Atomic Theory II

Atomic Theory II
Ions, Isotopes and Electron Shells

by Anthony Carpi, Ph.D.

In Atomic Theory I: The Early Days (see our Atomic Theory I module), we learned about the basic structure of the atom. Normally, atoms contain equal numbers of protons and electrons. Because the positive and negative charges cancel each other out, atoms are normally electrically neutral. But, while the number of protons is always constant in any atom of a given element, the number of electrons can vary.

Ions
When the number of electrons changes in an atom, the electrical charge changes. If an atom gains electrons, it picks up an imbalance of negatively charged particles and therefore becomes negative. If an atom loses electrons, the balance between positive and negative charges is shifted in the opposite direction and the atom becomes positive. In either case, the magnitude (+1, +2, -1, -2, etc.) of the electrical charge will correspond to the number of electrons gained or lost. Atoms that carry electrical charges are called ions (regardless of whether they are positive or negative). For example, the animation below shows a positive hydrogen ion (which has lost an electron) and a negative hydrogen ion (which has gained an extra electron). The electrical charge on the ion is always written as a superscript after the atom's symbol, as seen in the animation.

Hydrogen Ion Simulation

Isotopes
The number of neutrons in an atom can also vary. Two atoms of the same element that contain different numbers of neutrons are called isotopes. For example, normally hydrogen contains no neutrons. An isotope of hydrogen does exist that contains one neutron (commonly called deuterium). The atomic number (z) is the same in both isotopes; however the atomic mass increases by one in deuterium as the atom is made heavier by the extra neutron.

Hydrogen Isotope Simulation

Electron Shells
Ernest Rutherford's view of the atom consisted of a dense nucleus surrounded by freely spinning electrons (see our Atomic Theory I module). In 1913, the Danish physicist Niels Bohr proposed yet another modification to the theory of atomic structure based on a curious phenomenon called line spectra.

When matter is heated, it gives off light. For example, turning on an ordinary light bulb causes an electric current to flow through a metal filament that heats the filament and produces light. The electrical energy absorbed by the filament excites the atoms' electrons, causing them to "wiggle". This absorbed energy is eventually released from the atoms in the form of light.

When normal white light, such as that from the sun, is passed through a prism, the light separates into a continuous spectrum of colors:
spectrum-light - Continuous (white light) spectra

Continuous (white light) spectra

Bohr knew that when pure elements were excited by heat or electricity, they gave off distinct colors rather than white light. This phenomenon is most commonly seen in modern-day neon lights, tubes filled with gaseous elements (most commonly neon). When an electric current is passed through the gas, a distinct color (most commonly red) is given off by the element. When light from an excited element is passed through a prism, only specific lines (or wavelengths) of light can be seen. These lines of light are called line spectra. For example, when hydrogen is heated and the light is passed through a prism, the following line spectra can be seen:
spectrum-hydrogen - Hydrogen line spectra

Hydrogen line spectra

Each element has its own distinct line spectra. For example:
spectrum-helium - Helium line spectra

Helium line spectra
spectrum-neon - Neon line spectra

Neon line spectra

To Bohr, the line spectra phenomenon showed that atoms could not emit energy continuously, but only in very precise quantities (he described the energy emitted as quantized). Because the emitted light was due to the movement of electrons, Bohr suggested that electrons could not move continuously in the atom (as Rutherford had suggested) but only in precise steps. Bohr hypothesized that electrons occupy specific energy levels. When an atom is excited, such as during heating, electrons can jump to higher levels. When the electrons fall back to lower energy levels, precise quanta of energy are released as specific wavelengths (lines) of light.

Under Bohr's theory, an electron's energy levels (also called electron shells) can be imagined as concentric circles around the nucleus. Normally, electrons exist in the ground state, meaning they occupy the lowest energy level possible (the electron shell closest to the nucleus). When an electron is excited by adding energy to an atom (for example, when it is heated), the electron will absorb energy, "jump" to a higher energy level, and spin in the higher energy level. After a short time, this electron will spontaneously "fall" back to a lower energy level, giving off a quantum of light energy. Key to Bohr's theory was the fact that the electron could only "jump" and "fall" to precise energy levels, thus emitting a limited spectrum of light. The animation linked below simulates this process in a hydrogen atom.

Bohr's Atom: Quantum Behavior in Hydrogen

Concept simulation - Reenacts electron's "jump" and "fall" to precise energy levels in a hydrogen atom.

(Flash required)

Not only did Bohr predict that electrons would occupy specific energy levels, he also predicted that those levels had limits to the number of electrons each could hold. Under Bohr's theory, the maximum capacity of the first (or innermost) electron shell is two electrons. For any element with more than two electrons, the extra electrons will reside in additional electron shells. For example, in the ground state configuration of lithium (which has three electrons) two electrons occupy the first shell and one electron occupies the second shell. This is illustrated in the animation linked below.

The Lithium atom

For further details, the table linked below shows the electron configurations of the first eleven elements.

Atomic structure animation table

Source: http://www.visionlearning.com/library/module_viewer.php?mid=51

Atomic Theory I

Atomic Theory I

The Early Days

by Anthony Carpi, Ph.D.

Until the final years of the nineteenth century, the accepted model of the atom resembled that of a billiard ball - a small, solid sphere. In 1897, J. J. Thomson dramatically changed the modern view of the atom with his discovery of the electron. Thomson's work suggested that the atom was not an "indivisible" particle as John Dalton had suggested but, a jigsaw puzzle made of smaller pieces.

Thomson's notion of the electron came from his work with a nineteenth century scientific curiosity: the cathode ray tube. For years scientists had known that if an electric current was passed through a vacuum tube, a stream of glowing material could be seen; however, no one could explain why. Thomson found that the mysterious glowing stream would bend toward a positively charged electric plate. Thomson theorized, and was later proven correct, that the stream was in fact made up of small particles, pieces of atoms that carried a negative charge. These particles were later named electrons.

After Eugen Goldstein’s 1886 discovery that atoms had positive charges, Thomson imagined that atoms looked like pieces of raisin bread, a structure in which clumps of small, negatively charged electrons (the "raisins") were scattered inside a smear of positive charges. In 1908, Ernest Rutherford, a former student of Thomson's, proved Thomson's raisin bread structure incorrect.

Rutherford performed a series of experiments with radioactive alpha particles.  While it was unclear at the time what the alpha particle was, it was known to be very tiny.  Rutherford fired tiny alpha particles at solid objects such as gold foil.  He found that while most of the alpha particles passed right through the gold foil, a small number of alpha particles passed through at an angle (as if they had bumped up against something) and some bounced straight back like a tennis ball hitting a wall.  Rutherford's experiments suggested that gold foil, and matter in general, had holes in it!  These holes allowed most of the alpha particles to pass directly through, while a small number ricocheted off or bounced straight back because they hit a solid object.

In 1911, Rutherford proposed a revolutionary view of the atom. He suggested that the atom consisted of a small, dense core of positively charged particles in the center (or nucleus) of the atom, surrounded by a swirling ring of electrons. The nucleus was so dense that the alpha particles would bounce off of it, but the electrons were so tiny, and spread out at such great distances, that the alpha particles would pass right through this area of the atom. Rutherford's atom resembled a tiny solar system with the positively charged nucleus always at the center and the electrons revolving around the nucleus.

Interpreting Rutherford's Gold Foil Experiment

The positively charged particles in the nucleus of the atom were called protons.  Protons carry an equal, but opposite, charge to electrons, but protons are much larger and heavier than electrons.  

In 1932, James Chadwick discovered a third type of subatomic particle, which he named the neutron. Neutrons help stabilize the protons in the atom's nucleus. Because the nucleus is so tightly packed together, the positively charged protons would tend to repel each other normally. Neutrons help to reduce the repulsion between protons and stabilize the atom's nucleus. Neutrons always reside in the nucleus of atoms and they are about the same size as protons. However, neutrons do not have any electrical charge; they are electrically neutral.

Atoms are electrically neutral because the number of protons (+ charges) is equal to the number of electrons (- charges) and thus the two cancel out.  As the atom gets larger, the number of protons increases, and so does the number of electrons (in the neutral state of the atom).  The illustration linked below compares the two simplest atoms, hydrogen and helium.

Simulated hydrogen and helium atoms

Atoms are extremely small. One hydrogen atom (the smallest atom known) is approximately 5 x 10^-8 mm in diameter. To put that in perspective, it would take almost 20 million hydrogen atoms to make a line as long as this dash -. Most of the space taken up by an atom is actually empty because the electron spins at a very far distance from the nucleus. For example, if we were to draw a hydrogen atom to scale and used a 1-cm proton (about the size of this picture - ), the atom's electron would spin at a distance of ~0.5 km from the nucleus. In other words, the atom would be larger than a football field!

Atoms of different elements are distinguished from each other by their number of protons (the number of protons is constant for all atoms of a single element; the number of neutrons and electrons can vary under some circumstances). To identify this important characteristic of atoms, the term atomic number (z) is used to describe the number of protons in an atom. For example, z = 1 for hydrogen and z = 2 for helium.

Another important characteristic of an atom is its weight, or atomic mass. The weight of an atom is roughly determined by the total number of protons and neutrons in the atom. While protons and neutrons are about the same size, the electron is more that 1,800 times smaller than the two. Thus the electrons' weight is inconsequential in determining the weight of an atom - it's like comparing the weight of a flea to the weight of an elephant. Refer to the animation above to see how the number of protons plus neutrons in the hydrogen and helium atoms corresponds to the atomic mass.

Source: http://www.visionlearning.com/library/module_viewer.php?mid=50