* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Atoms
Bremsstrahlung wikipedia , lookup
Matter wave wikipedia , lookup
Molecular orbital wikipedia , lookup
Elementary particle wikipedia , lookup
Ferromagnetism wikipedia , lookup
Molecular Hamiltonian wikipedia , lookup
Theoretical and experimental justification for the Schrödinger equation wikipedia , lookup
Wave–particle duality wikipedia , lookup
Hydrogen atom wikipedia , lookup
X-ray fluorescence wikipedia , lookup
Electron scattering wikipedia , lookup
Auger electron spectroscopy wikipedia , lookup
X-ray photoelectron spectroscopy wikipedia , lookup
Chemical bond wikipedia , lookup
Tight binding wikipedia , lookup
Atomic orbital wikipedia , lookup
1 of 47 © Boardworks Ltd 2007 2 of 47 © Boardworks Ltd 2007 Elements – different types of atom Elements are the simplest substances. There are about 100 different elements. Each element is made up of very tiny particles called atoms, and each element is made up of just one particular type of atom, which is different to the atoms in any other element. Gold is an element made up of only gold atoms. Carbon is an element made up of only carbon atoms. 3 of 47 © Boardworks Ltd 2007 Atoms – the building blocks John Dalton had the first ideas about the existence of atoms over 200 years ago. However, it is only relatively recently that special microscopes (called electron microscopes) have been invented that can actually ‘see’ atoms. This image is highly magnified. What could it be showing? The yellow blobs are individual gold atoms, as seen through an electron microscope. 4 of 47 © Boardworks Ltd 2007 Dalton’s Atomic Theory: 5 of 47 © Boardworks Ltd 2007 How was atomic structure discovered? 6 of 47 © Boardworks Ltd 2007 What particles are atoms made of? For some time, people thought that atoms were the smallest particles and could not be broken into anything smaller. Scientists now know that atoms are actually made from even smaller particles. There are three types: proton neutron electron How are these particles arranged inside the atom? 7 of 47 © Boardworks Ltd 2007 How did our understanding change? J.J Thomson discovered electrons, and proposed the existence of a (+) particle. It wasn’t until Rutherford’s famous gold foil experiment that the + proton was discovered, and atoms were thought to me mostly empty space. He named the centre of atoms the nucleus. Bohr improved on this model proposing that electrons move around the nucleus in specific layers called shells. It was James Chadwick who discovered particles with no charge, which he named neutrons. 8 of 47 © Boardworks Ltd 2007 What is the structure of an atom? Protons, neutrons and electrons are not evenly distributed in an atom. The protons and neutrons exist in a dense core at the centre of the atom. This is called the nucleus. The electrons are spread out around the edge of the atom. They orbit the nucleus in layers called shells. 9 of 47 © Boardworks Ltd 2007 Labelling the atom 10 of 47 © Boardworks Ltd 2007 Mass and electrical charge There are two properties of protons, neutrons and electrons that are especially important: mass electrical charge. Particle Mass Charge proton 1 +1 neutron 1 0 electron almost 0 -1 The atoms of an element contain equal numbers of protons and electrons and so have no overall charge. 11 of 47 © Boardworks Ltd 2007 How many protons? The atoms of any particular element always contain the same number of protons. For example: hydrogen atoms always contain 1 proton carbon atoms always contain 6 protons magnesium atoms always contain 12 protons. The number of protons in an atom is known as the atomic number or proton number. It is the smaller of the two numbers shown in most periodic tables. 12 of 47 © Boardworks Ltd 2007 What is the atomic number? What are the atomic numbers of these elements? 13 of 47 sodium 11 iron 26 tin 50 fluorine 9 © Boardworks Ltd 2007 What is mass number? Electrons have a mass of almost zero, which means that the mass of each atom results almost entirely from the number of protons and neutrons in the nucleus. The sum of the protons and neutrons in an atom’s nucleus is the mass number. It is the larger of the two numbers shown in most periodic tables. Atoms Protons Neutrons Mass number hydrogen 1 0 1 lithium 3 4 7 aluminium 13 14 27 14 of 47 © Boardworks Ltd 2007 What’s the mass number? mass number = number of protons + number of neutrons What is the mass number of these atoms? Atoms Protons Neutrons Mass number helium 2 2 4 copper 29 35 64 cobalt 27 32 59 iodine 53 74 127 germanium 32 41 73 15 of 47 © Boardworks Ltd 2007 How many neutrons? number of neutrons = mass number - number of protons = mass number - atomic number How many neutrons are there in these atoms? Atoms Mass number Atomic number Neutrons helium 4 2 2 fluorine 19 9 10 strontium 88 38 50 zirconium 91 40 51 uranium 238 92 146 16 of 47 © Boardworks Ltd 2007 Building a nucleus 17 of 47 © Boardworks Ltd 2007 How many electrons? Atoms have no overall electrical charge and are neutral. This means atoms must have an equal number of positive protons and negative electrons. The number of electrons is therefore the same as the atomic number. Atoms Protons Neutrons Electrons helium 2 2 2 copper 29 35 29 iodine 53 74 53 Atomic number is the number of protons rather than the number of electrons, because atoms can lose or gain electrons but do not normally lose or gain protons. 18 of 47 © Boardworks Ltd 2007 From atoms to ions How can reactive metal atoms become stable positive ions? 19 of 47 © Boardworks Ltd 2007 How do atoms form ions? An ion is an atom or group of atoms that has an electrical charge, either positive or negative. Atoms have an equal number of protons and electrons and so do not have an overall charge. Atoms with incomplete outer electron shells are unstable. By either gaining or losing electrons, atoms can obtain full outer electron shells and become stable. When this happens, atoms have an unequal number of protons and electrons and so have an overall charge. This is how atoms become ions. How does an atom become a positive or negative ion? 20 of 47 © Boardworks Ltd 2007 Positive and negative ions? An atom that loses electrons has more protons than electrons and so has a positive overall charge. This is called a positive ion. An atom that gains electrons has more electrons than protons and so has a negative overall charge. This is called a negative ion. The electron configuration of an atom shows how many electrons it must lose or gain to have a filled outer shell. Atoms with a nearly empty outer shell, will lose electrons to obtain a full outer shell. Atoms with a nearly full outer shell, will gain electrons to obtain a full outer shell. 21 of 47 © Boardworks Ltd 2007 How do atoms form positive ions? An atom that loses one or more electrons forms a positive ion. Metal atoms, such as sodium, magnesium and iron, form positive ions. Positive ions have a small ‘+’ symbol and a number by this to indicate how many electrons have been lost. This number is usually the same as the number of electrons in the atom’s outer shell. For example: lithium atom 2.1 lithium ion [ 2 ] = Li+ magnesium atom 2.8.2 magnesium ion [ 2.8 ] = Mg2+ aluminium atom 2.8.3 aluminium ion [ 2.8 ] = Al3+ 22 of 47 © Boardworks Ltd 2007 How is a sodium ion formed? Sodium atom: 11 protons = +11 11 electrons = -11 Sodium ion: 11 protons = +11 10 electrons = -10 Total charge = Total charge = 0 +1 + Na loses 1 electron 2.8.1 (partially full outer shell) 23 of 47 Na [2.8] (full outer shell) © Boardworks Ltd 2007 How is a magnesium ion formed? Magnesium atom: 12 protons = +12 12 electrons = -12 Magnesium ion: 12 protons = +12 10 electrons = -10 Total charge = Total charge = 0 +2 2+ Mg loses 2 electrons 2.8.2 (partially full outer shell) 24 of 47 Mg [2.8]2+ (full outer shell) © Boardworks Ltd 2007 How do atoms form negative ions? An atom that gains one or more electrons forms a negative ion. Non-metal atoms, such as chlorine, oxygen and nitrogen, form negative ions. Negative ions have a small ‘-’ symbol and a number by this to indicate how many electrons have been gained to fill their outer shell. For example: chlorine atom 2.8.7 chloride ion [ 2.8.8 ] = Cl- oxygen atom 2.6 oxide ion [2] = O2- nitrogen atom 2.5 nitride ion [2] = N3- The name of the ion is slightly different to the atom’s name. 25 of 47 © Boardworks Ltd 2007 How is a fluoride ion formed? Fluorine atom: 9 protons = +9 9 electrons = -9 Fluoride ion: 9 protons = +9 10 electrons = -10 Total charge = Total charge = -1 0 - F gains 1 electron 2.7 (partially full outer shell) 26 of 47 F [2.8](full outer shell) © Boardworks Ltd 2007 How is a sulfide ion formed? Sulfur atom: 16 protons = +16 16 electrons = -16 Sulfide ion: 16 protons = +16 18 electrons = -18 Total charge = Total charge = -2 0 2- S gains 2 electrons 2.8.6 (partially full outer shell) 27 of 47 S [2.8.8]2(full outer shell) © Boardworks Ltd 2007 28 of 47 © Boardworks Ltd 2007 31 of 26 © Boardworks Ltd 2009 32 of 26 © Boardworks Ltd 2009 Subatomic particles Atoms are composed of three subatomic particles: protons, neutrons and electrons. The two important properties of these particles are mass and charge: Particle Relative mass Relative charge proton 1 +1 neutron 1 0 electron 1/1840 -1 The mass of electrons is negligible when compared to the mass of protons and neutrons, so their mass is not included when calculating the mass of the atom. 33 of 26 © Boardworks Ltd 2009 Atomic number and mass number The number of protons in an atom is known as the atomic number or proton number and is represented by the symbol Z. The mass number of an atom is the number of protons plus the number of neutrons, and is represented by the symbol A. When an atom is represented by its symbol, the mass number, and sometimes the atomic number, are shown. 34 of 26 mass number (A) atomic number (Z) © Boardworks Ltd 2009 What are isotopes? Isotopes are atoms of the same element that contain different numbers of neutrons. mass number is different atomic number is the same carbon-12 carbon-13 The reactivity of different isotopes of an element is identical because they have the same number of electrons. The different masses of the atoms means that physical properties of isotopes are slightly different. 35 of 26 © Boardworks Ltd 2009 Isotopes of chlorine About 75% of naturally-occurring chlorine is chlorine-35 (35Cl) and 25% is chlorine-37 (37Cl). 36 of 26 17 protons 17 protons 18 neutrons 20 neutrons 17 electrons 17 electrons © Boardworks Ltd 2009 Isotopes of carbon There is also more than one isotope of carbon: Isotope Protons Neutrons 12C 6 6 13C 6 7 14C 6 8 All isotopes of carbon have 6 protons and so have 6 electrons. Because chemical reactivity depends on the number of electrons the reactivity of the isotopes of carbon is identical. 37 of 26 © Boardworks Ltd 2009 What’s the number? 38 of 26 © Boardworks Ltd 2009 39 of 26 © Boardworks Ltd 2009 Isotopes & Mass Spectrometry The number of naturally occurring isotopes of an element and their relative masses and abundances may be determined by a mass spectrometer, an instrument based on the same principles employed by JJ Thomson to discover the electron. Mass Spectrometry uses the fact that charged particles moving through a magnetic field are deflected from their original path based on their charge-to-mass ratio (e/m or e/z) The greater the ratio, the more deflection occurs. 40 of 26 © Boardworks Ltd 2009 ‘Weighing’ atoms Mass spectrometry is an accurate instrumental technique used to determine the relative isotopic mass (mass of each individual isotope relative to carbon-12) and the relative abundance for each isotope. From this, the relative atomic mass of the element can be calculated. Some uses of mass spectrometry include: 41 of 26 carbon-14 dating detecting illegal drugs forensic science space exploration. © Boardworks Ltd 2009 Mass spectrometry 42 of 26 © Boardworks Ltd 2009 Parts of the mass spectrometer 43 of 26 © Boardworks Ltd 2009 Process of mass spectrometry 44 of 26 © Boardworks Ltd 2009 Mass spectra of monatomic elements 45 of 26 © Boardworks Ltd 2009 Mass spectra of diatomic elements 46 of 26 © Boardworks Ltd 2009 Process of mass spectrometry 47 of 26 © Boardworks Ltd 2009 48 of 26 © Boardworks Ltd 2009 What is relative atomic mass? The relative atomic mass (Ar) of an element is the mass of one of its atoms relative to 1/12 the mass of one atom of carbon-12. relative atomic mass average mass of an atom × 12 = (Ar) mass of one atom of carbon-12 Most elements have more than one isotope. The Ar of the element is the average mass of the isotopes taking into account the abundance of each isotope. This is why the Ar of an element is frequently not a whole number. 49 of 26 © Boardworks Ltd 2009 Using mass spectra to calculate Ar The mass spectrum of an element indicates the mass and abundance of each isotope present. For example, the mass spectrum of boron indicates two isotopes are present: abundance (%) 100 11B 80 (80%) 60 10B 40 (20%) 20 0 0 2 4 6 m/z 8 10 12 How can this be used to calculate the Ar of boron? 50 of 26 © Boardworks Ltd 2009 Calculating Ar Most elements have more than one isotope. The relative atomic mass of the element is the average mass of the isotopes taking into account the abundance of each isotope. Example: what is the Ar of boron? In a sample of boron, 20% of the atoms are 10Br and 80% are 11Br. If there are 100 atoms, then 20 atoms would be 10Br and 80 atoms would be 11Br. The relative atomic mass is calculated as follows: Ar of Br = (20 × 10) + (80 × 11) 100 Ar of Br = 10.8 51 of 26 © Boardworks Ltd 2009 Calculating Ar of magnesium In a sample of magnesium, 79.0% of the magnesium atoms are 24Mg, 10.0% are 25Mg and 11.0% are 26Mg. Example: What is the Ar of magnesium? 1. Calculate mass × abundance of each isotope 24 × 79.0 25 × 10.0 26 × 11.0 2. Add these values, and divide by 100 (1896 + 250 + 286) / 100 Ar of Mg = 24.3 52 of 26 © Boardworks Ltd 2009 53 of 26 © Boardworks Ltd 2009 54 of 26 © Boardworks Ltd 2009 Ar calculations 55 of 26 © Boardworks Ltd 2009 60 of 39 © Boardworks Ltd 2009 How are atoms arranged? Where are the electrons found in the atom? 61 of 47 © Boardworks Ltd 2007 How are electrons arranged? Electrons are not evenly spread but exist in layers called shells. (The shells can also be called energy levels). The arrangement of electrons in these shells is often called the electron configuration. 1st shell 2nd shell 3rd shell Note that this diagram is not drawn to scale – the atom is mostly empty space. If the electron shells were the size shown, the nucleus would be too small to see. 62 of 47 © Boardworks Ltd 2007 How many electrons per shell? Each shell has a maximum number of electrons that it can hold. Electrons will fill the shells nearest the nucleus first. 1st shell holds a maximum of 2 electrons 2nd shell holds a maximum of 8 electrons 3rd shell holds a maximum of 8 electrons This electron arrangement is written as 2,8,8. 63 of 47 © Boardworks Ltd 2007 Models of atoms The model of the atom states that a nucleus is surrounded by shells of electrons. Each shell holds a different maximum number of electrons: 1st shell = 2 electrons 2nd shell = 8 electrons 3rd shell = 8 electrons. At the Chemistry 11 level, this model is slightly different. Instead of electrons being arranged in shells that are a different distance from the nucleus, they are arranged in energy levels. 64 of 34 © Boardworks Ltd 2009 Waves behaving like particles? Well hit me with a Planck! A serious challenge to Rutherford’s atomic model arose almost immediately. By the end of the 1800s the physics available stated that accelerating charges should radiate energy. This meant that atoms should collapse in a fraction of a second as their electrons lost energy and spiralled around the nucleus. Rutherford’s conclusions were correct, the problems was with the physics of the day… 65 of 34 © Boardworks Ltd 2009 The Bohr model of the atom In 1900, Max Planck (right) developed his ‘Quantum theory’, which states that energy could be shown to behave like particles in fixed amounts he called quanta. In 1913, Niels Bohr (left) applied Plank’s theory to electrons, and improved upon Rutherford’s model. He proposed that electrons could only exist in fixed energy levels. The main energy levels are called principal energy levels and are given a number called the principal quantum number (n) with the lowest in energy being 1. 66 of 34 © Boardworks Ltd 2009 The Bohr Model Each electron has a fixed energy = an energy level. Electrons can jump from one energy level to another. Electrons can not be or exist between energy levels. A quantum of energy is the amount of energy needed to move an electron from one energy level to another energy level. 67 of 34 © Boardworks Ltd 2009 Bhor’s experiments… Bhor knew that when high voltage was applied across the electrodes of a sealed glass tube containing a gas such as hydrogen, the gas was heated and emitted light…he looked at this light through a spectroscope. He saw that the spectroscope separated the light into its component wavelengths, and for hydrogen he saw a series of coloured lines against a black background. Bhor applied quantum principles to explain the bright-line spectrum he saw for hydrogen. 68 of 34 © Boardworks Ltd 2009 Quantum Theory rescues the Nuclear Model To move from one level to another, the electron must gain or lose the right amount of energy. The higher the energy level, the farther it is from the nucleus. • Gain energy to move to higher energy levels (away from nucleus) • Lose energy to move to lower energy levels (closer to nucleus) The degree to which they move from level to level determines the frequency of light they give off. 69 of 34 © Boardworks Ltd 2009 Bhor proposed… The amount of energy required to go from one energy level to another is the not same for the electrons. Higher energy levels are closer together. This means it takes less energy to change levels in the higher energy levels. Once in a higher energy orbit, any electron could then return to a lower energy orbit by emitting a specific amount of energy corresponding to the energy difference. If the frequency of that emitted energy corresponding to any part of the visible spectrum, then a bright line of that specific colour would be seen. 70 of 34 © Boardworks Ltd 2009 An Explanation of Atomic Spectra • The three groups of lines in the hydrogen spectrum correspond to the transition of electrons from higher energy levels to lower energy levels. 71 of 34 © Boardworks Ltd 2009 The shortcomings… Bohr’s model of the hydrogen atom was successful in explaining the mystery of bright line spectra. His calculations and predictions worked for hydrogen and he even calculated the radius of the orbit for hydrogen’s electron in its ground state. BUT his model failed to explain the energies absorbed and emitted by atoms with more than one electron. 72 of 34 © Boardworks Ltd 2009 If an excited electron emits energy and drops to n = 2 from a higher energy orbit, the wavelength of the emitted energy corresponds to a particular colour of visible light. If an electron drops from n = 3 to n = 2, the energy difference between the two orbits (and therefore the energy emitted) corresponds to that of red light. Hence the red line appears in the emission spectrum. The blue-green line results from an electron transition from n = 4 to n = 2, the blue line from an electron transition from n = 5 to n = 2, and the violet line from an electron transition from n = 6 to n = 2. This series of four bright lines in the visible spectrum is called the Balmer series, named after the Swiss schoolteacher who first derived a mathematical relationship between the lines in hydrogen’s visible emission spectrum. 73 of 34 © Boardworks Ltd 2009 Did you know that an element can be identified by its emission spectra? When atoms absorb energy, electrons move into higher energy levels. These electrons then lose energy by emitting light when they return to lower energy levels. Mercury 74 of 34 Nitrogen © Boardworks Ltd 2009 “Fingerprints” of certain atoms 75 of 34 © Boardworks Ltd 2009 What is ionization energy? Ionization is a process in which atoms lose or gain electrons and become ions. The first ionization (I1) energy of an element is the energy required to remove one electron from a gaseous atom. M(g) → M+(g) + e-(g) The second ionization (I2) energy involves the removal of a second electron: M+(g) → M2+(g) + e-(g) Looking at trends in ionization energies can reveal useful evidence for the arrangement of electrons in atoms and ions. 76 of 34 © Boardworks Ltd 2009 Ionization energy definitions 77 of 34 © Boardworks Ltd 2009 ionization energy Evidence for energy levels Plotting the successive ionization energies of magnesium clearly shows the existence of different energy levels, and the number of electrons at each level. 6 Successive ionization energies increase as more 5 electrons are removed. Large jumps in the ionization energy reveal where electrons are being removed from the next principal energy level, such as between the 2nd and 3rd, and 10th and 11th ionization energies for magnesium. 4 3 2 electron removed 78 of 34 © Boardworks Ltd 2009 More evidence for energy levels first ionization energy (kJ mol-1) The first ionization energies of group 2 elements also show evidence for the existence of different principal energy levels. Even though the nuclear charge increases down the group, the first ionization energy decreases. 900 800 700 600 500 400 Be 79 of 34 Mg Ca Sr element This means electrons are being removed from successively higher energy levels, which lie further from the nucleus and are less attracted to Ba the nucleus. © Boardworks Ltd 2009 Trends in first ionization energies 80 of 34 © Boardworks Ltd 2009 Energy levels 81 of 34 © Boardworks Ltd 2009 82 of 39 © Boardworks Ltd 2009 Beyond Bhor: Particles Behaving Like Waves The beginnings of quantum mechanics was in 1923 with Louis de Broglie: “If waves behaved like particles, could it also be possible for particles to have wave properties?” • Applied waveparticle theory to electrons n=4 n=5 • electrons exhibit wave properties n=6 Forbidden n = 3.3 83 of 34 © Boardworks Ltd 2009 Heisenberg: Quantum Mechanics • Treated the electron as a pure particle with quantum behavior. • Heisenberg Uncertainty Principle: impossible to know exactly both where any particle was located and where it was going as the same time (velocity) • Werner Heisenberg ~1926 Ie: It is impossible to describe precisely both the location and the speed of particles that exhibit wavelike behavior. g Microscope Electron 84 of 34 © Boardworks Ltd 2009 The Quantum Mechanical Model • Rutherford’s and Bohr’s model focused on describing the path of the electron around the nucleus like a particle (like a small baseball). • Austrian physicist Erwin Schrödinger rejected the idea of an electron as a particle, he treated the electron as a wave. The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation. 86 of 34 © Boardworks Ltd 2009 Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS VISIBLE LIGHT 87 of 34 ELECTRONS © Boardworks Ltd 2009 The electron as a WAVE Schrödinger’s wave equation Used to determine the probability of finding the H electron at any given distance from the nucleus Electron best described as a cloud • Effectively covers all points at the same time (fan blades) 88 of 34 © Boardworks Ltd 2009 Quantum Mechanics • Schrödinger Wave Equation (1926) finite # of solutions quantized energy levels which defines probability of finding an electron Ψ1s Erwin Schrödinger ~1926 1 Z 3/2 σ π a0 e The regions in 3D space around the nucleus in which there is the highest probability of finding electrons are called atomic orbitals. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 89 of 34 © Boardworks Ltd 2009 Atomic Orbitals • (fuzzy cloud) = An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron. 90 of 34 © Boardworks Ltd 2009 first ionization energy (kJ mol-1) Evidence for sub-levels 1600 1400 1200 The first ionization energies for the elements in period 3 show a general increase. 1000 However, aluminium’s 800 value is below that of magnesium. This suggests 600 that the third principal 400 energy level is not one Na Mg Al Si P S Cl Ar single energy level. element All principal energy levels contain one or more sub-levels, with different but exact energy values. 91 of 34 © Boardworks Ltd 2009 The sub-levels There are four sub-levels, labelled in order of increasing energy: s, p, d and f. Each holds a different number of electrons. Each principal energy level contains a different number of sub-levels. 92 of 34 sub-level max no. electrons s 2 p 6 d 10 f 14 principal energy level, n 1 2 3 4 sub-levels max no. electrons 1s 2 2s, 2p 8 3s, 3p, 3d 18 4s, 4p, 4d, 4f 32 © Boardworks Ltd 2009 Levels and sub-levels 93 of 34 © Boardworks Ltd 2009 Atomic Orbitals • The numbers and kinds of atomic orbitals depend on the energy sublevel. Energy Level, n 94 of 34 # of sublevels Letter of sublevels # of # of Total orbitals per electrons in electrons in sublevel each orbital energy level © Boardworks Ltd 2009 Atomic Orbitals • The numbers and kinds of atomic orbitals depend on the energy sublevel. Energy Level, n # of sublevels Letter of sublevels 1 1 s 1 2 2 2 2 s p 1 3 2 6 8 3 s p d 1 3 5 2 6 10 18 4 s p d f 1 3 5 7 2 6 10 14 32 3 4 95 of 34 # of # of Total orbitals per electrons in electrons in sublevel each orbital energy level © Boardworks Ltd 2009 Blocks of the periodic table 96 of 34 © Boardworks Ltd 2009 Order of sub-levels 97 of 34 © Boardworks Ltd 2009 The Aufbau principle As part of his work on electron configuration, Niels Bohr developed the Aufbau principle, which states how electrons occupy sub-levels. The Aufbau principle states that the lowest energy sub-levels are occupied first. This means the 1s sub-level is filled first, followed by 2s, 2p, 3s and 3p. However, the 4s sub-level is lower in energy than the 3d, so this will fill first. 98 of 34 © Boardworks Ltd 2009 99 of 39 © Boardworks Ltd 2009 Writing electron configuration 100 of © Boardworks Ltd 2009 Electron configuration: true or false? 104 of © Boardworks Ltd 2009 Electron configuration of transition metals Although the 3d sub-level is in a lower principal energy level than the 4s sub-level, it is actually higher in energy. This means that the 4s sub-level is filled before the 3d sub-level. Example: what is the electron structure of vanadium? 1. Count number of electrons in atom 23 2. Fill sub-levels, remembering 4s is filled before 3d 1s22s22p63s23p64s23d3 105 of © Boardworks Ltd 2009 Electronic configuration: atoms 106 of © Boardworks Ltd 2009 107 of © Boardworks Ltd 2009 Electron configuration of ions When writing the electron configuration of ions, it is important to add or subtract the appropriate number of electrons. For non-transition metals, the sublevels are then filled as for atoms. For negative ions add electrons. For positive ions remove electrons. Example: what is the electron structure of O2-? 1. Count number of electrons in atom 8 2. Add or remove electrons due to charge 8 + 2 = 10 3. Fill sub-levels as for uncharged atom 1s22s22p6 108 of © Boardworks Ltd 2009 Electronic configuration of transition metal ions When transition metals form ions, it is the 4s electrons that are removed before the 3d electrons. Example: what is the electron structure of Ni2+? 1. Count number of electrons in atom 28 2. Fill sub-levels, remembering 4s is filled before 3d 1s22s22p63s23p64s23d8 3. Count number of electrons to be removed 2 4. Remove electrons starting with 4s 1s22s22p63s23p63d8 109 of © Boardworks Ltd 2009 Electron configuration: ions 110 of © Boardworks Ltd 2009 111 of © Boardworks Ltd 2009 Electron orbitals As Heisenberg suggested, it is impossible to exactly locate the position of an electron within an energy sub-level. But Schrödinger’s wave equation showed by measuring the electron density around the nucleus, it is possible to define regions where electrons are most likely to be found at any one time. These regions are called orbitals. sub-level no. orbits max no. Each energy sub-level electrons has one or more orbitals, s 1 2 each of which can p 3 6 contain a maximum of two electrons. d 5 10 f 112 of 7 14 © Boardworks Ltd 2009 Shapes of electron orbitals 113 of © Boardworks Ltd 2009 Atomic Orbitals • Different atomic orbitals are denoted by letters. • The s orbitals are spherical, and p orbitals are dumbbell-shaped. • Four of the five d orbitals have the same shape but different orientations in space. 114 of © Boardworks Ltd 2009 The Pauli exclusion principle and spin The Pauli exclusion principle states that each orbital may contain no more than two electrons. It also introduces a property of electrons called spin, which has two states: ‘up’ and ‘down’. The spins of electrons in the same orbital must be opposite, i.e. one ‘up’ and one ‘down’. A spin diagram shows how the orbitals are filled. Orbitals are represented by squares, and electrons by arrows pointing up or down. 115 of spin diagram for magnesium, 1s22s22p63s2 © Boardworks Ltd 2009 Rules for filling electrons When two electrons occupy a p sub-level, they could either completely fill the same p orbital or half fill two different p orbitals. Hund’s rule states that single electrons occupy all empty orbitals within a sub-level before they start to form pairs in orbitals. If two electrons enter the same orbital there is repulsion between them due to their negative charges. The most stable configuration is with single electrons in different orbitals. 116 of © Boardworks Ltd 2009 Evidence for Hund’s rule first ionization energy (KJ mol-1) The first ionization energies for the elements in period 3 show a general increase. 1600 1400 1200 1000 800 600 However, sulfur’s value is below that of phosphorus. As the highest energy electrons of both are in the 3p sub-level this is evidence for Hund’s rule. 400 Na Mg Al Si P S Cl Ar element 117 of © Boardworks Ltd 2009 Evidence for Hund’s rule: P vs. S Phosphorus has three electrons in its 3p sub-level and sulfur has four. The lower first ionisation energy for sulfur is because it has a pair of electrons in one of the 3p orbitals. Mutual repulsion between these two electrons makes it easier to remove one of them. phosphorus 118 of sulfur © Boardworks Ltd 2009 Electron configuration of Cr and Cu The electron configurations of chromium and copper are exceptions to the normal rules of orbital filling: chromium copper 1s22s22p63s23p64s13d5 1s22s22p63s23p64s13d10 In each case the 4s orbital contains one electron. This is because the 4s and 3d sub-levels lie very close together in energy, and the 3d being either half full or completely full is a lower energy arrangement. With larger atoms like this it can be useful to shorten the electron arrangement. Copper can be shortened to [Ar]4s13d10. 119 of © Boardworks Ltd 2009 Creating spin diagrams 120 of © Boardworks Ltd 2009 Summary of Quantum Numbers 121 of © Boardworks Ltd 2009 122 of © Boardworks Ltd 2009 Glossary 123 of © Boardworks Ltd 2009 What’s the keyword? 124 of © Boardworks Ltd 2009 Energy levels and electrons 125 of © Boardworks Ltd 2009 Multiple-choice quiz 126 of © Boardworks Ltd 2009 127 of © Boardworks Ltd 2009 128 of © Boardworks Ltd 2009 What is periodicity? The term periodicity describes a repeating pattern in properties of elements across periods of the periodic table. The Russian chemist Dmitry Mendeleev is credited with being the creator of the first version of the periodic table. He observed that when the elements are arranged in order of atomic mass, there are recurring patterns in certain properties. The modern periodic table can be used to analyse trends in properties such as atomic radius across periods and down groups. 129 of © Boardworks Ltd 2009 What is atomic radius? The atomic radius of an element is difficult to precisely define because of the uncertainty over the size of the electron cloud. Several definitions are used. One definition is half the shortest internuclear distance found in the structure of the element. For non-metallic elements, the covalent radius is often used as the atomic radius. This is half the internuclear distance between two identical atoms in a single covalent bond. 130 of covalent radius © Boardworks Ltd 2009 More on atomic radius For non-bonded adjacent atoms (e.g. in a covalent crystal of a non-metallic element), the van der Waals radius is used as a value for atomic radius. This is half the shortest internuclear distance between two similar non-bonded atoms. van der Waals radius For metallic elements, the metallic radius is often used as the atomic radius. This is half the shortest internuclear distance between two adjacent atoms in a metallic bond. 131 of © Boardworks Ltd 2009 Trends in atomic radius in period 3 132 of © Boardworks Ltd 2009 Trends in atomic radius in period 3 Atomic Element radius (nm) Na 0.190 Mg 0.145 Al 0.118 Si 0.111 P 0.098 S 0.088 Cl 0.079 Ar 0.071 133 of The atomic radius of the elements across period 3 decreases. This might seem counterintuitive, because as the numbers of sub-atomic particles increase, the radius might be expected to also increase. However, more than 99% of the atom is empty space – the nucleus and electrons themselves occupy a tiny volume of the atom. © Boardworks Ltd 2009 Increase in proton number The number of protons in the nucleus of the atoms increases across period 3. proton number Element 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar This increase in the number of protons increases the nuclear charge of the atoms. The nucleus has stronger attraction for the electrons, pulling them in closer and so the atomic radius decreases across the period. increased nuclear charge pulls electrons closer 134 of © Boardworks Ltd 2009 What is shielding? 135 of © Boardworks Ltd 2009 Explaining atomic radius in period 3 Atomic Proton Element number radius (nm) Na 11 0.190 Mg 12 0.145 Al 13 0.118 Si 14 0.111 P 15 0.098 S 16 0.088 Cl 17 0.079 Ar 18 0.071 136 of Proton number increases across period 3, but shielding remains approximately constant. This causes an increase in effective nuclear charge, leading to a greater attraction between the nucleus and the outermost electrons. This pulls these electrons closer to the nucleus and results in a smaller radius. © Boardworks Ltd 2009 Atomic radius in period 3 137 of © Boardworks Ltd 2009 Atomic radius: true or false? 138 of © Boardworks Ltd 2009 139 of © Boardworks Ltd 2009 What is first ionization energy? Ionization is a process in which atoms lose or gain electrons and become ions. The first ionization energy of an element is the energy required to remove one electron from a gaseous atom. M(g) → M+(g) + eThe first ionization energy is therefore a measure of the strength of the attraction between the outermost electrons and the nucleus. The first ionization energies of the elements in periods 2 or 3 can give information about their electronic structure. 140 of © Boardworks Ltd 2009 Plot of the first ionization energies 141 of © Boardworks Ltd 2009 General trend in first ionization energy 1600 ionization energy (kJ mol-1) There is a general increase in the first ionization energies across period 3. 1400 1200 1000 800 Across period 3, the 600 proton number 400 increases but the Na Mg Al Si P S Cl Ar amount of shielding element does not change significantly. The effective nuclear charge therefore increases. The greater attraction between the nucleus and the outermost electrons means that more energy is required to remove an electron. 142 of © Boardworks Ltd 2009 Trend in first ionization energy: exceptions ionization energy (kJ mol-1) There are two exceptions to the general trend in first ionization energy: both aluminium and sulfur have lower ionization energies than might be expected. 1600 1400 1200 1000 800 600 400 Na Mg Al lower ionization energies than expected Si P S Cl Ar element 143 of © Boardworks Ltd 2009 First ionization energy of Al vs. Mg The first ionization energy of aluminium is less than that of magnesium, even though aluminium has a higher nuclear charge. The electron removed when aluminium is ionized is in a 3p sub-level, which is higher in energy than the 3s electron removed when magnesium is ionized. Removing an electron from a higher energy orbital requires less energy. magnesium 144 of aluminium © Boardworks Ltd 2009 First ionization energy of S vs. P The first ionization energy of sulfur is less than that of phosphorus, even though sulfur has a higher nuclear charge. The highest energy electron in both phosphorus and sulfur is in the 3p sub-level. However, in sulfur this electron is paired, while in phosphorus each 3p orbital is singly occupied. Mutual repulsion between paired electrons means less energy is required to remove one of them. phosphorus 145 of sulfur © Boardworks Ltd 2009 Ionization energy in period 3 146 of © Boardworks Ltd 2009 Ionization energy in period 3 147 of © Boardworks Ltd 2009 148 of © Boardworks Ltd 2009 Plot of the melting and boiling points 149 of © Boardworks Ltd 2009 Na, Mg and Al: melting and boiling points The melting and boiling points increase for the three metallic elements from sodium to aluminium. temperature (K) 3000 2500 2000 boiling point 1500 1000 500 0 Na 150 of melting point Mg element Al This is because the strength of the metallic bonds increases. More energy is needed to break the stronger metallic bonds, so melting and boiling points are higher. © Boardworks Ltd 2009 Na, Mg and Al: metallic bond strength The increase in metallic bond strength from sodium to aluminium is due to two factors: 1. Charge density. This is the ratio of an ion’s charge to its size. Na+ ions are large with a small charge, so have a low charge density. Al3+ ions are smaller with a larger charge, and so have a higher charge density. They are therefore more strongly attracted to the delocalized electrons. 2. Number of free electrons. Sodium has one free electron per metal ion, whereas aluminium has three. This leads to more attractions that must be broken in aluminium. 151 of © Boardworks Ltd 2009 Silicon Silicon has a macromolecular structure similar to that of diamond. Each silicon atom is bonded to four neighbouring silicon atoms by strong covalent bonds. These must be broken in order for silicon to melt. This requires a lot of energy, so silicon's melting and boiling points are high. 152 of © Boardworks Ltd 2009 Period 3 non-metals This is because they have a simple molecular structure with weak van der Waals forces holding the molecules together. temperature (K) The melting and boiling points of phosphorus, sulfur and chlorine are much lower than those of silicon. 3500 3000 2500 2000 1500 1000 500 0 Na Mg Al Si P element S Cl Breaking these forces of attraction requires much less energy than breaking covalent bonds. 153 of © Boardworks Ltd 2009 Ar Period 3 non-metals: structure 154 of © Boardworks Ltd 2009 Melting points in period 3 155 of © Boardworks Ltd 2009 Melting points in period 3 156 of © Boardworks Ltd 2009 157 of © Boardworks Ltd 2009 Glossary 158 of © Boardworks Ltd 2009 What’s the keyword? 159 of © Boardworks Ltd 2009 Multiple-choice quiz 160 of © Boardworks Ltd 2009