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Slide 2 Picturing the atom Slide 3 Contents The Structure of Atoms Isotopes of elements The mass spectrometer Electron Arrangement The nature of the electron HL Slide 4 Objectives State the position of protons, neutrons and electrons in the atom State the relative masses and charges of protons, neutrons and electrons. Define the terms mass number (Z) and isotopes of an element. Deduce the symbol for an isotope given its mass number and atomic number. Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge. Compare the properties of the isotopes Discuss the uses of radio isotopes Slide 5 A single Hydrogen-1 atom contains one proton and one electron. An electron has negligible mass compared with that of a proton. As 1 mol of hydrogen atoms contains 6.02 x 10 23 hydrogen-1 atoms, and the mass of 1 mol of hydrogen-1 atoms is 1.00g, we can calculate the mass of a single proton. It will simply be equal to the reciprocal of Avogadros constant in grams. Properties of sub atomic particles Slide 6 A neutron has a mass that is almost the same as, but very slightly larger than, that of a proton, whereas an electrons mass is only about 1/1836 that of a proton. These masses can be simplified by assigning relative masses ( known as atomic mass units). Similarly, electrons and protons have relative charges. Can you remember the relative masses and charges of each sub atomic particle? Copy out and complete the table below. Sub atomic particle Relative MassRelative Charge Proton Neutron Electron 1 1 0 or 5 x 10 -4 +1 0 Slide 7 All atoms can be characterized by two numbers: the atomic number and the mass number. The atomic number (Z) is simply equal to the number of protons in the nucleus of the atom. As atoms are electrically neutral this will also be equal to the number of electrons in the atom. Atoms of different elements will have different atomic numbers The mass number (A) is equal to the number if protons and neutrons (collectively known as nucleons) in the nucleus of the atom. Different atoms of the same element may have the same mass number, or a different mass number if they contain a different number of neutrons. Atoms of the same elements must contain the same number of protons, but if they contain a different number of neutrons they are known as ISOTOPES Isotopes of elements Slide 8 Isotopes are atoms of the same element which have the same number of protons but a different number of neutrons Slide 9 The symbol for an isotope of an atom is written in the form: A Z X Examples include: H 1 1 H 2 1 3 H 1 Cl UU 35 17 37 17 235 92 238 92 hydrogendeuteriumtritium Chlorine-35Chlorine-37 Uranium-235Uranium-237 More on isotopes Slide 10 Isotopic ions If the symbol for an ion of an isotope rather than that of an atom is required, then the charge carried by the ion is written on the top right hand side, some examples are shown below: SymbolAtomic number Z Mass number A Number of protons Number of neutrons Number of electrons 27 13 Al 23 11 Na + 31 15 P 3- 1327131413 1123111210 1531 151618 Slide 11 Properties of isotopes The chemical properties of atoms depend on their outer electrons. As all isotopes of the same element have the same arrangement of electrons, their chemical properties are identical. Isotopic Water Boiling point at 1atm H2OH2O100.0 D2OD2O101.4 T2OT2O101.5 H 2 17 O100.1 H 2 18 O100.2 D 2 18 O101.5 HDO100.7 HTO100.8 However because they have different masses, their physical properties such as density, rate of diffusion, melting point and boiling point will differ. Slide 12 Radio isotopes Isotopes have many uses in chemistry and beyond. Many but by no means all, isotopes of elements are radioactive, because the nuclei of these atoms break down spontaneously. When they break down, these radioisotopes emit radiation, which may be one of three types: Gamma () radiation is highly penetrating, Beta () radiation which can be stopped by a thin sheet of aluminium Alpha () radiation which can be stopped by a few centimeters of air. Slide 13 Some uses of radioisotopes Radioisotopes can occur naturally or be created artificially. Their uses include Nuclear power generation, The sterilization of surgical instruments in hospitals Crime detection To find cracks and stresses in metals The preservation of food Dating artifacts Treating and diagnosing illness in medicine. Slide 14 Slide 15 Describe and explain the operation of a mass spectrometer Describe how the mass spectrometer may be used to determine relative atomic mass using the 12 Carbon scale Calculate non-integer relative atomic masses and abundance of isotopes from given data. Slide 16 The Mass Spectrometer Atoms have masses in the range of about 1x10 -24 to 1x10 -22 grams, and you cant weigh them in any conventional sense. You can, however, get around the problem. If you throw three balls, which all have the same diameter, at exactly the same speed on a very windy day. They present the same profile to the wind, but their masses are distinctly different e.g. a foam ball, a tennis ball and a cricket ball. Where would each ball land? Atoms can be deflected in a similar way to the balls in the wind by magnetic fields provided the atom is first turned into an ion. wind Throw direction Cricket ball Tennis ball Foam ball Slide 17 The positive ions pass through slits in negatively charged parallel plates, where they are accelerated. The ions are deflected by a magnetic field. The amount of the deflection depends both on the mass of the ion and on its charge. Heavier and less highly charged ions will be deflected less than lighter and more highly charged ions. Ions with a particular mass to charge ratio (m/z) are then recorded on a detector, which measures both the mass-to-charge ratio and the relative amounts of all the ions present. In practice, the machines electron beam energy can be adjusted so that only positive ions with a single charge are detected, so that the mass-to-charge ratio is the same as the mass. The sample is first vaporized turned to a gas The vapour is ionized by bombarding it with a stream of high-energy electrons from an electron gun to generate positive ions. M(g)+e- M + (g)+2e - Vaporised atom High-energy electron Unipositive ion Slide 18 Slide 19 The mass spectrometer produces a mass spectrum. From the mass spectrum of an element its relative atomic mass A r can be calculated, as it is equal to the weighted mean mass of all the naturally occurring isotopes of that element relative to 1/12 carbon-12. Detector current (arbitrary units) m/z (mass charge ratio) 6.83 9.13 12.17 2.60 The mass spectrum above is of naturally occurring germanium. From the mass spectrum it can be seen that; total detector current is = (6.83 + 9.13 + 2.60 + 12.17 + 2.60) = 33.33 The relative abundance of germanium-70 = 6.83 33.33 x100 = 20.5% Slide 20 The relative abundance of all the isotopes can be calculated in a similar way. Copy out this diagram in your notes and calculate the relative abundance of the other germanium isotopes. IsotopeRelative abundance / % 70 72 73 74 76 The relative atomic mass of germanium is given by A r = (70 x 20.5) + (72 x 27.4) + (73 x 7.8) + (74 x 36.5) + (76 + 7.8) 100 = 72.7 This is the number we see on the periodic table 20.5 27.4 7.8 36.5 7.8 Slide 21 The mass spectrum for chlorine Chlorine has two isotopes, 35 Cl and 37 Cl, in an approximate ratio of 3 atoms of 35 Cl to 1 atom of 37 Cl. So obviously the mass spectrum will consist of two lines at m/z 35 and 37, with the 35 line three times higher than the 37 line. But there is more Chlorine consists of molecules, not individual atoms So when chlorine passes into the ionisation chamber, an electron is knocked off the molecule to give a molecular ion, Cl 2 + Some of these ions fragment (fall apart) to give a chlorine atom and a Cl + ion. Cl 2 + Cl+Cl + (fragmentation) Slide 22 If the Cl atom formed is not then ionized by collision with an electron, it simply gets lost in the machine neither accelerated nor deflected. The Cl + ions will pass through the machine and will give lines at 35 and 37, depending on the isotope, with the m/z = 35 line 3 times taller than the 37 line. This is what you would expect. It will also record lines for unfragmented Cl 2 + ions. There are three different possible masses for a Cl 2 + ion depending on what combination of 35 Cl and 37 Cl atoms it contains. What could the masses be? Slide 23 The masses could be: 35 + 35 = 70 35 + 37 = 72 37 + 37 = 74 So in addition to the lines at 35 and 37, there will also be lines at 70, 72 and 74 You should be able to work out the relative height ratio to be 9:6:1 but you cannot make predictions about the relative heights of the 35/37 compared to those at 70/72/74. That depends on what proportion of the molecular ions break up into fragments. Slide 24 Problems on isotopes: 1) Calculate the relative atomic mass of silicon, given: 2) Calculate the relative atomic mass of gallium given the percentage abundances: 69 Ga 60.2%, 71 Ga 39.8% 3) Bromine has two isotopes, 79 Br and 81 Br. At what values of m/e would you expect to find lines in the mass spectrum of bromine, Br 2 ? (assume that only 1+ ions are formed), Relative isotopic massRelative abundance 28100 295.10 303.36 Slide 25 Slide 26 Describe the electromagnetic spectrum Distinguish between a continuous spectrum and a line spectrum Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels. Deduce the electron arrangement for atoms and ions up to Z = 20 Slide 27 The Electromagnetic spectrum Electromagnetic waves can travel through space and, depending on the wavelength, also through matter. The velocity of travel, c, is related to its wavelength, and its frequency, f. Velocity is measured in m s -1, wavelength in m and frequency in s -1 so it is easy to remember the relationship between them. c (m s -1 ) = (m) x f (s -1 ) Slide 28 The electromagnetic spectrum Slide 29 Although all e-m waves travel at the same speed, their wavelength [ ] and frequency [] can be different. The properties, dangers and uses of e-m waves depends on the wavelength [ ]. Waves that cook food. Waves that cause sun-tans. Slide 30 Electromagnetic radiation is a form of energy. The smaller the wavelength and thus the higher the frequency, the more energy the wave possesses. Electromagnetic waves have a wide range of wavelengths ranging from low energy radio waves to high energy gamma ( ), radiation. As you have seen visible light occupies a very narrow part of the spectrum. Slide 31 The Atom It was the Greek philosopher Democritus who first considered the idea that matter is made up of particles in about 400BC The idea was not accepted because there was no experimental evidence for it. John Dalton revived the discussion around 1801 and compiled experimental evidence which convinced people. Around the year 1900, physicists began to find evidence that atoms are made up of smaller particles Slide 32 Atomic spectra If sunlight or light from an electric light bulb is formed into a beam by a slit and passed through a prism on to a screen, a rainbow of separated colours is seen. The spectrum of colours is composed of visible light of all wavelengths and is called a continuous spectrum Slide 33 The colours Red, Orange, Yellow, Green, Blue, Indigo and Violet make up the colours in the visible part of what is known as the electromagnetic spectrum. Each colour is a wavelength which represents a particular amount of energy. Slide 34 If atoms and molecules are heated to sufficiently high temperatures, they emit light of certain wavelengths. The observed spectrum is called an atomic emission spectrum or line spectrum. Slide 35 All substances give emission spectra when they are excited in some way, by the passage of an electric discharge or by a flame. The atomic emission spectra of elements are in the visible and ultraviolet regions of the spectrum. When sodium or a sodium compound is put into a flame, it colours the flame yellow. A tube of hydrogen gas which has been excited by an electric discharge glows a reddish-pink colour. Slide 36 The Hydrogen Spectrum Viewed through a spectrometer, the emission spectrum of hydrogen is seen to be number of separate sets of lines or series of lines. These series of lines are named after their discoverers, the only series of lines we can observe with our eyes are in the visible spectrum called the Balmer series. Slide 37 The Balmer series of hydrogen The Balmer series is in the visible part of the spectrum. In each series, the intervals between the frequencies of the lines become smaller and smaller towards the high frequency end of the spectrum until the lines run together or converge to form a continuum of light. Slide 38 Why do atomic spectra consist of discrete (separate) lines? Why do atoms absorb or emit light of certain frequencies? Why do the spectral lines converge to form a continuum? The Rutherford picture of the atom offers no explanation. The theories continued Slide 39 Niels Bohr In 1913, Niels Bohr (1885-1962) put forward his picture of the atom to answer these questions. Bohr referred to Max Planks recently developed quantum theory, according to which energy can be absorbed or emitted in certain amounts, like separate packets of energy, called quanta. Slide 40 The Bohr model: Bohr Suggested An electron moving in an orbit can have only certain amounts of energy, not an infinite number of values: its energy is quantised The energy that an electron needs in order to move in a particular orbit depends on the radius of the orbit. An electron in an orbit distant from the nucleus requires higher energy than an electron in an orbit near the nucleus. Slide 41 If the energy of the electron is quantised, the radius of the orbit also must be quantised. There is a restricted number of orbits with certain radii, not an infinite number of orbits. An electron moving in one of these orbits does not emit energy. In order to move to an orbit farther away from the nucleus, the electron must absorb energy to do work against the attraction of the nucleus. If an atom absorbs a photon (a quantum of light energy), it can promote an electron from an inner orbit to an outer orbit. Slide 42 For an electron to move from an orbit of energy E1 to one of energy E2, the light absorbed must have a frequency given by Plancks equation: hv = E 2 E 1 where v = frequency, h = Plancks constant The emission spectrum arises when electrons which have been excited (raised to orbits of high energy) drop back to orbits of lower energy. They emit energy as light with a frequency given by Plancks equation Slide 43 Bohr assigned quantum numbers to the orbits. He gave the orbit of lowest energy (nearest to the nucleus) the quantum number 1 an electron in this orbit is in its ground state. The next energy level has quantum number 2 and so on. If the electron receives enough energy to remove it from the attraction of the nucleus completely, the atom is ionised. Slide 44 To summarise When energy is supplied to an atom electrons are excited (gain energy) from their lowest state to n excited state. Electrons can only exist in certain fixed energy levels. When electrons drop from a higher level to a lower level they emit energy. This energy corresponds to a particular wavelength and shows up as a line in the spectrum. Slide 45 When it returns to its ground state it releases the energy it absorbed as a specific wavelength HEAT n = 1 e e e e n = 2n = 3n = 4 Each shell of an atom represents an energy level, these are shown by n = 1, n = 2, n = 3 to n = infinity When an electron is excited (gains energy) it can move from a lower energy to a higher energy level When electrons return to the first level (n = 1) the series of lines occurs in the UV region as this involves the largest energy change The visible region spectrum is formed by electrons dropping back to the n = 2 level The first series in the infra red is due to electrons falling to the n = 3 level. Slide 46 656486 365 1 2 3 4 5 /nm Slide 47 Development of ideas The idea of the atom is still developing but what you must appreciate is that these are theories backed up by scientific evidence, but evidence is not proof!!!!! Bohrs model is an accepted model of the atom and is used as a building block for other ideas and theories, it is by no means the end of the story. Slide 48 Using Bohrs model We can use Bohrs model to help illustrate bonding in metallic, ionic and covalent compounds. For this we need to be able to show electron arrangements. Can you remember how many electrons can fit into the first shell? How many can fit into the second? How do you know this?????????? First we need to look at ionisation energies. Slide 49 Ionisation energies When an electron receives enough energy to remove it from the attraction of the nucleus completely, the atom is ionised. The ionisation energy of an element is defined as: The energy required to remove one mole of electrons from one mole of atoms in the gas phase to form one mole of cations in the gas phase A (g) A + (g) + e- 1 2 3 4 5 nucleus Slide 50 Successive Ionisation energies Successive ionisation energies for the same element also provide evidence for the electron arrangement in each energy level. X (g) X + (g) + e first ionisation energy X + (g) X 2+ (g) + e second ionisation energy X 2+ (g) X 3+ (g) + e third ionisation energy As more electrons are removed, the electrostatic pull of the protons holds the remaining electrons more tightly, so considerably more energy is required to remove them; hence a logarithmic scale is usually used. Slide 51 Activity: Using the information in the table below calculate the log of each ionisation energy and plot a graph of number of electrons lost and log 10 energy Can you compare the pattern of ionisation energies with the electron arrangement of potassium that you learnt in Grade 10? Number of electrons removedEnergylog10 1419 23051 34412 45877 57975 69649 711343 814942 916964 1048577 1154433 1260701 1368896 1475950 1583152 1693403 1799771 18444911 19476075 Slide 52 Slide 53 Electron arrangement Each of the energy levels described by the principal quantum number can contain only a certain number of electrons. Generally, when the level is full, extra electrons begin to fill up the next available energy level. In the case of the third level, after eight electrons have been added the next two electrons occupy the fourth level before the next ten electrons then complete the third level. Principal quantum numberMaximum number of electrons 12 28 38 or 18 Slide 54 Objectives AHL Explain how evidence from first ionisation energies across periods accounts for the existence of main energy levels and sub-levels in atoms. Explain how successive ionisation energy data is related to the electron configuration of an atom. State the relative energies of s,p,d and f orbitals in a single energy level. State the maximum number of orbitals in a given energy level. Draw the shape of an s orbital and the shapes of p x, p y and p z orbitals Apply the Aufbau principle, Hunds rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 54 Slide 55 The nature of the electron The Bohr model describes electrons as very small particles occupying fixed energy levels or shells. In the 1920s Louis de Broglie (1892-1987) proposed that electrons possess a dual nature, and that as well as behaving like particles they also have the properties of waves. Evidence for this is that electrons, in a similar manner to light, can be diffracted by passing them through very narrow slits. Slide 56 The uncertainty principle Werner Heisenberg (1901-1976) published his uncertainty principle shortly after Louis de Broglies proposal of the dual nature of the electron. The principle states that it is impossible to know the exact position and momentum of an electron. As momentum is related to time, what this implies is that it is impossible to know the exact location of an electron at an exact moment in time. The more precisely the time is known, the less precisely the location is known, and vice versa. Heisenbergs uncertainty principle actually applies to any particle with mass, but the more massive the particle the less the uncertainty. Slide 57 Quantum mechanics This concept has been extended by philosophers to the whole of the physical universe, which can be seen to exist as a collection of probabilities, because no individual particle or event can be located precisely in time. This led to Albert Einstein to retort, I cannot believe that God would choose to play dice with the universe. The new science of quantum mechanics was further developed in the 1920s by the Austrian physicist Erwin Schrodinger (1887-1961). The mathematical solution to the Schrodinger wave equation describe the three dimensional shapes of the atomic orbitals where there is a high probability that electrons are located. Slide 58 Sub-levels and types of orbital If you look closely at the emission spectrum of sodium the yellow colour is not actually one single discrete line, but two lines very close together. One has a wavelength of 5.889 x 10 -7 and the other 5.895 x 10 -7. This is evidence that the energy levels can be split into sub-levels. The difference in energy between the two lines will be equivalent to the energy difference between the split sub-levels. This also explains why the graph of the first ionisation energy against atomic number does not show a regular increase as the electrons are added. Slide 59 Patterns of ionisation energies in the periodic table Look up the first ionisation energies for each of the first 20 elements in the periodic table and draw (as accurately as is possible) a graph of atomic number (x axis) versus first ionisation energy (y axis) Ionisation energy values can be found from emission spectra, and the values can be found in any chemical book of data. Slide 60 2 8 8 Slide 61 Explaining Ionisation Energies A graph of first ionization energies plotted against atomic number shows a repeating pattern. It can be seen that the highest value is for Helium, an atom that contains two protons and two electrons. The two electrons are in the lowest level (n = 1) and are held tightly by the two protons. For lithium it is relatively easy to remove an electron. Can you suggest a reason for this? Slide 62 It suggests that the third electron in lithium is in a higher energy level than the first two. The value then generally increases until element 10, neon, is reached before it drops sharply for sodium. This graph provides evidence that the levels can contain different numbers of electrons before they become full. Slide 63 n = 1n = 2n = 3 Relating Ionisation Energies to Atomic Structure Slide 64 Slide 65 Stability points You can see on the graph of successive ionisation energies that there is not a regular increase in the ionisation energies across a period although there is a general one. As we noticed before there are points on the graph which show particular stability; i.e. Helium has a particularly more stable electron arrangement than hydrogen in period 1, and Neon has a more stable arrangement than Lithium in period 2. But we can also see in period 2 that Beryllium has a slightly more stable arrangement than Boron and that Nitrogen has a slightly more stable arrangement than Oxygen. Slide 66 There is a particular stability at Beryllium And another slight stability at nitrogen The same pattern repeats itself in the periodic table Slide 67 Sub-levels n = 1 n = 2 n = 3 n = 4 n = 5 In fact it is thought that the quantum levels (apart from n = 1) consist of more than 1 level Each quantum level consists of an s orbital, all but n = 1 consist of p orbitals, all but n = 1 and n = 2 consist of d orbitals and all but n = 1, 2 and 3 consist of f orbitals 1s 2s 3s 4s 2p 3p 4p 3d 4d Slide 68 Orbitals From quantum theory it can be shown that electrons in an atom exist in orbitals, and that each orbital can hold a maximum of two electrons. Each electron in an atom is uniquely described by four different quantum numbers. The first number is the Principal number which describes the main energy level. The second number is the type of orbital this can be an s, p, d or f orbital The third number is the number of each type of orbital Fourth number describes the spin of the electron. Slide 69 Pauli exclusion principle The fourth quantum number refers to the spin of the electrons. Electrons in an orbital can either be spinning in one direction or in the opposite direction. The Pauli exclusion principle states that no two electrons in the same atom can have exactly the same four quantum numbers. Hence each orbital can contain a maximum of only two electrons, because if they hold more, at least two of the electrons would have the same four quantum numbers. Slide 70