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Atoms: very small Comparing the size of a helium atom to a tennis ball... ... is like comparing a tennis ball to the Earth. The ancient Greeks theorised that everything was made of fire, earth, water and air. Only after over 1600 years did scientists realise matter was composed of different elements. In 1805 John Dalton refined a vague idea that matter was indivisible. In 1805, John Dalton proposed explanations for some observations he had made during experiments. He developed an atomic theory. 1. Elements are made of extremely small particles called atoms. 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. 3. Atoms cannot be subdivided, created, or destroyed. 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged. Dalton predicted that each element was made of a type of atom and he created symbols for them. He also put them together to make compounds. After he discovered the electron in 1897, J J Thomson suggested in 1904 that electrons are embedded in a blob of positive charge. It was nicknamed the “plum pudding model”. Rutherford’s gold foil experiment In 1909, Ernest Rutherford bombarded a thin gold foil with alpha particles. He thought they would go through, but be “bent” by the positive charge in the atoms Instead, most went straight through and a few bounced straight back. Rutherford had discovered the positivelycharged nucleus and disproved Thomson’s theory. But where were the electrons? Niels Bohr suggested they orbited the nucleus in “shells” around it, like planets orbiting the sun. Quantum Theory, 1920s The modern atomic model Electrons exist in a “cloud” around the nucleus. They can be thought of as waves, particles or mathematical functions How big is a nucleus? The helium atom is not in the right proportions. The three subatomic particles are wrongly enormous in comparison to the atom’s radius. Most of the atom is empty space. The nucleus is extremely small. But we draw it like this anyway, for obvious reasons. Ease of drawing. The number of protons is what defines an element A is the mass number: protons + neutrons Z is the atomic number, or proton number. 4 2 He Z = 2 4 Helium He 2 No. of protons, Z Number of protons = Electrons 2 Protons 2 Neutrons 2 Number of electrons Subatomic particles in more detail Subatomic particle Relative charge Relative mass Common depiction Proton +1 1 + Neutron 0 1 -1 1 ____ 1836 Electron Actual mass of electron, me = 9.11 × 10−31 kg - Lithium Mass number, A 7 Li 3 Atomic number, Z Number of protons (+) = Number of electrons (-) The charges must be balanced in a neutral atom. Lithium Electrons 3 Protons 3 Neutrons 4 Number of protons does not always equal the number of neutrons. Isotopes The number of protons is the “soul” of an atom. Most elements have different isotopes: atoms of the same element with different numbers of neutrons. Chlorine (Z = 17) has 2 naturally-occurring isotopes: 37Cl Neutrons 35Cl 37Cl 18 20 % abundance 75.53 24.47 35Cl Thus the Ar of chlorine is: ( )+( ) = and Isotopes Write the symbol for each of the following isotopes 1. Xenon-133 2. Oxygen-15 3. Iodine-125 4. Phosphorus with 18 neutrons 5. Cobalt with number of neutrons = Z + 3 6. The three isotopes of hydrogen 7. The most abundant isotope of carbon Relative atomic mass Symbol: Ar Since elements have different isotopes, the mass stated in the periodic table is a weighted average. It is stated relative to a particular “standard”. It has a mass of 1 and is equal to 1/12 of the mass of an atom of 12C. Relative atomic mass = average mass of the naturally-occurring isotopes of an element relative to C-12. Physical properties of isotopes 10B and 11B have different atomic masses. This affects their properties: • Density • Boiling point • Melting point • Deflection by a magnetic field • Rate of diffusion Heavier things require more energy to move! This means it is possible to separate isotopes. Useful isotopes 14C “Carbon dating” of organic matter up to about 60,000 years old. In use since 1949. 131I Radiotherapy, particularly to test functioning of thyroid gland, or to treat cancer of the thyroid gland, medical tracer 125I Used to treat brain and prostate cancer. 60Co Radiotherapy, sterilisation of medical equipment, industrial radiography, thickness and levelling tools, medical tracer t½ : 5730 years β- decay to 14N t½ : 8.07 days β- & γ decay to 131Xe t½ : 60 days γ decay to 125Te t½ : 5.27 years β- decay to 60Ni Electronic Structure The first 20 elements have their electrons arranged in a simple manner: The first energy level can take up to 2 electrons. The second energy level can take up to 8 electrons. The third energy level can take up to 8 electrons. After calcium (Ca) things start to get more complicated. Studying the transition metals requires a new method of arranging electrons on top of the first 20 elements (orbital theory). Electronic Structure All the Group 1 elements have 1 electron in the outermost energy level. 2,1 Li Lithium Na Sodium K Potassium Rb Rubidium Cs Caesium 2,8,1 2,8,8,1 The group number tells you how many electrons the atom has in its outer energy level. Magic! How many protons, electrons and neutrons do these have? Atom Na Rh phosphorus The last of the halogens Xe The only liquid non-metal Li+ FCarbon-14 (14C) 4 He 2 A helium atom P N 2 protons 2 electrons 2 neutrons E How many protons, electrons and neutrons do these have? Atom Na Rh phosphorus The last of the halogens Xe The only liquid non-metal Li+ FCarbon-14 (14C) 4 He 2 A helium atom P 11 45 15 85 54 35 3 9 6 N 12 58 16 125 77 45 4 10 8 2 protons 2 electrons 2 neutrons E 11 45 15 85 54 35 2 10 6 Fill in the table Species 238 92 Protons Electron Neutrons Electrons configuration U X 40Ca2+ 35 Cl Br-81 Fluoride ion X 10 Carbon-14 Mg ion in MgO 133 Xe 12 X Fill in the table Species 238 Protons Electron Neutrons Electrons configuration U 92 146 92 X 40Ca2+ 20 20 18 2,8,8 Cl 17 18 17 2,8,7 Br-81 35 46 35 2,8,8,17 Fluoride ion 9 10 10 2,8 Carbon-14 6 8 6 2,4 Mg ion in MgO 12 12 10 2,8 54 79 54 X 92 35 133 Xe Exciting electrons Electrons can jump from one energy level to another. Absorption of energy Emission of energy Electron jumps from n=3 to n=4 Electron jumps from n=4 to n=3 Exciting electrons – so what? Well, without these transitions the world would be boring and colourless! Yes, it’s true! Electron transitions in atoms are responsible for colour!! We can disturb these poor electrons by heating atoms up. When they go back to their original arrangement, energy is emitted in the form of light. The light pattern is called a spectrum In the 1850s, Robert Bunsen and Gustav Kirchhoff invent spectroscopes. They use them to study the spectra of blood and other substances for criminal investigations. Just like a fingerprint, each substance has its own spectrum Lockyer 1868 Discovered helium while looking through a spectroscope at the sun during an eclipse The spectroscope A simple device: 1. A slit allows light in. 2. The light is split by a prism into a spectrum. 3. The resulting spectrum is projected onto a screen. Not the whole thing The continuous spectrum (“rainbow”) is the result of white light. If you start off with only a fraction of white light (i.e. a colour) then you won’t get a “rainbow”. Instead, you’ll get only certain bits of it. This is a line spectrum. Hydrogen Iron Light source Spectroscope Projection Continuous spectrum White light Slit Heated gas Prism Emission spectrum Cold gas Absorption spectrum White light Light source Spectroscope Projection Continuous spectrum White light Slit Heated gas Prism Emission spectrum Cold gas Absorption spectrum White light Absorption + emission spectra Hydrogen The Balmer series In the visible light spectrum (about 380 to 750nm) From n > 2 to n = 2 Ultraviolet visible region Infrared 656 nm n = 3 to n = 2 486 nm 434 nm 410 nm n = 4 to n=2 n = 5 to n=2 n = 6 to n=2 Balmer series CONVERGENCE Infrared region Visible region Ultraviolet region Atomic Structure (HL) WRONG Electrons don’t orbit the nucleus. They don’t spin around it along a set path. Bohr model (1913) Also rather wrong. It’s not that simple. The electrons don’t spend their time a specific distance from the nucleus. But this diagram is partly useful if one considers the “shells” not to be orbits but energy levels. A better model The electron can be considered to be a particle, but also a wave or wave packet. This means there is a probability associated with the location of the electron – it’s not in a definite place. Heisenberg’s uncertainty principle “You cannot determine both the position and momentum of an electron at the same time.” Electron density How do we describe where the electron is, then? Electron density is a measure of the probability of finding the electron in a particular place. Covalent bonds are areas of high electron density Pentacene, C22H14 How do you take a photograph of a molecule? Atomic Force Microscopy (AFM) Very fine tip, one atom thick. Carbon monoxide molecule attached to tip to make it sharper and produce better images The tip oscillates above the sample, measuring the forces between them February 2012: Image of charge distribution around a naphthalocyanine molecule September 2012 – New images of a similar molecule showing different bond lengths s orbitals d orbitals p orbitals z The long form of the periodic table Normally rather inconvenient, so the f-block is detached out of the way. Note that helium is actually in the s-block. s f d p The f block and its orbitals Starts to be filled after 6s. 7 different f orbitals, meaning up to 14 electrons go in each f sub-shell. Relative Energies of Sub-shells This explains the order or the s, p, d and f blocks in the periodic table Rules for arranging electrons Aufbau (“building up”) principle “Electrons enter the lowest available energy level.” Hund’s rule of “maximum multiplicity” “When in orbitals of equal energy, electrons will try to remain unpaired.” This is because electrons repel each other Pauli exclusion principle “No two electrons can have the same four quantum numbers.” i.e. A full orbital must contain oppositespin electrons. n >0 principal quantum number (main shell) l 0 to n-1 azimuthal q.n. (sub-shell: s, p, d or f) m -l to +l magnetic q.n. (orbital within sub-shell) s ½ spin q.n. (up or down) 3p E •n=3 3s 2p 2s Example: Silicon (Z=14) The last electron has: • l = 1 (i.e. p orbital) • m = 0 (py - middle p) • s = + ½ (spin up) 1s No other electron in this atom has the same set of four quantum numbers Matching orbitals to quantum numbers Evidence: ionisation energies Ionisation Energy is the energy required to remove a mole of electrons from a mole of gaseous atoms. It is determined largely by the Effective Nuclear Charge This is the attractive force from the nucleus, minus the effect of inner electrons counteracting it. I.E. is also affected by electron-electron repulsion within the same main energy level. Successive ionisation energies increase due to decreased electron-electron repulsion, and increasing ENC. First n=2 electron First n=1 electron Variation in 1st I.E. with atomic number Why the drop between Mg-Al? Why the drop between P-S?