Inorganic Chemistry
THE DISCOVERY OF ATOMIC PARTICLES
The 3 fundamental atomic particles
are protons, neutrons and electrons.
Democritus (an ancient Greek philosopher) taught that if any object was repeatedly cut into smaller and smaller pieces, eventually a smallest particle would be obtained that could not be further divided. He called this smallest particle of matter an ‘atom’. (Gr. ‘a’ = not, ‘tom’ = cuttable).
John Dalton (1807), a British schoolteacher, pictured atoms as solid billiard ball-like spheres. He measured the masses of elements that reacted to form various compounds and proposed his ‘atomic hypothesis’:
- an element is composed of only 1 kind of atom, e.g., the element carbon contains only carbon atoms.
- Atoms of different elements have unique (different) masses, e.g., a carbon atom has a mass of 12 atomic mass units (amu), a hydrogen atom has a mass of 1 amu.
- Chemical compounds are formed from specific ratios of different elements, e.g., H2O always forms in the ratio of 2 hydrogen atoms per oxygen atom.
- Atoms are exchanged (not created or destroyed) in chemical reactions, e.g.,
H2O + Na NaOH + H2
We know this as the law of conservation of mass. Thus, chemists always balance equations.
Balance the preceding equation.
Electrons: In 1897, the first subatomic particle (the electron) was discovered by a British physicist, J. J. Thomson using cathode ray tubes. Two electrodes are sealed in a glass tube containing gas at a low pressure. When a high voltage is applied across the electrodes, current flows as a visible stream of electrons are emitted from the negative electrode (cathode) to the positive electrode (anode). Thomson found that the rays were the same regardless of the metal used for the cathode and he correctly concluded that the particles were part of the makeup of all atoms. Thomson found that cathode rays were deflected by nearby electric and magnetic fields.
Thus electrons have electric charge. In fact, the charge on an electron is the smallest unit of electric charge that can exist (1.602 10-19 Coulombs). Although electrons have a negative charge, atoms have an overall charge of zero. Therefore scientists around 1900 knew that each atom must contain enough positive charge to cancel out the negative charge. Thomson proposed a ‘plum pudding’ model of the atom, in which the positive charge was distributed evenly throughout the atom and the negative charges were pictured as being imbedded in the atom like plums in a pudding.
Atomic Nucleus: By 1909, Ernest Rutherford had determined that alpha () particles (helium nuclei, He+2) are positively charged particles and are emitted by some radioactive atoms – atoms that spontaneously disintegrate. Rutherford bombarded a thin gold foil with particles from a radioactive source. A fluorescent, ZnS screen was placed around the foil to observe the scattering of the particles by the gold atoms. Scintillations (flashes) on the screen caused by the impact of individual particles were counted to determine the relative number of particles deflected at various angles of deflection.
As expected, most particles passed through the foil with little or no deflection, however, to his amazement, a few were deflected at large angles and a few particles bounced straight back at the source.
Rutherford proposed that the positive charge in atoms is not evenly distributed but exists as dense, point-like centers surrounded by a large volume of empty space.
Rutherford named these centers of positive charge – ‘atomic nuclei’. He was able to calculate the magnitude of the positive charges and estimate the diameter of the nucleus at ca. 1/100,000 of an atom.
Since the atom is mostly empty space, the atomic nucleus, containing virtually all the mass, is extremely dense. In fact, an atomic nucleus the size of a grain of sand would weigh ca. 50 106 tons!!
We now know that every nucleus contains an integral (whole) number of protons equal to the number of electrons in the atom (atoms are electrically neutral). The number of protons in an atom (called the atomic number, symbol (Z) determines an atoms identity, e.g., all atoms with 3 protons are lithium atoms.
Neutrons: The 3rd kind of fundamental particle was discovered by James Chadwick in 1932. He bombarded beryllium atoms with high-energy particles and dislodged uncharged particles (neutrons) from the nucleus. It was soon after understood that the nuclei of all atoms (except the common form of hydrogen) contain 1 or more neutrons. Neutrons are almost identical in size and mass with protons. Both neutrons and protons are collectively termed ‘nucleons’ since they both reside in the nucleus of the atom.
Identifying the Elements: H.G.J. Moseley directed high-energy electrons at samples of pure elements. Electrons decelerate rapidly on impact and in so doing emit x-rays. The x-rays emitted are recorded photographically as a series of lines – their patterns varying with the atomic mass of the element. On the basis of mathematical analysis of these x-ray data, it was concluded that each element (H, He, Li, Be, etc.) differs from the preceding element by having one more positive charge in its nucleus. For the first time it was possible to arrange all know elements in order of increasing nuclear charge.
Properties of Subatomic Particles
Particle
/ Symbol / Charge* / Mass (g) / Mass (amu)**electron / e / -1 / 9.109 10-28 / 0.00055
proton / p / +1 / 1.673 10-24 / 1.0078
neutron / n / 0 / 1.675 10-24 / 1.0090
* charges are given as multiples of the charge on an electron (1.602 1019 Coulombs)
** amu (atomic mass unit) = 1/12 of the mass of carbon 12 (12C), the most common form of C.
The Mass Spectrometer and Isotopes: The mass spectrometer (or mass spec) is one of the most powerful analytical instruments available. It permits chemists to identify and quantify (measure the concentration of) all known elements and almost all known compounds.
A portion of the element to be analyzed is injected into a heated sample chamber. Its vapors are drawn into the evacuated instrument and bombarded with high-energy electrons from a cathode ray. The colliding electrons dislodge electrons from the sample atoms producing positive ions of the element. These electrically charged atoms (positive ions) are accelerated through the instrument by a strong electric field applied between 2 metal grids. The ions’ speeds vary with their masses, lighter ions reaching higher speeds. The path of the ions is bent as they travel between the poles of a variable electromagnet. As the magnetic field is varied, each type of positive ion is, in turn, directed to a detector which produces electric signals integrated into a ‘mass spectrum’ – a plot of concentration (signal intensity) versus atomic (positive ion) mass.
Using mass specs, the atomic mass of all 112 known elements have been measured with great accuracy. The use of early mass specs led to the discovery of isotopes. Researchers found that not all atoms of a single element have the same mass. For example, all atoms of boron have 5 protons, however, in a sample of pure boron, 20.0% of the atoms have 5 neutrons while the remaining 80.0% have 6 neutrons. Atoms of the same atomic number (same number of protons, hence same element) but different number of neutrons (hence different mass) are called isotopes.
Naturally Occurring Isotopic Abundance of Some Elements
Element / Nuclide Symbol of Isotope / Number of Neutrons / % Natural Abundance / Atomic Mass(amu) / Weighted Average Mass (amu)
Boron / / 5 / 20.0 / 10.01294 / 10.81
/ 6 / 80.0 / 11.00931
Carbon / / 98.89 / 12.0000
/ 1.11 / 13.0034
Silicon / / 92.21 / 27.9769
/ 4.70 / 28.9765
/ 3.09 / 29.9738
Protium / / 99.98 / 1.0078
Deuterium / / 0.02 / 2.0141
Tritium / / trace / 3.0
Notation of Nuclide Symbols:
* On the periodic table, the symbol A refers to the atomic mass, i.e., the weight average atomic weight of the element in its naturally occurring form – as a mixture of isotopes. Atomic mass is not an integral (whole) number whereas mass number is a count of the number of protons + neutrons and is always an integral number
Problem: Complete the empty cells in the table of Isotopic Abundance.
Problem: Write the nuclide symbol and state the number of protons, neutrons, and electrons:
- an atom of nitrogen 14
- an atom of iron 56
- an atom of uranium 236
Problem: The weight average mass of gallium is 69.72 amu. The masses of the naturally occurring isotopes are 68.9257 for 69Ga and 70.9249 for 71Ga. Calculate the % abundance of each isotope. (Answer: 69 Ga = 60.0%, 71Ga = 40.0%)
Atomic Weight Scale and Atomic Weights: Even before the masses of various kinds of atoms could be measured (as with the mass spec) scientists determined a relative scale of atomic masses for many of the elements. For example, experiments showed that carbon and hydrogen have relative atomic masses (atomic weights) of 12 to 1, respectively.
The atomic weight scale approved in 1962 by the International Union of Pure and Applied Chemists (IUPAC) is based on the carbon-12 isotope.
This is approximately the mass of one atom of protium (1H), the lightest isotope of the element with lowest mass.
Problem: Calculate the mass of 1 amu in grams. Recall Avogadro’s Number, N = 6.022 1023 atoms per mole. (Answer = 1.66110-24g)
Quantum Mechanics: The Rutherford model of the atom, while basically correct, did not answer important questions such as the following.
- Why do different elements have different physical and chemical properties?
- Why and how does chemical bonding occur?
- Why do atoms of different elements give off or absorb light of characteristic colors?
Early scientists found that classical mechanics (Newton’s laws) which successfully describe the motion of visible objects like balls and planets, fail when applied to electrons in atoms. New laws, which came to be known as quantum mechanics were developed in the early 1900’s.
Electromagnetic Radiation: All of the previous unanswered questions (above) can be explained with an understanding of electron arrangement (configuration) within atoms. What we currently know about electron configuration has largely been determined from the analysis of electromagnetic radiation emitted or absorbed by substances (spectroscopy)
Electromagnetic radiations are forms of radiant energy, some natural and some synthetic, that possess no mass or weight and are electrically neutral. They also share 4 other common characteristics:
1)all pass through a vacuum in wavelike motion;
2)all travel at the speed of light (3.00 108 m/s), denoted ‘c’
3)all give off electric and magnetic fields
4)all have different energies, wavelengths, and frequencies
Problem: Calculate the speed of light in miles per hour. 1 mile = 1.609 km. (Ans. = 6.71108mi/h)
Problem: How many minutes will it take light from the sun to reach earth assuming an average distance of 93 106 miles?(Ans. = 8.31min.)
The spectrum of electromagnetic radiation, in order of increasing energy (decreasing wavelength), includes radio and TV waves, microwaves, infrared radiation, visible light, ultraviolet radiation, X-rays, and gamma rays. No clear-cut separation exists between the bands so overlap of wavelengths, as shown below, is reported in various literature sources.
The Electromagnetic Spectrum
Wavelength (m)
10-14 10-12 10-10 10-8 10-6 10-4 10-2 1 102 104
Gamma raysVisible light Microwaves AM radio
Ultraviolet
X-rays Infrared (heat)TV and FM radio
1022 1020 1018 1016 1014 1012 1010 108 106 104
Frequency (Hz)
As far as we know, there is neither and upper nor a lower limit to the wavelength of EMR.
Since all types of electromagnetic radiation (EMR) travel as waves, they can be described in terms of their frequency (, Gr. ‘nu’) and wavelength (, Gr. ‘lambda’).
Wavelength is the distance between any 2 identical points of a wave, for instance, 2 adjacent crests. The frequency is the number of wave crests passing a given point per unit time, usually expressed in cycles/second (cps or s-1 or Hertz, Hz).
For a wave traveling at some speed, the wavelength and frequency are related to each other by…
Thus and are inversely proportional to each other. A shorter the equals a higher . For water waves it is the surface of the water that oscillates. For a vibrating guitar string it is the string that moves repetitively. EMR consists of regular, repetitive variation in electrical and magnetic fields.
The EMR most familiar to us is visible light. It has wavelengths varying from about
400 m (violet) to 800 m (red).
Problem: Calculate the frequency of:
- violet light (Ans. = 7.51014s-1)
- red light (Ans. = 3.81014Hz)
Quanta and Photons: In addition to behaving as waves, EMR can be described as particles called photons. Max Plank, 1900, discovered that each photon has a fixed amount (a quantum) of energy. The amount of energy possessed by a photon depends on its frequency and wavelength.
The energy of a photon of light (e) is given by Plank’s Equation.
h = Plank’s constant = 6.62 10-34Js
= frequency of radiation in Hz
= wavelength of radiation in m
Problem:
- Calculate the energy of 1 photon of violet light ( = 7.31 1014 s-1) [Ans. = 4.8410-19 J]
- Calculate the energy of 1 photon of x-rays with = 2.5 m. [Ans. = 7.9410-17 J]
The Photoelectric Effect: Evidence for the particle nature of light came from the photoelectric effect. When light of sufficiently high energy strikes the negative electrode (cathode) in an evacuated tube, electrons are knocked off the electrode surface and travel to the positive electrode (anode) creating an electric current in the circuit. However, the following behaviors were noted …
- no electrons were ejected unless the incident radiation has a frequency above a certain value characteristic of the metal cathode, no matter how long or how brightly the light shines.
- the electric current (number of electrons emitted per second) increases with increasing brightness (intensity) of the light.
Classical physics said that even low energy light should cause a current to flow if the metal is irradiated long enough. Electrons should accumulate energy and be released when they have enough energy to escape from the metal atoms. This is not observed. In addition, old theory suggests that if light is more energetic, then the current should increase even though the light intensity remains the same. This also is not observed.
The answer to the puzzle was provided by Albert Einstein. In 1905 he extended Plank’s idea that light behaves as though it were composed of photons, each with a particular amount (quantum) of energy. According to Einstein, each photon can transfer its energy to a single electron in a collision. If the energy of the photon is equal to or greater than the amount needed to liberate the electron, then the electron can escape the metal surface. Increased intensity means that the number of photons striking a given area per second is increased.
Atomic Spectra and the Bohr Atom:
Incandescent (red hot or white hot) solids, liquids, and high-pressure gases emit continuous spectra. For example, a white hot (nearly 1000 C) tungsten light bulb filament emits a continuous band of visible radiation (white).
However, when an electric current is passed through a gas in a vacuum tube at very low pressure, the light that the gas emits is dispersed by a prism into distinct lines. Such emission spectra are described as bright line spectra. The lines can be recorded photographically and the wavelength of light that produced each line can be calculated.
Similarly, we can shine a beam of white light (containing a continuous spectrum) through a gas and analyze the beam that emerges. We find that only certain wavelengths have been absorbed. The wavelengths that are absorbed in this absorption spectrum are the same as those given off in the emission experiment.
Each element displays its own characteristic set of lines in its emission or absorption spectrum. These spectra can serve as 'fingerprints' to allow us to identify different elements in a sample, even in trace amounts.
Emission spectra of various elements were intensely studied by scientists. J.J. Rydberg, a British school teacher discovered that the wavelengths of the hydrogen spectrum can be related by a mathematical equation:
The Rydberg equation: where R = 1.097 107 m-1 (the Rydberg constant)
In 1913, Neils Bohr, a Danish physicist, provided an explanation for Rydberg's observations. He wrote equations that described the electron of a hydrogen atom as revolving around the nucleus of the atom in circular orbits (planetary model of the atom). He included assumptions that the electronic energy is quantized; that is, only certain values of electron energy are possible. This led him to the suggestion that electrons can only be in certain discrete orbits, and that they absorb or emit energy in discrete amounts as they move from one orbit to another. Each orbit thus corresponds to a definite energy level for the electron. When an electron is promoted from a lower energy level to a higher one, it absorbs a definite (or quantized) amount of energy. When the electron falls back to the original energy level, it emits exactly the same amount of energy it absorbed in moving from the lower to the higher energy level.
Bohr's equation for the energy of each orbit was:
where
The larger the value of n, the farther from the nucleus is the orbit being described. For orbits farther from the nucleus, the electronic potential energy is higher (less negative - the electron is in a higher energy level or less stable state). As n approaches infinity, the electron is completely removed from the nucleus.
With this equation, Bohr was able to predict the wavelengths observed in the hydrogen emission spectrum. Although the Bohr theory explained the spectra of hydrogen and other species containing only one electron (He+, Li+2) it could not calculate the wavelengths observed in spectra of more complex species. Bohr's approach was doomed to failure because it modified classical mechanics. It was a contrived solution. There was a need to literally invent a new physics, quantum mechanics, to deal with subatomic particles.
However, Bohr's theory did support the ideas that only certain energy levels are possible and that energy levels could be described by quantum numbers.
Wave Particle Duality of Matter: Once it was learned that EMR can exhibit both wave properties and particle properties, French scientist, Louie de Broglie (1925) suggested that all particles have wavelength properties. de Broglie predicted that a particle with a mass m and velocity v should have a wavelength given by …