Chapter 7: Atomic Structure and Periodicity I. Review of Atomic Structure a. Ernest Rutherford i. Gold Foil Experiment ii. Developed the Nuclear Atom iii. Electrons Float in Space Around the Nucleus iv. Key to changes in electron behavior found in electromagnetic radiation. II. Electromagnetic Radiation a. Energy that travels through space as a wave. b. Electromagnetic Spectrum i. Gamma Rays ii. X Rays iii. Ultraviolet iv. Visible Light v. Infrared vi. Microwaves vii. Radio Waves c. Waves i. Wavelength (λ) 1. Distance between identical points on two adjacent waves. ii. Frequency (ν) 1. Number of waves to pass a point in a given unit of time, usually a second. 2. Measured in Hertz (Hz) iii. Frequency and Wavelength are inversely proportional. 1. When one goes up the other goes down. 2. When numbers in inverse proportion are multiplied, the product is a constant. a. λν = c b. c is the speed of electromagnetic radiation c. c = 2.9979 x 108 m/s iv. Concept of Electromagnetic Radiation as a wave was first proposed by Sir Isaac Newton. 1. Waves of Electromagnetic Radiation amplify each other. 2. Waves of Electromagnetic Radiation interfere with each 3. other. III. The Nature of Matter a. Classical Physics – Before 1900 i. Matter 1. Made of Particles 2. Had Mass 3. Position in Space was known. ii. Energy 1. Massless 2. Delocalized iii. Existing theories of physics explain all observed phenomena when matter and energy are considered as separate things. b. Modern Physics – After 1900 i. Max Planck (1858 – 1947) 1. German Physicist 2. Examined light given off by very hot objects. a. Many objects give off colored light 3. Classical physics predicted that matter could absorb or emit all types of energy. a. Objects should only give off white light, since all wavelengths of visible light would be produced. 4. Planck determined that energy was given off in units or packets called quanta. a. The energy of a quantum is determined by its frequency. b. E = hν i. E is energy of a quantum. ii. h is a value called Planck’s Constant 1. h = 6.626 x 10-34 J s 5. Planck proposes the idea of quantized energy. ii. Albert Einstein (1879 – 1955) 1. German Physicist 2. Proposed that since energy was quantized, and electromagnetic radiation is energy, it must also be quantized. 3. Electromagnetic radiation can be viewed as a stream of particles. a. Particle of electromagnetic radiation is a photon. b. Each photon carries a quantum of energy. c. The energy of the quantum is determined by the frequency of the photon. i. Ephoton = hν ii. Ephoton = (hc)/λ 4. Einstein developed the theory of special relativity in 1905. a. E = mc2 b. m = E/c2 i. Important implication of this is that energy has mass. ii. Mass of a photon is determined by its wavelength. 1. m = E/c2 a. m = (hc/λ)/c2 b. m = h/(λc) 5. Dual nature of light a. Light exists as both a wave and a particle. b. Breaks down the early theories of Newton. c. Classical theories of physics do not work here. iii. Louis de Broglie (1892 – 1987) 1. If light exists as a wave and a particle, is the opposite true? 2. Can matter exist as both a particle and a wave? a. Particles cannot move at the speed of light. i. According to E = mc2 a particle must gain mass as it accelerates, and at the speed of light it must have an infinite mass, requiring an infinite amount of energy to move it at that speed. ii. Particles do move at some velocity (υ) b. de Broglie rewrote Einstein’s formulas, substituting velocity for c. i. m = h/(λυ) ii. λ = h/(mυ) – de Broglie’s Equation c. To test the de Broglie equation, we bombard a crystal with X Rays. i. The X Rays form a diffraction pattern of light and dark spots. 1. The light spots indicate where the X Rays were amplified. 2. The dark spots indicate where the X Rays were interfered with. ii. The crystal is then bombarded with electrons, and a similar pattern is seen. 1. This pattern is only possible if electrons have a wavelength. 3. This experiment verifies that electrons do behave as both particles and waves. 4. Larger objects have experimental wavelengths too small to observe, but it is commonly believed that all matter behaves according to the de Broglie equation. iv. Niels Bohr (1885 – 1962) 1. Danish Physicist 2. Examined the Atomic Emission Spectra of Hydrogen. a. When energy is absorbed by an atom it becomes “excited”. b. They release this excess energy and return to a ground state by emitting a photon of light. c. If all possible wavelengths of light are emitted, the light given off by all atoms would be white. d. Only a few wavelengths are given off “lines” of the visible spectrum. This produces colored light. 3. The energy in an atom is quantized. a. This ties in to the theories of Max Planck. 4. Bohr developed the Bohr model of the atom in 1913. a. Electrons do not simply float in space as Rutherford proposed, but travel around the nucleus on set, circular paths called orbits. b. Electrons can exist on an orbit, and when it absorbs a photon of energy, it can jump up to a higher orbit. c. When electrons lose a photon of energy, they jump down to lower orbits. i. These jumps are called quantum leaps. d. Electrons can exist on any orbit in an atom, but cannot exist between two orbits. i. This results in only specific photons given off, not all possible photons. e. The energy levels in an atom can be determined by the Bohr equation: i. E = -2.178 x 10-18 J(Z2/n2) 1. n is an integer – the higher the value of n, the higher the energy level. 2. Z is the nuclear charge (number of protons). ii. The energy change as an electron moves from one orbit to another can be determined by the following: 1. ΔE = -2.178 x 10-18 J (Z2/nfinal2 – Z2/ninitial2) 5. The Bohr model appeared very promising at first, but it only works for Hydrogen. For atoms with more than one electron, the mathematics breaks down. IV. The Quantum Model of the Atom a. Collaborative work of three Physicists i. Werner Heisenberg (1901 – 1976) ii. de Broglie iii. Erwin Schrödinger (1887 – 1961) b. Schrödinger focused on the wave properties of electrons to examine atomic structure. i. Assumed the electron was a standing wave. ii. Calculated a wave function – function of the coordinates x, y, and z of the electron’s position in space. iii. A specific wave function is called an orbital. 1. An orbital is an area of three dimensional space where an electron is probably going to be found (90% of the time). 2. We do not know how an electron moves within the orbital. c. Werner Heisenberg i. There is a limit to how precisely we can know the position and momentum of a particle at any time. ii. The better we know a particle’s position, the less accurately we know its velocity, and vice-versa. iii. This is the Heisenberg Uncertainty Principle 1. Δx . Δ (mυ) ≥ h/4π a. Δx = Uncertainty in particle position b. Δ (mυ) = Uncertainty in particle momentum c. h = Planck’s Constant d. Square of the Wave Function is a Probability Distribution i. Also referred to as an electron density map – shows where electrons are most likely to be. ii. A physical drawing of an orbital is really an electron density map. iii. All possible orbitals overlapping each other in an atom make up something called the electron cloud. e. Quantum Mechanical Atomic Model i. Electrons exist in an electron cloud around the nucleus, which is an area where electrons can be found 90% of the time. f. Linus Pauling (1901 -1994) i. American Chemist ii. Described Schrödinger’s orbitals using a series of values he called quantum numbers. 1. Each Quantum Number describes a different property of the orbital. 2. There are four in total. 3. No two electrons in the same atom can be described by the same set of quantum numbers. (Pauli Exclusion Principle) a. Principal Quantum Number (n) i. Always an integer (1,2,3,…) ii. Indicates size and energy of the orbital. iii. Often described as the energy level. b. Angular Momentum Quantum Number (l) i. Always an integer ranging from 0 to (n-1) ii. Indicates the shape of the orbital. iii. The value of l is usually represented by a letter 1. l = 0 is s 2. l = 1 is p 3. l = 2 is d 4. l = 3 is f 5. l = 4 is g c. Magnetic Quantum Number (ml) i. An integer ranging from l to –l. ii. The value of ml relates to orientation in three dimensional space (x, y, or z). d. Spin Quantum Number (ms) i. Two possible values, +1/2 or -1/2 ii. Value indicates the spin of the electron on its axis. iii. +1/2 = clockwise spin iv. -1/2 = counterclockwise spin v. No idea which direction an electron spins, but only two can occupy the same orbital and the must spin in opposite directions to do this, so the values are assigned arbitrarily. V. The History of the Periodic Table a. Dmitri Mendeleev (1834-1907) i. Russian Chemist ii. Wrote a chemistry textbook in which he grouped elements together based on: 1. Increasing Atomic Mass 2. Similar Properties iii. Used his table to predict existence and properties of undiscovered elements. 1. Gallium (Ga) 2. Scandium (Sc) 3. Germanium (Ge) iv. In some cases, he did not follow the rule of increasing atomic mass to make elements with similar properties group together. 1. This problem was resolved when atomic number was discovered and placed into the table. 2. Mendeleev’s table perfectly followed the order of increasing atomic number. v. Because Mendeleev’s table has powerful predictive properties, it is widely used today, and Mendeleev is considered the father of Periodic Law. VI. Electron Configurations a. Aufbau Principle i. Lowest energy orbitals are always filled with electrons first. ii. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, … b. Hund’s Rule i. Orbitals with equal energy must each be half filled before any one orbital can be totally filled. ii. s – 1 orbital iii. p – 3 orbitals iv. d – 5 orbitals v. f – 7 orbitals vi. g – 9 orbitals c. Electrons in Highest Occupied Energy Level i. Valence Electrons ii. Determine all of an atoms chemistry iii. Number of Valence Electrons is equal to Group # for all Representative Elements (Group A) iv. Eight Valence Electrons is Most Stable (Lowest Energy) configuration possible – Octet Rule v. For smaller elements (He – B) Two Valence Electrons fills s orbital and makes them very stable as well – Duet Rule. d. Electrons in lower energy orbitals i. Core Electrons ii. Not important for chemistry of atoms. e. Transition Metals i. Because of overlapping orbitals number of valence electrons and electron configurations can be hard to predict. 1. Reason multiple charges are possible. ii. Only exception we need to worry about is that it is lower in energy to steal one electron from an s orbital to fill or half fill a d orbital. 1. Copper, Chromium VII. Periodic Trends a. Ionization Energy i. Energy required to remove an electron from a gaseous atom. ii. There can be multiple ionization energies. 1. First Ionization energy – energy needed to remove one electron. Always the lowest. 2. Second Ionization energy – energy required to remove a second electron. 3. Biggest jump in ionization energy always occurs when a core electron is removed. 4. Across a period (L to R) 1st ionization energy tends to go up. (Some exceptions) 5. Down a column, 1st ionization energy always goes down. b. Electron Affinity i. Energy change associated with the addition of an electron to a gaseous atom. ii. The more negative electron affinity is, the more energy is released, and the more favorable it will be. 1. Nonmetals tend to have very negative electron affinities. 2. Metals tend to have positive (endothermic) ones. 3. L to R across a period, electron affinity tends to get more negative. (Excluding Group 8A) 4. As you go down a column, electron affinity becomes more positive. (Many exceptions) c. Atomic Radius i. ½ the distance between the nuclei of two identical bonded atoms. ii. L to R across a period, atomic radius goes down. 1. More protons added to the nucleus pull electrons in tighter. iii. Down a column, atomic radius goes up. 1. Electrons placed into successively higher energy levels farther away from the nucleus. d. Electronegativity i. Ability of an atom in a compound to attract electrons to itself. ii. Typically, elements with high electron affinities have high electronegativities for similar reasons. iii. L to R across a period, electronegativity goes up. (Excluding Group 8A). iv. Down a column, electronegativity goes down. v. F is most electronegative element on the periodic table.
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