Topic 2: Atomic Theory 2.1 The atom Introduction to atomic structure - Atoms, molecules, elements, compounds and mixtures Atoms are the basic building blocks of matter; they are the smallest units of an element. A molecule is a larger unit of matter in which two or more atoms are joined together. An element is composed of only one kind of atom. In compounds the atoms of two or more elements combine in definite arrangements. Mixtures do not involve the specific interactions between elements found in compounds, and the substances which comprise the mixture can be of varying ratios. Dalton s Atomic Theory, 1803 According to the first two postulates of Dalton s Atomic theory: All matter consists of tiny particles (atoms) Atoms are indestructible and unchangeable J.J. Thomson At the end of the nineteenth century, a scientist called J.J. Thomson discovered the electron. This is a tiny charged particle that is much, much smaller than any atom. Thomson was experimenting by applying high voltages to gases at low pressure. Since they were so small, Thomson suggested that they could only have come from inside atoms. Dalton's idea of the indestructible atom had to be revised atoms are indestructible and unchangeable. Subatomic Particles Since an atom is electrically neutral you would also expect there to be positively charged particles within an atom Thomson thought that the electrons were spread in a cloud of positive charge - the Plum Pudding model of the atom. Rutherford s model for the atom In 1911, the scientist Ernest Rutherford suggested a new model for the atom. In his model, the positive charge was concentrated in a tiny volume at the center of the atom. Subatomic particles We now know that the positive charge is also in the form of a particle and that there are at least three types of particle inside an atom: Electrons (negatively charged), Protons (positively charged), Neutrons (electrically neutral). The Neutral Atom Since, under normal conditions, the atom as a whole is electrically neutral, the number of electrons must equal the number of protons in the nucleus. 2.1.1. State the position of protons, neutrons and electrons in the atom. The structure of an atom Experiments have shown that both the protons and neutrons are crammed into the centre of the atom, called the nucleus. This is where the mass of an atom is concentrated. Mr. Johns 15/9/08 1
The electrons are found orbiting (circling) the nucleus. 2.1.2 State the relative masses and relative charges of protons, neutrons and electrons. The Mass of Subatomic Particles Now let's consider the masses of the sub-atomic particles. The masses are so small it is easier to compare them with each other. The relative masses of the three subatomic particles Sub-atomic particle Relative mass Proton 1 Neutron 1 Electron 1/1840 (so small we can usually ignore their mass) 2.1.3 Define the terms mass number (A), atomic number (Z) and isotopes of an element. The Mass of an Atom Electrons contribute very little to the mass of an atom. The mass of an atom can be determined by simply adding up the number of protons and neutrons. The number of protons and neutrons is known at the mass number of an atom. Mass number (A) The mass number of an atom tells us how many heavy particles protons and neutrons there are in the nucleus of an atom. Mass number = number of protons + number of neutrons in an atom. Nucleon Number? Some people refer to the mass number as the nucleon number because nucleon means a proton, or a neutron. Superscript We show the mass number by writing it as a superscript before the symbol for the element as shown below: 1 H 4 He 7 Li 9 Be 11 B etc. Atomic number (Z) Each chemical element also has its own unique atomic number. The atomic number tells us the number of protons in an atom, hence it is sometimes referred to as the proton number. Atomic Number = No. of Electrons All atoms are neutral - they carry no overall charge. therefore, the atomic number of an element also tells us the number of electrons in an atom. Subscript We show the atomic number by writing it as a subscript before the symbol for the element as shown below: 1H 2He 3Li 4Be 5B etc. Mr. Johns 15/9/08 2
Note: The atomic number is written above the symbol for an element in the periodic table. Usefulness of mass no. and atomic number Taking lithium (Li) as an example, we know from its mass number that its nucleus must contain a total of seven protons and neutrons, but how many of each? To answer this, we need the atomic number as well. 7 Li The element lithium has a mass number of 7 and an atomic number of 3 given that the atomic number of Li is 3, we know that Li has three protons therefore, it must have four neutrons, to make up its mass number of 7. 7 Li 3Li 7 3 Li nucleus containing three protons and 7-3 = 4 neutrons Relative atomic masses and isotopes The mass of atoms is compared on a relative scale. This is more convenient than quoting the tiny masses of atoms in standard units of mass, like grams or kilograms. Calculating relative atomic masses On the relative atomic mass scale, a carbon-12 atom has a mass of exactly 12.00 The relative atomic mass of hydrogen-1 is 1, making it the lightest of all atoms. Calculating relative atomic masses If you look up the mass of elements in the periodic table, you will find that they are not a whole number e.g. chlorine actually has a relative atomic mass of 35.5 At first sight this seems a bit strange! how can you have an atomic mass that is a fraction? after all, you can't get half a proton or half a neutron in an atom! The answer lies in the existence of isotopes. Isotopes What is an isotope? Each chemical element has its own unique atomic number (the number of protons in an atom). That means that all the atoms of any particular element must have the same number of protons (and therefore electrons). so, any atom atom of hydrogen has one proton, otherwise it wouldn't be hydrogen! if the atom had two protons, it would be helium. Number of Neutrons But there is no reason why the number of neutrons can't vary in fact, there are three types of hydrogen atom, each containing a different number of neutrons We call these atoms isotopes of hydrogen. Isotopes Isotopes are atoms of the same element that have different numbers of neutrons. Different isotopes of an element have the same atomic number, but different mass numbers. Mr. Johns 15/9/08 3
2.1.4 Deduce the symbol for an isotope given its mass number and atomic number. Use the notation A ZX, eg 12 6C. Isotopes of Hydrogen The isotopes of hydrogen are: Hydrogen Deuterium Tritium (no neutrons) (1 neutron) (2 neutrons) The two heavier isotopes of hydrogen may be written as: 2 1H (deuterium) and 3 1H (tritium) Isotopes of Hydrogen Isotopes are also sometimes shown by the name of the element, followed by the mass number. thus the three isotopes of hydrogen would be: Hydrogen-1, Hydrogen-2, Hydrogen-3 How many isotopes can one element have? There are preferred combinations of neutrons and protons, at which the forces holding nuclei together seem to balance best. Light elements tend to have about as many neutrons as protons Heavy elements apparently need more neutrons than protons in order to stick together. Atoms with a few too many neutrons, or not quite enough, can sometimes exist for a while, but they're unstable. 2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge. Calculating relative atomic masses Chlorine exists as two isotopes: chlorine-35 chlorine-37 When we calculate the relative atomic mass of an element we have to take into account the proportions of each isotope present. In a naturally occurring sample of chlorine, we find that 75 per cent is chlorine-35 atoms and the other 25 per cent is chlorine-37 atoms. Calculating relative atomic masses If you had 100 atoms of chlorine: 75 would have a relative atomic mass of 35 and 25 would have a relative atomic mass of 37. 75 x 35 = 2625 25 x 37 = 925 3550/100 = 35.5 Mr. Johns 15/9/08 4
2.1.6 Compare the properties of the isotopes of elements. Chemical Properties of Isotopes Chemical properties of an element are related to the number and arrangement of electrons. Because isotopes have the same number, and therefore arrangement, of electrons, isotopes of an element undergo the same chemical reactions. The only difference is their mass. Physical Properties of Isotopes Isotopes have similar, but not identical physical properties e.g. the density of the isotope and its compounds will be slightly different: -this may allow separation of the isotopes e.g. Uranium 238 is separated from Uranium 235 Uranium is first converted to uranium hexafluoride (a gas) and then separated by centrifugation into separate isotopes. 2.1.7 Discuss the uses of radioisotopes http://www.uic.com.au/peac.htm Some isotopes are referred to as 'stable' and others as 'unstable' or 'radioactive'. It is the radioactive nature of these unstable isotopes, usually referred to as 'radioisotopes', which gives them so many applications in modern science and technology. Agriculture Fertilisers 'labelled' with a particular isotope, such as nitrogen-15 and phosphorus-32 provide a means of finding out how much is taken up by the plant and how much is lost, allowing better management of fertiliser application. Radio carbon dating Radio carbon dating is a technique that is used to date things that are composed of organic material. Radiocarbon dating was recognised as a way to date articles in 1960 and W.F. Libby was awarded the Nobel Prize in Chemistry for this finding. One isotope of carbon is radioactive and unstable whilst the others are stable. Carbon -14 is the radioactive isotope of Carbon and is the isotope that is relied upon in radiocarbon dating. Carbon-14, 14 6C, has a half life of 5730 years and decays as follows: 14 C 14 N + ß - The ratio of 14 C: 12 C are constant in a living plant. Carbon 14 in the environment decays, but is reformed at the same rate as it is lost by the following equation: 14 N + 1 n 14 C + 1 H Where 1 n is a high-energy neutron In living plants, the process of photosynthesis ensures that the amount of Carbon-14 used as CO 2 is continuous along with the amount of Carbon-12 and Carbon-13. Once the plant has died however, no more carbon-14 is taken in and the Carbon-14 that is present starts to decay. As the carbon-14 decays, the ratio of 14 C: 12 C changes with time. This change in the ratio allows the species to be dated with respect to the amount of Carbon-14 present. Mr. Johns 15/9/08 5
Radiotracers Radiotracers depend on the fact that an isotope of the same element behaves with very little difference in chemical behavior. A radioisotope inserted into the body will take the route of the element that is normally present. The radioisotope can be traced whilst traveling through the body and any problems can be detected. For example, 131 I and 125 I may be used to evaluate iodine uptake by the thyroid gland. Radiotherapy The uses of radioisotopes in therapy are comparatively few, but important. Cancerous growths are sensitive to damage by radiation, which may be external- using a gamma beam from a cobalt-60 source, or internal - using a small gamma or beta radiation source. Iodine-131 is commonly used to treat thyroid cancer, probably the most successful kind of cancer treatment, and also for non-malignant thyroid disorders. Now read Green & Damji 2.2 The mass spectrometer Textbook http://www.cem.msu.edu/~reusch/virtualtext/spectrpy/massspec/masspec1.htm 2.2.1 Describe and explain the operation of a mass spectrometer A simple diagram of a single beam spectrometer is required. The following stages of operation should be considered: vaporization, ionisation, acceleration, deflection and detection. In order to measure the characteristics of individual molecules, a mass spectrometer converts them to ions so that they can be moved about and manipulated by external electric and magnetic fields. The three essential functions of a mass spectrometer, and the associated components, are: 1. A small sample of compound is ionized, usually to cations by loss of an electron - the Ion Source. 2. The ions are sorted and separated according to their mass and charge - the Mass Analyzer. 3. The separated ions are then detected and tallied, and the results are displayed on a chart - the Detector. Because ions are very reactive and short-lived, their formation and manipulation must be conducted in a vacuum. Atmospheric pressure is around 760 torr (mm of mercury). The pressure under which ions may be handled is roughly 10-5 to 10-8 torr (less than a billionth of an atmosphere). Each of the three tasks listed above may be accomplished in different ways. In one common procedure, ionization is effected by a high energy beam of electrons, and ion separation is achieved by accelerating and focusing the ions in a beam, which is then bent by an external magnetic field. The ions are then detected electronically and the resulting information is stored and analyzed in a computer. A mass spectrometer operating in this fashion is outlined in the diagram on the next page. Mr. Johns 15/9/08 6
The heart of the spectrometer is the ion source. Here molecules of the sample (black dots) are bombarded by electrons (light blue lines) issuing from a heated filament. This is called an EI (electron-impact) source. Gases and volatile liquid samples are allowed to leak into the ion source from a reservoir (as shown), but nonvolatile solids and liquids may be introduced directly. Cations formed by the electron bombardment are pushed away by a charged repellor plate (anions are attracted to it), and accelerated toward other electrodes, having slits through which the ions pass as a beam. Some of these ions fragment into smaller cations and neutral fragments. When the ion beam experiences a strong magnetic field perpendicular to its direction of motion, the ions are deflected in an arc whose radius is inversely proportional to the mass of the ion. Lighter ions are deflected more than heavier ions. By varying the strength of the magnetic field, ions of different mass can be focused progressively on a detector fixed at the end of a curved tube (also under a high vacuum). Molecular and Fragment Ions When a high energy electron collides with a molecule it often ionizes it by knocking away one of the molecular electrons (either bonding or non-bonding). This leaves behind a molecular ion. Residual energy from the collision may cause the molecular ion to fragment into neutral pieces and smaller fragment ions. Mr. Johns 15/9/08 7
The Nature of Mass Spectra: Molecular Fingerprints A mass spectrum is usually presented as a vertical bar graph, in which each bar represents an ion having a specific mass-to-charge ratio (m/z), the length of each bar indicates the relative abundance of the ion. The most intense ion is assigned an abundance of 100, and it is referred to as the base peak. Most of the ions formed in a mass spectrometer have a single charge, so the m/z value is equivalent to mass itself. Modern mass spectrometers easily distinguish (resolve) ions differing by only a single atomic mass unit (amu), and thus provide completely accurate values for the molar mass of a compound. The highest-mass ion in a spectrum is normally the molecular ion, and lower-mass ions are fragments from the molecular ion, assuming that the sample is a pure compound. The following diagrams are the mass spectra of three simple gaseous compounds, carbon dioxide, propane and cyclopropane: The molecules of these compounds are similar in size, CO 2 and C 3 H 8 both have a mass of 44 amu, and C 3 H 6 has a mass of 42 amu. The molecular ion is the strongest ion in the spectra of CO 2 and cyclopropane (C 3 H 6 ), and is moderately strong in propane. The mass resolution is readily apparent in these spectra (note ions having m/z=39, 40, 41 and 42 in the cyclopropane spectrum). Even though these compounds are very similar in size, it is a simple matter to identify them from their individual mass spectra. Even with simple compounds like these, it should be noted that it is rarely possible to explain the origin of all the fragment ions in a spectrum. Since a molecule of carbon dioxide is composed of only three atoms, its mass spectrum is very simple. The molecular ion is also the base peak, and the only fragment ions are CO (m/z=28) and O (m/z=16). The molecular ion of propane also has m/z=44, but it is not the most abundant ion in the spectrum. Cleavage of a carbon-carbon bond gives methyl and ethyl fragments, but the larger ethyl cation (m/z=29) is the most abundant. A similar bond cleavage in cyclopropane does not give two fragments, so the molecular ion is stronger than in propane, and is responsible for the base peak. Loss of a hydrogen atom, either before or after ring opening, produces the stable allyl cation (m/z=41). The third strongest ion in the spectrum has m/z=39 (C 3 H 3 ). The small m/z=39 ion in propane and the absence of a m/z=29 ion in cyclopropane are particularly significant in distinguishing these hydrocarbons. Mr. Johns 15/9/08 8
Most stable organic compounds have an even number of total electrons, reflecting the fact that electrons occupy orbitals in pairs. When a single electron is removed from a molecule to give an ion, the total electron count becomes an odd number, and we refer to such ions as radical cations. The molecular ion in a mass spectrum is always a radical cation, but the fragment ions may either be even or odd-electron cations, depending on the neutral fragment lost. Fragment ions themselves may fragment further. Isotopes: Atoms of Differing Mass Since a mass spectrometer separates and detects ions of slightly different masses, it easily distinguishes different isotopes of a given element. This is most dramatic for compounds containing bromine and chlorine. Since molecules of bromine have only two atoms, the spectrum on the left will come as a surprise if a single atomic mass of 80 amu is assumed for Br. The five peaks in this spectrum demonstrate clearly that natural bromine consists of a nearly 50:50 mixture of isotopes having atomic masses of 79 and 81 amu respectively. Thus, the bromine molecule may be composed of two 79 Br atoms (mass 158 amu), two 81 Br atoms (mass 162 amu) or the more probable combination of 79 Br- 81 Br (mass 160 amu). Fragmentation of Br 2 to a bromine cation then gives rise to equal sized ion peaks at 79 and 81 amu. <> bromine vinyl chloride methylene chloride The center and right hand spectra show that chlorine is also composed of two isotopes, the more abundant having a mass of 35 amu, and the minor isotope a mass 37 amu. The precise isotopic composition of chlorine and bromine is: Chlorine: 75.77% 35 Cl and 24.23% 37 Cl Bromine: 50.50% 79 Br and 49.50% 81 Br The presence of chlorine or bromine in a molecule or ion is easily detected by noticing the intensity ratios of ions differing by 2 amu. In the case of methylene chloride, the molecular ion consists of three peaks at m/z=84, 86 and 88 amu, and their diminishing intensities may be calculated from the natural abundances given above. Loss of a chlorine atom gives two isotopic fragment ions at m/z=49 & 51amu, clearly incorporating a single chlorine atom. Fluorine and iodine, by contrast, are mono-isotopic, having masses of 19 amu and 127 amu respectively. Mr. Johns 15/9/08 9
2.2.2 Describe how the mass spectrometer may be used to determine relative atomic mass using the 12C scale. High Resolution Mass Spectrometry In assigning mass values to atoms and molecules, we have so far assumed integral values for isotopic masses. However, accurate measurements show that this is not strictly true. Because the strong nuclear forces that bind the components of an atomic nucleus together vary, the actual mass of a given isotope deviates from its nominal integer by a small but characteristic amount (remember E = mc 2 ). Thus, relative to 12 C at 12.0000, the isotopic mass of 16 O is 15.9949 amu (not 16) and 14 N is 14.0031 amu (not 14). By designing mass spectrometers that can determine m/z values accurately to four decimal places, it is possible to distinguish different formulas having the same nominal mass. The table below illustrates this important feature, and a double-focusing high-resolution mass spectrometer easily distinguishes ions having these compositions. Mass spectrometry therefore not only provides a specific molecular mass value, but it may also establish the molecular formula of an unknown compound. Formula C 6 H 12 C 5 H 8 O C 4 H 8 N 2 Mass 84.0939 84.0575 84.0688 Tables of precise mass values for any molecule or ion are available in libraries. Since a given nominal mass may correspond to several molecular formulas, lists of such possibilities are especially useful when evaluating the spectrum of an unknown compound. Composition tables are available for this purpose. 2.2.3 Calculate non-integer relative atomic masses and abundance of isotopes from given data. If a sample of magnesium is vaporized in a mass spectrometer, three peaks are found on the mass spectrum. The table below shows the isotopes and their relative abundance. Use the data to determine the relative atomic mass of magnesium. Isotope Abundance 24 Mg 100 25 Mg 12.8 26 Mg 14.4 Total 127.2 Mr. Johns 15/9/08 10
2.3 Electron Arrangement 2.3.1 Describe the electromagnetic spectrum (in terms of variation in wavelength, frequency and energy. Students should be able to identify the ultraviolet, visible and infrared regions and to describe the variation in wavelength, frequency and energy across the spectrum. Electromagnetic Radiation is a type of energy that travels in waves Types of EMR: Radio waves Microwaves Infrared Visible light Ultra Violet X-Rays Gamma Rays EMR exists in a range of wavelengths that extends from radio waves that many be thousands of meters long to gamma rays with wavelengths as short as a million-millionth (10-12) of a meter. Introduction to Electromagnetic Radiation (EMR) EMR is a type of energy that is transmitted through space at enormous velocities Electromagnetic waves are produced by the motion of electrically charged particles. Electromagnetic radiation because they radiate from the electrically charged particles. EMR takes numerous forms (spectrum): visible region is small compared to other spectral regions 3 Mr. Johns 15/9/08 11
Electromagnetic spectrum 4 Energy of a Photon http://imagers.gsfc.nasa.gov/ems/waves4.html Electromagnetic waves can be described by their wavelengths, energy, and frequency. They are related to each other mathematically: E = h ν or, E = h c / λ h = Plank s constant ν = frequency λ = wavelength c = speed of light Energy is inversely proportional to wavelength 9 Planck's constant h = 6.626 x 10-34 Js -1 Photon energies are usually quoted in electron-volts. 1eV = 1.6 x 10-19 J The greater the frequency and smaller the wavelength, the greater the energy of the radiation. Hence, uv radiation is more energetic than visible light. X-rays and gamma-rays are extremely penetrating. 2.3.2 Distinguish between a continuous spectrum and a line spectrum http://www.bcpl.net/~kdrews/bohr/bohr.html http://www.iun.edu/~cpanhd/c101webnotes/modern-atomic-theory Solids, liquids and dense gases glow at high temperatures. The emitted light, examined using a spectroscope, consists of a continuous band of colours as in a rainbow. A continuous spectrum is observed. A continuous spectrum, or rainbow, when observed in nature, is usually the result of the spectra of many elements superimposed on top of each other. In addition, the spectrum produced by sunlight is a continuous spectrum. Mr. Johns 15/9/08 12
This is typical of matter in which the atoms are packed closely together. Gases at low pressure behave quite differently. The excited atoms emit only certain frequencies, and when these are placed as discreet lines along a frequency scale a line spectrum (atomic emission spectrum) is formed. Spectral lines exist in series in the different regions (infra-red, visible and ultra-violet) of the spectrum of electromagnetic radiation. The spectral lines in a series get closer together with increasing frequency. Each element has its own unique atomic emission spectrum. By the early 1900 s it was well known that an atomic (emission) spectrum existed for the various elements. The process of exciting an atom, involves adding energy to the atom. This can be done in a variety of ways: Simply heating a sample of an element in an open flame will excite electrons -the colored lights observed when sky rockets explode are a result of burning gunpowder exciting electrons within atoms of elements packed with the gun powder. Passing electricity through a sample of an element e.g. hydrogen at low pressure, will also excite electrons. The emitted light is passed through a prism - which separates it into its constituent wavelengths. The wavelengths can then recorded on light sensitive film. Each element gives a unique spectrum. The position of the lines changes from element to element and, from molecule to molecule. In essence, the atomic emission spectrum is a fingerprint of the element that generates it. The emission spectrum is related to the locations of the electrons in the atom and their relationship with the nucleus (see the next section). Atomic spectra were fundamental pieces of experimental information used by chemists in the development of the electronic structures of atoms. By studying the colors emitted by different elements, it was possible to work backwards to the sources of those colors. In this way, it was possible to determine the electronic structures of the elements. 2.2.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels. http://www.avogadro.co.uk/light/bohr/spectra.htm It was necessary to explain how electrons are situated in atoms and why atoms are stable. Hydrogen atoms were studied as these contained only one proton and one electron, making them convenient to study. In 1913, Neils Bohr solved many of the problems at the time by proposing that the electron revolves around the nucleus of the atom with a definite fixed energy in a fixed path, without emitting or absorbing energy. The electron in the hydrogen atom exists only in certain definite energy levels. When the electron occupies the energy level of lowest energy the atom is said to be in its ground state. An atom can have only one ground state. If the electron occupies one of the higher energy levels then the atom is in an excited state. An atom has many excited states. Mr. Johns 15/9/08 13
When a gaseous hydrogen atom in its ground state is excited by an input of energy, its electron is 'promoted' from the lowest energy level to one of higher energy. The atom does not remain excited but re-emits energy as electromagnetic radiation. This is as a result of an electron 'falling' from a higher energy level to one of lower energy. This electron transition results in the release of a photon from the atom of an amount of energy (E = hn) equal to the difference in energy of the electronic energy levels involved in the transition. In a sample of gaseous hydrogen where there are many trillions of atoms all of the possible electron transitions from higher to lower energy levels will take place many times. A prism can be used to separate the emitted electromagnetic radiation into its component frequencies (wavelengths or energies). These are then represented as spectral lines along an increasing frequency scale to form an atomic emission spectrum. The Bohr theory was a success in explaining the spectrum of the hydrogen atom. His calculated wavelengths agreed perfectly with the experimentally measured wavelengths of the spectral lines. Bohr knew that he was on to something; matching theory with experimental data is successful science! More recent theories about the electronic structure of atoms have refined these ideas, but Bohr's 'model' is still very helpful. For clarity, it is normal to consider electron transitions from higher energy levels to the same level. Electron transitions from higher energy levels to the n = 3 level produce a series of lines in the infrared region of the electromagnetic spectrum, called the Paschen Series. Electron transitions from higher energy levels to the n = 2 level produce a series of lines in the visible region of the electromagnetic spectrum, called the Balmer Series. The series of lines in the ultra-violet region, called the Lyman Series, are due to electron transitions from higher energy levels to the n = 1 level, and these were discovered after Bohr predicted their existence. Within each series, the spectral lines get closer together with increasing frequency. This suggests that the electronic energy levels get closer the more distant they become from the nucleus of the atom. Convergence is a mathematical term for a series of numbers that gradually decrease, but never quite get to 0, and add up to infinity (but technically reach some limit). The diagram below illustrates the formation of three series of spectral lines in the atomic emission spectrum of hydrogen. Mr. Johns 15/9/08 14
Additional information Electron Transition and Quantum Jumps An electron transition is the movement of an electron from one energy level to another. If an electron moves from a low energy level to a higher energy level it does so by gaining energy. If it moves from a higher energy level to a lower energy level, it does so by releasing energy. The released energy is in the form of electromagnetic radiation. The energy content and wavelength of the released electromagnetic radiation will correspond to the difference in energy content of the two levels. This transition is also known as a Quantum Jump. Energy Level An energy level is a specific location on an energy level diagram that corresponds to an allowed electron energy content. Bohr Theory said that electrons could only exist at specific levels of energy. This component of the theory was based on work performed by Max Planck. (These energy levels are called Principal Quantum Levels, denoted by the Principal Quantum Number, n. Principal Quantum Level n = 1 is closest to the nucleus of the atom and of lowest energy). Excited State The Excited State according to Bohr Theory is a position on an energy level diagram that contains more energy than the Ground State. When an electron is exposed to energy it may absorb some of that energy. If it does, it will rise up the energy level diagram to a new position that corresponds to the higher energy content. An electron in an excited state is not in its most stable position. An electron in an excited state will eventually return to the Ground State. Ground State The Ground State according to Bohr Theory is the position on an energy level diagram that the electron is located in under normal conditions, such as room temperature. When located at Ground State, the electron is in the orbit that is closest to the nucleus and has the least allowed quantity of energy. It is the position of maximum stability for an electron on an atom. An electron that is not located in the Ground State, or position of greatest stability, will be located in the Excited State. Bohr Theory The Bohr model of the atom In 1914 the Danish physicist Niels Bohr revised the model of the atom. Mr. Johns 15/9/08 15
Bohr suggested that the electrons must be orbiting the nucleus in certain fixed energy levels (or shells) The evidence for this came from light energy emitted from heated atoms e.g. hydrogen line spectra. Bohr Theory was the foundation on which many people began to look at the behavior of electrons in new ways. The theory has some clearly defined problems, ideas that we now believe were incorrect. Those problems have been modified over the years so that the final description of the behavior of the electron is now more accurate. Ultimately, the Bohr Theory led scientists, specifically Erwin Schrodinger, to the Modern Theory of Atomic Structure. The primary limitation to Bohr Theory is that it was limited to a description of a one electron system, namely hydrogen. The description of multiple electron systems is much more complex, and was only adequately handled by the Modern Theory of Atomic Structure. Orbit According to Bohr Theory an orbit is the path that an electron follows as it moves around the nucleus. The orbits appear as a series of concentric circles with their centers located at the nucleus. This idea was eventually replaced with the concept of the orbital in the Modern Theory of Atomic Structure. The concept of the orbit should not be confused with the term known as the orbital. This term was introduced in the Modern Theory of Atomic Structure and represents a different idea. 2.3.4 Deduce the electron arrangement for atoms and ions up to Z=20 Energy levels or shells Electrons don't just orbit the nucleus in a random way, they occupy energy levels or shells at different distances from the centre of the atom. There are 7 energy levels around an atom, although they are not always filled. Each energy level or shell can hold only a certain number of electrons. Moving outwards from the nucleus, the shells get larger and can hold more electrons. Electrons always occupy the lowest available energy level. The lowest energy level is close to the nucleus. Mr. Johns 15/9/08 16
Energy Levels The lowest energy level (the one found nearest to the nucleus) can hold just two electrons Some people refer to this as the first or innermost shell. Energy Levels The second energy level can hold eight electrons, as can the third energy level. Summary of electron arrangement Energy Levels The energy levels fill up with electrons from the lowest energy level (innermost shell) and build up outwards They only start occupying a new energy level when the previous one has been filled. Electronic Configuration We can represent the arrangement of electrons in an atom using a shorthand called the electronic structure or electronic configuration This shows the numbers of electrons in each energy level, starting with the lowest level. Electronic Structures So, for atoms of the elements below, we have the following electronic structures: Helium: 2 Carbon: 2, 4 Sodium: 2, 8, 1 Calcium: 2, 8, 8, 2 For ions: Na + = 2, 8 Ca 2+ = 2, 8, 8 Note that there is a relationship between atomic structure and position in the periodic table. Mr. Johns 15/9/08 17