The Atom: Nature's Building Blocks

Written by tutor Sam R.

Discovery, Structure, and Application

John Dalton: The Father of the Atom

Nearing the end of the 18th century, John Dalton observed three fundamental aspects of all matter:

1. Mass Conservation: Matter cannot be created or destroyed
2. Definite Composition: No matter the origin, a particular compound is made up of the same matter/elements in identical fractions.
3. Multiple Proportions: When one or more elements are combined together to form a compound, it is always done in the exact same proportion.

Discovery: Electricity?!

To understand the nature of the atom, one must return to the beginning and analyze observations which lead to our modern day schema. 19th century scientists began their odyssey into understanding the atom by observing electricity. The notion of matter and electric charge were understood but the actual components of the current were an anathema. First among many experiments are the Cathode Ray Tube Experiments. These cathode rays originate at a negative cathode and proceed towards a positive cathode. However, once the scientists introduced a magnet to the setting, they noticed these rays did not travel in straight lines. Rather, the rays were attracted towards the positive magnet, making them appear bent. Scientists understood that opposites attract and the current, which bent toward the positive end, must be composed of negatively charged "things" (we later call them electrons).

Discovery: From electrons to a nucleus!

After the discovery of the electron, physicists J.J. Thompson and Robert Milliken proceeded to illustrate the charge (-1.602x10-19 C) and mass (9.109x10-28g) of electrons as well. Following their observations, Thompson introduced the "Plum Pudding" model as shown below where the atom is just a positive sphere with negatively charged particles residing inside.

Scientist intuitively understand nature demands equilibrium, and that if an atom contains negative portions (electrons) it also must consist of positively charge parts to provide balance. Furthermore, if electrons are exceedingly small, what makes up the mass of the atom itself? Enter Ernest Rutherford, and his Gold Foil Experiments.

Approaching the end of the 20th century, Rutherford utilized alpha particles (a.k.a. helium) and beamed them at a thin sheet of gold foil. He noticed that rather than all atoms going through the foil, some began to rebound. The rebounding nature of these few alpha atoms implied that perhaps the negatively charged portions were not part of the nucleus itself and that the center contained a positive portion with substantial mass. Electrons, being so very small, could not have reflected the bulky alpha particles thus there must a large portion in the center having substantial mass with electrons revolving around this center. Rutherford is the first person to coin the term "nucleus" in relation to the atomic center.

Following Rutherford, James Chadwick in 1932 discovered the presence of an uncharged portion of the nucleus called the neutron. These developments led to the modern day structure of the atom. Namely, a nucleus containing protons/neutrons representing the bulk of the atoms mass with electrons orbiting around the center is the modern day representation of the atom.

Modern Day Structure of the Atom

Today’s atom is a spherical existence formulated from a positively charged nucleus consisting of protons/neutrons surrounded by one or more negatively charged electrons. These electrons orbit the nucleus and make up the greatest portion of the atom’s volume. However, it is the exceedingly dense center made up of protons and neutrons which contain the bulk of the atom’s mass. Protons have a charge of +1, electrons -1, and neutrons are uncharged but have mass nearly equivalent to that of protons. The neutral state of atoms exists because the charges of protons and electrons cancel each other out. All atoms of a unique element have the same atomic number and each element has a unique atomic number different from other elements. Each atom has a given atomic number (Z), mass number (A), and atomic symbol (X) as shown below.

The importance of these numbers cannot be understated and one must understand how to calculate them.

I.     Atomic Number (Z)
       a. This is the number of protons in the atom
       b. The number of protons of each atom is unique to an element
               i. Ex) Carbon: Z=6 because it has 6 protons in the nucleus
               ii. Ex) Oxygen Z=8 because it has 8 protons in the nucleus
       c. Forms the basis of the periodic tables organization

II.    Mass Number (A)
       a. This is the total sum of the protons and neutrons residing in the nucleus
       b. Protons = 1 mass unit; Neutrons = 1 mass unit
               i. Ex) Carbon: A=12
                       1. 6 protons + 6 neutrons = 12 mass units
               ii. Ex) Oxygen: A=14
                       1. 8 protons + 8 neutrons = 16 mass units
       c. To calculate the number of neutrons:
               i. A – Z = N (neutrons)

III.   Atomic Symbol (X)
       a. Every element has a unique symbol
               i. Oxygen: O
               ii. Carbon: C

Isotopes, Ions, and Atomic Mass...

All atoms of the same element have the same number of protons but can have different numbers of neutrons, leading to different atomic masses. Atoms with the same number of protons but different neutrons are called isotopes. Isotopes cause an element’s atomic mass to be different than its mass number because the atomic number is the average of the isotopes masses whereas the atomic number is the sum of protons and neutrons. For instance:

I.     Carbon: Mass number= 12 BUT Atomic Mass=12.01... why?
       a. Mass Number= neutrons (N) + protons (P)= 6+6=12
       b. Atomic Mass is average of masses, based upon abundance, of isotopes
               i. 12C (98.88%), 13C (1.11%), 14C(0.01%)
               ii. Atomic Mass= 12(.9988) + 13(.0111) + 14(.0001)= 12.01amu

But wait, where did the notion of atomic mass come from? The atomic mass unit (amu) is based upon the standard of a carbon-12 atom, and states that 1/12 of a carbon-12 atom is one amu.

With atomic mass and isotopes affecting the mass of the atom, the next portion deals with the charge of the atom. Ions are atoms of the same number of protons but different number of electrons. In the calculation of an atom’s charge, protons are +1, electrons are -1 and neutrons are neutral. The process of losing an electron makes an atom more positive while gaining an electron makes it more negative. An atom’s charge is indicated in the top right portion of the atomic symbol.

I.     Ex) Carbon
       a. 12C: 6 protons, 6 electrons
               i. (+6) + (-6) = 0
       b. 12C+1: 6 protons, 5 electrons
               i. (+6) + (-5)= +1
       c. 12C-1: 6 protons, 7 electrons
               i. (+6) + (-7)= -1

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