental evidence.
In 1896, Henri Becquerel (French scientist, 1852-1908) discovered that pitchblende (a uranium
containing mineral) spontaneously released high-energy radiation (radioactivity). Marie Curie
(Polish chemist, 1867-1934)
Ernest Rutherford (British physicist, 1871-1937) helped identify three types of radioactivity:
alpha (α, a helium nucleus with a 2+ charge and mass number = 4), beta (β, an electron with a 1charge and a mass number = 0), and gamma (γ, high energy electromagnetic radiation with 0
charge and a mass number = 0).
Based on the results of Hans Geiger’s (German physicist, 1882-1945) and Ernest Marsden’s
(British physicist, 1888-1970) “gold foil experiment”, Rutherford proposed that atoms have an
extremely small, massive, and positively charged central core – the nucleus. Geiger and
Marsden bombarded thin sheets of gold foil with alpha particles and discovered that about 1 in
8000 alpha particles were scattered at a large angle (sometimes backward). This experimental
evidence lead to the conclusion that almost all of an atom’s mass is concentrated in a small,
positive, central core. Typical atomic diameters are about 10-10 m and the diameter of the
nucleus is about 10-15 m.
In 1911, Rutherford discovered that protons (hydrogen nuclei) were released when nitrogen
atoms were bombarded with alpha particles. The proton has a charge of 1+ (the electron has a
charge of 1-) and a mass about 1800 times greater than an electron. Rutherford later (around
1920) proposed the existence of a neutral particle in the nucleus.
Another of Rutherford’s students, Henry Gwyn-Jefferies Moseley (British physicist, 1887-1915)
used his experimental work with x-rays to develop the concept of atomic number (Z) – the
number of protons in an atom’s nucleus. The discovery of the atomic number lead to a change in
one of Dalton’s postulates – atoms of the same element have the same atomic number (number
of protons in their nucleus). The atomic number concept helped explain and correct the apparent
misplacement of nickel (Ni) and cobalt (Co) as well as Iodine (I) and Tellurium (Te) on
Mendeleev’s original Periodic Table. Remember that Mendeleev arranged the elements by
increasing atomic weight!
Francis William Aston (British physicist, 1877-1945), inventor of the mass spectrograph, used
mass spectrographic data to propose that many elements had more than one type of atom having
different masses – isotopes. Hydrogen has three naturally occurring isotopes: 1H, 2H, and 3H .
An atom’s mass number (A) is the sum of its protons and neutrons. Carbon-14 is a radioactive
isotope that has six protons and eight neutrons in its nucleus. The symbol 146 C shows that carbon
has an atomic number (lower left) of six and a mass number (upper left) of 14.
James Chadwick (British physicist, 1891-1974) used alpha particle bombardment of beryllium
atoms to initiate the release of highly penetrating radiation consisting of neutral particles –
neutrons. Neutrons are slightly more massive than protons.
Particle
Mass (kg)
Charge (C)
Mass (amu)
-31
-19
Electron
9.10939 x 10
-1.60218 x 10
0.00055
Proton
1.67262 x 10-27
+1.60218 x 10-19
1.00728
-27
Neutron
1.67493 x 10
0
1.00866
th
Table 2.2, p. 51, General Chemistry, Ebbing & Gammon, 7 Ed. (2002)
Charge (e)
-1
+1
0
Due to the existence of isotopes, the atomic masses (weights) shown on the Periodic Table is a
“weighted average” based on the relative abundance of the naturally occurring isotopes and their
respective relative atomic mass. Dalton assigned atomic weights relative to hydrogen. In 1961,
the International Union of Pure and Applied Chemistry (IUPAC) redefined atomic weights to be
relative to the most abundant isotope of carbon: carbon-12.
The atomic mass unit (amu) is defined as being equal to 1/12th the mass of a single atom of 12C.
The mass spectrograph is used to determine the relative masses of the atoms of other isotopes
and/or elements. The mass spectrograph can show the relative mass and fractional abundance
for the isotopes for an element.
Isotope
Neon-20
Neon-21
Neon-22
Isotopic Mass (amu)
19.992
20.994
21.991
Fractional Abundance
0.9051
0.0027
0.0922
Atomic Mass =
Mass Contribution
18.09
0.06
2.03
20.18