The Atom

Because of the definition of the unified atomic mass unit, each carbon-12 atom has an atomic mass of exactly 12 u, and so a mole of carbon-12 atoms weighs exactly 0.012 kg. Neils Bohr’s model a. About The Atom Bistro and Coffee Bar is a small coffee shop born in a photography gallery on the North Umpqua in Glide, Oregon.

1. The Atomic Number
2. Atom Ray Palmer
3. The Atomic City Film

The Atomic Number

The number of protons, neutrons, and electrons in an atom can be determined from a set of simple rules.

• The number of protons in the nucleus of the atom is equal to the atomic number (Z).
• The number of electrons in a neutral atom is equal to the number of protons.
• The mass number of the atom (M) is equal to the sum of the number of protons and neutrons in the nucleus.
• The number of neutrons is equal to the difference between the mass number of the atom (M) and the atomic number (Z).

Examples: Let's determine the number of protons, neutrons, and electrons in the following isotopes.

The different isotopes of an element are identified by writing the mass number of the atom in the upper left corner of the symbol for the element. ¹²C, ¹³C, and ¹⁴C are isotopes of carbon (Z = 6) and therefore contain six protons. If the atoms are neutral, they also must contain six electrons. The only difference between these isotopes is the number of neutrons in the nucleus.

¹²C: 6 electrons, 6 protons, and 6 neutrons

¹³C: 6 electrons, 6 protons, and 7 neutrons

¹⁴C: 6 electrons, 6 protons, and 8 neutrons

Much of what is known about the structure of the electrons in an atom has been obtained by studying the interaction between matter and different forms of electromagnetic radiation. Electromagnetic radiation has some of the properties of both a particle and a wave.

Particles have a definite mass and they occupy space. Waves have no mass and yet they carry energy as they travel through space. In addition to their ability to carry energy, waves have four other characteristic properties: speed, frequency, wavelength, and amplitude. The frequency (v) is the number of waves (or cycles) per unit of time. The frequency of a wave is reported in units of cycles per second (s^-1) or hertz (Hz).

The idealized drawing of a wave in the figure below illustrates the definitions of amplitude and wavelength. The wavelength (l) is the smallest distance between repeating points on the wave. The amplitude of the wave is the distance between the highest (or lowest) point on the wave and the center of gravity of the wave.

If we measure the frequency (v) of a wave in cycles per second and the wavelength (l) in meters, the product of these two numbers has the units of meters per second. The product of the frequency (v) times the wavelength (l) of a wave is therefore the speed (s) at which the wave travels through space.

vl = s

Light is a wave with both electric and magnetic components. It is therefore a form of electromagnetic radiation.

Visible light contains the narrow band of frequencies and wavelengths in the portion of the electro-magnetic spectrum that our eyes can detect. It includes radiation with wavelengths between about 400 nm (violet) and 700 nm (red). Because it is a wave, light is bent when it enters a glass prism. When white light is focused on a prism, the light rays of different wavelengths are bent by differing amounts and the light is transformed into a spectrum of colors. Starting from the side of the spectrum where the light is bent by the smallest angle, the colors are red, orange, yellow, green, blue, and violet.

As we can see from the following diagram, the energy carried by light increases as we go from red to blue across the visible spectrum.

Because the wavelength of electromagnetic radiation can be as long as 40 m or as short as 10^-5 nm, the visible spectrum is only a small portion of the total range of electromagnetic radiation.

The electromagnetic spectrum includes radio and TV waves, microwaves, infrared, visible light, ultraviolet, x-rays, g-rays, and cosmic rays, as shown in the figure above. These different forms of radiation all travel at the speed of light (c). They differ, however, in their frequencies and wavelengths. The product of the frequency times the wavelength of electromagnetic radiation is always equal to the speed of light.

vl = c

As a result, electromagnetic radiation that has a long wavelength has a low frequency, and radiation with a high frequency has a short wavelength.

Unlike planets orbiting the Sun, electrons cannot be at any arbitrary distance from the nucleus; they can exist only in certain specific locations called allowed orbits. This property, first explained by Danish physicist Niels Bohr in 1913, is another result of quantum mechanics—specifically, the requirement that the angular momentum of an electron in orbit, like everything else in the quantum world, come in discrete bundles called quanta.

In the Bohr atom electrons can be found only in allowed orbits, and these allowed orbits are at different energies. The orbits are analogous to a set of stairs in which the gravitational potential energy is different for each step and in which a ball can be found on any step but never in between.

The laws of quantum mechanics describe the process by which electrons can move from one allowed orbit, or energy level, to another. As with many processes in the quantum world, this process is impossible to visualize. An electron disappears from the orbit in which it is located and reappears in its new location without ever appearing any place in between. This process is called a quantum leap or quantum jump, and it has no analog in the macroscopic world.

Because different orbits have different energies, whenever a quantum leap occurs, the energy possessed by the electron will be different after the jump. For example, if an electron jumps from a higher to a lower energy level, the lost energy will have to go somewhere and in fact will be emitted by the atom in a bundle of electromagnetic radiation. This bundle is known as a photon, and this emission of photons with a change of energy levels is the process by which atoms emit light. See alsolaser.

Atom Ray Palmer

In the same way, if energy is added to an atom, an electron can use that energy to make a quantum leap from a lower to a higher orbit. This energy can be supplied in many ways. One common way is for the atom to absorb a photon of just the right frequency. For example, when white light is shone on an atom, it selectively absorbs those frequencies corresponding to the energy differences between allowed orbits.

Each element has a unique set of energy levels, and so the frequencies at which it absorbs and emits light act as a kind of fingerprint, identifying the particular element. This property of atoms has given rise to spectroscopy, a science devoted to identifying atoms and molecules by the kind of radiation they emit or absorb.

The Atomic City Film

This picture of the atom, with electrons moving up and down between allowed orbits, accompanied by the absorption or emission of energy, contains the essential features of the Bohr atomic model, for which Bohr received the Nobel Prize for Physics in 1922. His basic model does not work well in explaining the details of the structure of atoms more complicated than hydrogen, however. This requires the introduction of quantum mechanics. In quantum mechanics each orbiting electron is represented by a mathematical expression known as a wave function—something like a vibrating guitar string laid out along the path of the electron’s orbit. These waveforms are called orbitals. See alsoquantum mechanics: Bohr’s theory of the atom.