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Atomic Theory notes.notebook February 01, 2016 John Dalton: (1803) Atomic Theory 1. All matter made up of tiny indivisible particles called atoms 2. All atoms of the same element are identical 3. Atoms of different elements are different 4. Atoms combine in simple whole number ratios to form compounds 5. Atoms cannot be divided John Dalton: (1810) Law of Multiple Proportions – the same two elements can combine in different ratios to form different compounds. When comparing the masses between compounds, they are still small whole number ratios. Feb 18:27 AM Dalton came up with his ideas based solely on Large Scale experimentation (Macro World). He didn't have access to the kinds of equipment that we do today that allow us to see and do Small Scale experimentation (Micro World) The scientists who do most of the Micro World experiments begin their work shortly after the first world war (1920's). But Just Before That... Feb 18:32 AM Atomic Theory notes.notebook February 01, 2016 JJ Thompson, a British Chemist, (1897) discovered the electron and calculated the chargetomass ratio of e. His Cathode ray tube experiment was instrumental in beginning the study of the SubAtomic Particles. (e/m = 1.759 x 108 coulomb/g) Based on his findings, Dalton's Solid Sphere Atom is transformed into the Plumpudding model of the atom (blueberry muffin/Chocolate Chip Cookie model) Feb 18:34 AM JJ Thompson becomes the head of Britain's Foremost research facility and attracts scientists from all over the world to come to Britain to perform experiments and study. Robert Millikan: (1911) Oil‐Drop experiment established charge on e‐ (1.62 x 10‐19 coulomb) 1. sprayed oil into the upper chamber, transferred e‐ to drops 2. oil drops fall through chamber due to gravity 3. charge on plates adjusted to offset gravity and suspend drop charge on oil drops calculated **charges varied, but were all multiples of one fundamental charge, the charge of an e‐** Ernest Rutherford: (1911) Gold Foil Experiment – discovered the nucleus, led to discovery and identification of protons 1. 2. Bombarded gold foil with alpha particles Most particles passed right through Some particles greatly deflected James Chadwick: (1932) discovered the neutron; high energy, neutral particle emitted from the nucleus Feb 18:37 AM Atomic Theory notes.notebook February 01, 2016 Neils Bohr, A Danish Scientist, begins studying chemistry and physics, writes incredible papers that astound the scientific community, Pwns Einstein, and creates the following atomic Model: Bohr’s Planetary Model of the Atom: (1913) 1. Electrons travel in definite energy levels without releasing energy 2. Electrons in each orbit have a certain amount of energy 3. Energy of orbits increases as distance from the nucleus increases 4. Electrons lose energy when they drop to lower energy levels Feb 18:41 AM All of which leads up to Heisenberg and Schrodinger Einstein hates this, Bohr loves it. Heisenberg’s Uncertainty Principle: (1927) It is impossible to know the position and velocity of an e‐ at the same time. Schrödinger’s Wave Equation: (1926) complex mathematical function that treats the e‐ as a wave. The equation determines the probability of finding the e‐ in any given place around the nucleus. When the changing probabilities are plotted in 3‐D, the probability area for finding the e‐ becomes a “cloud”. Schrödinger’s wave equation used “quantum numbers” Feb 18:43 AM Atomic Theory notes.notebook February 01, 2016 So what do you need to know? Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Chm.1.1.2 • Analyze diagrams related to the Bohr model of the hydrogen atom in terms of allowed, discrete energy levels in the emission spectrum. • Describe the electron cloud of the atom in terms of a probability model. • Relate the electron configurations of atoms to the Bohr and electron cloud models. Chm.1.1.3 • Understand that energy exists in discrete units called quanta. 1. Describe the concepts of excited and ground state of electrons in the atom: 1. When an electron gains an amount of energy equivalent to the energy difference, it moves from its ground state to a higher energy level. 2. When the electron moves to a lower energy level, it releases an amount of energy equal to the energy difference in these levels as electromagnetic radiation (emissions spectrum). • Articulate that this electromagnetic radiation is given off as photons. • Understand the inverse relationship between wavelength and frequency, and the direct relationship between energy and frequency. • Use the “Bohr Model for Hydrogen Atom” and “Electromagnetic Spectrum” diagrams from the Reference Tables to relate color, frequency, and wavelength of the light emitted to the energy of the photon. • Explain that Niels Bohr produced a model of the hydrogen atom based on experimental observations. This model indicated that: 1. an electron circles the nucleus only in fixed energy ranges called orbits; 2. an electron can neither gain or lose energy inside this orbit, but could move up or down to another orbit; 3. that the lowest energy orbit is closest to the nucleus. • Describe the wave/particle duality of electrons. Feb 18:45 AM We'll take these one at a time: Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Particle Symbol location relative relative charge mass Proton 1 1 H1+, 11P1+ Nucleus Plus 1 1 amu Neutron 1 n0 Nucleus Zero 1 amu Electron 0 Cloud Negative1 1/1857th amu 0 1 e1 Feb 18:46 AM Atomic Theory notes.notebook February 01, 2016 Next: Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Atomic number (Z) – the numbers of protons in the nucleus, identifies the element Mass number (A) – The total number of Protons and Neutrons in an element Feb 19:02 AM Next: Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Isotopes – atoms of the same element with different numbers of neutrons. Examples: protium (11H), deuterium (21H), tritium (31H) Ion – an atom that has gained or lost electrons. Lost electrons (cations) positive charge Gained electrons (anions ) negative charge Symbolic notation – shorthand method of describing a nuclide or ion Example – A nuclide with A = 17 (8p+ + 9n), Z = 8 (8p+), and charge = ‐2 (8p+ ‐ 10e‐) Feb 19:04 AM Atomic Theory notes.notebook February 01, 2016 Next: Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Isotope Worksheet Feb 18:48 AM Next: Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Atomic mass The atomic mass system is a relative system based on the mass of a carbon‐12 nuclide. The unit for measuring atomic mass is the atomic mass unit (amu). 1 amu = 1/12 the mass of a carbon‐12 nuclide. Atomic masses are specific to nuclides! Average atomic mass All naturally occurring elements exist as a combination of nuclides. For any element, the average atomic mass is an average of all the masses of all the isotopes in a naturally occurring sample of the element. Avg atomic mass = Σ (abundance in decimal form)(mass of nuclide) Example: Find the average atomic mass for boron, given that 10B is 19.781% and 11B is 80.219%. avg atomic mass = (0.19781)(10) + (0.80219)(11) = 10.812 amu Beanium Activity Feb 19:23 AM