Thomson Atom:
The next great step forward in the understanding of atoms was accomplished by John Thomson. Using a cathode ray scope, Thomson determined that all matter, whatever its source, contains particles of the same kind that are much less massive than the atoms of which they form a part. They are now called electrons, although he originally called them corpuscles.
His discovery was the result of an attempt to solve a long-standing controversy regarding the nature of cathode rays, which occur when an electric current is driven through a vessel from which most of the air or other gas has been pumped out.
By applying an improved vacuum technique, Thomson was able to put forward a convincing argument that these rays were composed of particles. Furthermore, these rays seemed to be composed of the same particles, or corpuscles, regardless of what kind of gas carried the electric discharge or what kinds of metals were used as conductors.
Thomson's conclusion that the corpuscles were present in all kinds of matter was strengthened during the next three years, when he found that corpuscles with the same properties could be produced in other ways; e.g., from hot metals. Thomson may be described as "the man who split the atom" for the first time, although "chipped" might be a better word, in view of the size and number of electrons.
Rutherford Atom:
Ernest Rutherford is considered the father of nuclear physics. Indeed, it could be said that Rutherford invented the very language to describe the theoretical concepts of the atom and the phenomenon of radioactivity. Particles named and characterized by him include the alpha particle, beta particle and proton. Rutherford overturned Thomson's atom model in 1911 with his well-known gold foil experiment in which he demonstrated that the atom has a tiny, massive nucleus.
By the turn of the 20th century, physicists knew that certain elements emitted fast moving particles of two flavors, alpha particles and beta particles. These elements were typically very heavy (i.e. their atom nuclei were massive) such as uranium and radium. Today we know that heavy nuclei are unstable and `decay', meaning that they spontaneously split into smaller nuclei and emit stray particles. This is called radioactivity.
The alpha particle was heavy and positively charged, we now know that it is the helium nuclei (2 protons and 2 neutrons). The beta particle was light and negatively charged, the electron. Rutherford designed an experiment to use the alpha particles emitted by a radioactive element as probes to the unseen world of atomic structure. His experiment looked like the following:
The Rutherford beamed alpha particles through gold foil and detected them as flashes of light or scintillations on a screen. The gold foil was only 0.00004 centimeter thick, meaning on a few hundreds of atoms thick.
The expectation is that they will strike the fluorescent screen directly behind the foil.
These results can best explained by a model for the atom as a tiny, dense, positively charged core called a nucleus, in which nearly all the mass is concentrated, around which the light, negative constituents, called electrons, circulate at some distance, much like planets revolving around the Sun.
The Rutherford atomic model has been alternatively called the nuclear atom, or the planetary model of the atom.
Planck's curve:
Parallel to efforts to understand the inner workings of the atom was research in the late-1800's to understand how matter emits energy. The energy an object emits as a function of the wavelength of light is called its continuous spectrum. The field of science that studies spectrum is called spectroscopy. One of the primary results from the field of spectroscopy was the discovery that the spectrum of an object changes with temperature.
In particular, was the formulation of the two laws of radiation:
Both these relationships were synthesized by physicist Max Planck into what is called Planck's curve. All objects emit energy in the shape of Planck's curve, where the amount and the peak energy vary only as the temperature of the body.
The problem with Planck's curve is that it does not agree with Rutherford's model of the atom. Atoms absorb and emit light through a process called scattering. Photons fly near the atoms and are deflected. Sometimes their motion pushes the atom (where push means electromagnetic forces), and the photon loses energy (i.e. becomes redder). Sometimes the atom pushes the photon, and the photon gains energy (i.e. becomes bluer). However, even the nature of atoms, the photons should receive more energy than the atoms, so there should be more and more blue photons, but clearly the Planck curve drops off at short wavelengths. This is called the UV catastrophe, which was resolved by quantum physics.
Stellar Spectroscopy:
The fact that the temperature and energy generation of an object can be determined by its Planck curve made the study of spectrum the key component to stellar astronomy. The amount of energy emitted from stars is determined by measuring their brightness or the amount of light they emit. This is called photometry. However, two major developments expanded our understanding of the chemical make-up of stars. They were:
The discovery of spectra lines was made by Fraunhofer who, in the early 1800's, magnified the Sun's spectrum and discovered dark lines which could be identified with particular elements (based on spectra in the laboratories).
English astronomer Lockyer, in the late-1800's, discovered an unknown element in the Sun, i.e. a set of spectral lines which did not correspond to elements in the lab. He named this element helium (Latin for Sun element).
By the 20th century, spectral lines for all the elements in the periodic table have been classified and astronomers could examine the chemical composition of stars and planets. What was missing was an explanation for why these lines should exist. Why did particular atoms absorb photons at particular wavelengths?
Kirchhoff's Laws:
The first clue to the origin of spectral lines was found by Kirchhoff who, in the mid-1800's, developed the three laws of spectroscopic analysis. Kirchhoff showed that there are three types of spectra emitted by objects:
1) Continuous spectrum - a solid or liquid body radiates an uninterrupted, smooth spectrum (called a Planck curve)
2) Absorption spectrum - a continuous spectrum that passes through a cool gas has specific spectral lines removed (inverse of an emission spectrum)
3) Emission spectrum - a radiating gas produces a spectrum of discrete spectral lines
There remains a problem with the spectra and the Rutherford model of the atom. Rutherford atoms can absorb or emit photons at any wavelength because the orbit of the electron can take on any distance from the nucleus. However, the spectral lines indicate that there are preferred orbits for the electrons, that they have certain discrete distances from the nuclei. This `quantization' of electron orbits leads to the next great breakthrough in modern physics, quantum mechanics.
Planck's constant:
The UV catastrophe and the dilemma of spectral lines were already serious problems for attempts to understand how light and matter interact. Planck also noticed another fatal flaw in our physics by demonstrating that the electron in orbit around the nucleus accelerates. Acceleration means a changing electric field (the electron has charge), when means photons should be emitted. But, then the electron would lose energy and fall into the nucleus. Therefore, atoms shouldn't exist!
To resolve this problem, Planck made a wild assumption that energy, at the sub-atomic level, can only be transfered in small units, called quanta. Due to his insight, we call this unit Planck's constant (h). The word quantum derives from quantity and refers to a small packet of action or process, the smallest unit of either that can be associated with a single event in the microscopic world.
Changes of energy, such as the transition of an electron from one orbit to another around the nucleus of an atom, is done in discrete quanta. Quanta are not divisible. The term quantum leap refers to the abrupt movement from one discrete energy level to another, with no smooth transition. There is no ``inbetween''.
The quantization, or ``jumpiness'' of action as depicted in quantum physics differs sharply from classical physics which represented motion as smooth, continuous change. Quantization limits the energy to be transfered to photons and resolves the UV catastrophe problem.
Bohr Atom:
Perhaps the foremost scientists of the 20th century was Niels Bohr, the first to apply Planck's quantum idea to problems in atomic physics. In the early 1900's, Bohr proposed a quantum mechanical description of the atom to replace the early model of Rutherford.
The Bohr model basically assigned discrete orbits for the electron, multiples of Planck's constant, rather than allowing a continuum of energies as allowed by classical physics.
The power in the Bohr model was its ability to predict the spectra of light emitted by atoms. In particular, its ability to explain the spectral lines of atoms as the absorption and emission of photons by the electrons in quantized orbits.
Our current understanding of atomic structure was formalized by Heisenberg and Schroedinger in the mid-1920's where the discreteness of the allowed energy states emerges from more general aspects, rather than imposed as in Bohr's model. The Heisenberg/Schroedinger quantum mechanics have consistent fundamental principles, such as the wave character of matter and the incorporation of the uncertainty principle.
In principle, all of atomic and molecular physics, including the structure of atoms and their dynamics, the periodic table of elements and their chemical behavior, as well as the spectroscopic, electrical, and other physical properties of atoms and molecules, can be accounted for by quantum mechanics => fundamental science.
Wave-Particle Dualism:
The wave-like nature of light explains most of its properties:
- reflection/refraction
- diffraction/interference
- Doppler effect
This dualism to the nature of light is best demonstrated by the photoelectric effect, where a weak UV light produces a current flow (releases electrons) but a strong red light does not release electrons no matter how intense the red light.
Einstein explained that light exists in a particle-like state as packets of energy (quanta) called photons. The photoelectric effect occurs because the packets of energy carried by each individual red photons are too weak to knock the electrons off the atoms no matter how many red photons you beamed onto the cathode. But the individual UV photons were each strong enough to release the electron and cause a current flow.
It is one of the strange, but fundamental, concepts in modern physics that light has both a wave and particle state (but not at the same time), called wave-particle dualism.
de Broglie Matter Waves:
Perhaps one of the key questions when Bohr offered his quantized orbits as an explanation to the UV catastrophe and spectral lines is, why does an electron follow quantized orbits? The response to this question arrived from the Ph.D. thesis of Louis de Broglie in 1923. de Broglie argued that since light can display wave and particle properties, then perhaps matter can also be a particle and a wave too.
One way of thinking of a matter wave (or a photon) is to think of a wave packet. Normal waves look with this:
having no beginning and no end. A composition of several waves of different wavelength can produce a wave packet that looks like this:
So a photon, or a free moving electron, can be thought of as a wave packet, having both wave-like properties and also the single position and size we associate with a particle. There are some slight problems, such as the wave packet doesn't really stop at a finite distance from its peak, it also goes on for every and every. Does this mean an electron exists at all places in its trajectory?
de Broglie also produced a simple formula that the wavelength of a matter particle is related to the momentum of the particle. So energy is also connected to the wave property of matter.
Lastly, the wave nature of the electron makes for an elegant explanation to quantized orbits around the atom. Consider what a wave looks like around an orbit, as shown below.
The electron matter wave is both finite and unbounded (remember the 1st lecture on math). But only certain wavelengths will `fit' into an orbit. If the wavelength is longer or shorter, then the ends do not connect. Thus, de Broglie explains the Bohr atom in that on certain orbits can exist to match the natural wavelength of the electron. If an electron is in some sense a wave, then in order to fit into an orbit around a nucleus, the size of the orbit must correspond to a whole number of wavelengths.
Notice also that this means the electron does not exist at one single spot in its orbit, it has a wave nature and exists at all places in the allowed orbit. Thus, a physicist speaks of allowed orbits and allowed transitions to produce particular photons (that make up the fingerprint pattern of spectral lines). And the Bohr atom really looks like the following diagram:
While de Broglie waves were difficult to accept after centuries of thinking of particles are solid things with definite size and positions, electron waves were confirmed in the laboratory by running electron beams through slits and demonstrating that interference patterns formed.
How does the de Broglie idea fit into the macroscopic world? The length of the wave diminishes in proportion to the momentum of the object. So the greater the mass of the object involved, the shorter the waves. The wavelength of a person, for example, is only one millionth of a centimeter, much to short to be measured. This is why people don't `tunnel' through chairs when they sit down.
Elementary Particles :
Modern physics speaks of fundamental building blocks of Nature, where fundamental takes on a reductionist meaning of simple and structureless. Many of the particles we have discussed so far appear simple in their properties. All electrons have the exact same characteristics (mass, charge, etc.), so is an electron fundamental because they are all non-unique.
The search for the origin of matter means the understanding of elementary particles. The understanding of elementary particles requires an understanding of not only their characteristics, but how they interact and relate to other particles and forces of Nature, the field of physics called particle physics.
The study of particles begins with the search for the primary constituent. More than 200 subatomic particles have been discovered so far, however most are not fundamental, most are composed of other, simpler particles. For example, Rutherford showed that the atom was composed of a nucleus and orbiting electrons. Later physicists showed that the nucleus was composed of neutrons and protons. More recent work has shown that protons and neutrons are composed of quarks.
Quarks and Leptons:
The two most fundamental types of particles are quarks and leptons, where each class is divided into 6 flavors corresponding to three generations of matter. Quarks (and antiquarks) have electric charges in units of 1/3 or 2/3's. Leptons have charges in units of 1 or 0.
Baryons and Mesons :
Quarks combine to form other types of matter, baryons and mesons. Baryons are made of three quarks and are found everywhere, they are the protons and neutrons of atomic nuclei (and also anti-protons and anti-neutrons). Mesons, made of quark pairs, are usually found in cosmic rays. Notice that a rule for quark combinations is that the sum of the charges must be an integer; -1, 0, or +1. We do not see fractional charges in Nature, thus there are no free quarks, they are only found in pairs or triplets.
No comments:
Post a Comment