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Unit 3 – Atomic Structure Bravo – 15,000 kilotons DO NOW • How many moles are in 78.54 g Magnesium? (1 mole = 24.31 g Mg) • How many atoms are in 45.7 g of sodium? (1 mole = 6.02 x 1023 atoms; 1 mole = 22.99 g Na) Agenda • DO NOW • Atomic Structure History Democritus • 400 BC • Greek philosopher • 1st to come up with idea of atoms John Dalton – 1800’s • 1) 2) 3) 4) Major contributor of Atomic Theory All matter made of atoms All atoms of an element are alike Atoms cannot be created or destroyed Atoms combine in whole-number ratios to form compounds JJ Thomson – late 1800’s • Cathode Ray Experiment – discovery of electrons • “Plum Pudding” model of atom • Measured charge to mass ratio of e- Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. Conclusions from the Study of the Electron Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons Electrons have so little mass that atoms must contain other particles that account for most of the mass Millikan - 1910 Oil Drop Experiment Determined actual charge and mass of an e- Rutherford - 1910 Discovered nucleus and that it was positive Gold Foil Experiment 1) Most of atom is empty space (majority of particles went straight through) 2) nucleus is small, dense and positively charged (some positive charges were greatly deflected) Rutherford’s Gold Foil Experiment Alpha particles are helium nuclei Particles were fired at a thin sheet of gold foil Particle hits on the detecting screen (film) are recorded Rutherford’s Findings Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected Conclusions: The nucleus is small The nucleus is dense The nucleus is positively charged Niels Bohr - 1915 • Proposed early model of atom • “Planetary Model” electrons orbit nucleus like planets orbit sun • Lacks math of modern version • Has some errors/violates current theory • Radiation is emitted when electrons move from one orbit to another Chadwick Discovered neutrons Modern Atomic Theory (changes from Dalton) Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions! Atoms of an element have a characteristic average mass which is unique to that element (isotopes) Foundations of Atomic Theory Law of Conservation of Mass: mass is neither created or destroyed in ordinary chemical reactions Law of Definite Proportions (composition): compounds contain same elements in same ratio by mass Example: NaCl is always 39.9% Na and 60.66% Cl by mass Law of Multiple Proportions: 2 or more different compounds composed of same two elements have ratios of small whole numbers Example: CO vs CO2 ratio of oxygen to oxygen is 2 to 1 Book Assignment • Use sections 4.1 and 4.2 of the textbook to answer the following questions: – P.110 #1,4,5 – P.115 #10, 13-15 What is AMU? Stands for atomic mass unit – used when describing “relative” atomic masses This system is used because the actual masses of atoms are so small Carbon-12 is the standard to which all other elements are compared (i.e. hydrogen-1 has a mass that is 1/12 that of carbon-12 so it’s mass would be 1 amu) Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element # of protons Atomic # (Z) 6 6 Phosphorus 15 15 Gold 79 79 Carbon Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0 Nuclide p+ n0 e- Mass # Oxygen - 18 8 10 8 18 Arsenic - 75 33 42 33 75 Phosphorus - 31 15 16 15 31 Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011 Isotope Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 6 protons 7 neutrons 1.11% Carbon-14 14C 6 protons 8 neutrons <0.01% Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Hydrogen–1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2 Nucleus ISOTOPES • Example: • Carbon, C, exists in 3 isotopes: Isotope Protons Neutrons Mass Carbon-12 6 6 12 Carbon-13 6 7 13 Carbon-14 6 8 14 Isotope symbols Hyphen notation Nuclear notation Carbon – 12 12C 6 6p+ and 6no Carbon – 13 13C 6p+ and 7no Carbon – 14 14C 6p+ and 8no 6 6 How to Calculate the Average Mass What is the average atomic mass of sample of Cesium with 3 isotopes: 75% 133Cs, 20% 0.75 x 133 = 0.20 x 132 = 0.05 x 134 = Total = 132Cs, and 5% 134Cs. 99.75 26.40 6.70 132.85 avg. atomic mass The Mole 1 dozen = 12 1 gross = 144 1 ream = 500 1 mole = 6.02 x 1023 There are exactly 12 grams of carbon-12 in one mole of carbon-12. Avogadro’s Number 6.02 x 1023 is called “Avogadro’s Number” in honor of the Italian chemist Amadeo Avogadro (1776-1855). I didn’t discover it. Its just named after me! Amadeo Avogadro Calculations with Moles: Converting moles to grams How many grams of lithium are in 3.50 moles of lithium? 3.50 mol Li 6.94 g Li 1 mol Li = 45.1 g Li Calculations with Moles: Converting grams to moles How many moles of lithium are in 18.2 grams of lithium? 18.2 g Li 1 mol Li 6.94 g Li = 2.62 mol Li Calculations with Moles: Using Avogadro’s Number How many atoms of lithium are in 3.50 moles of lithium? 3.50 mol Li 6.022 x 1023 atoms Li 1 mol Li = 2.11 x 1024 atoms Li Calculations with Moles: Using Avogadro’s Number How many atoms of lithium are in 18.2 g of lithium? 18.2 g Li 1 mol Li 6.94 g Li 6.022 x 1023 atoms Li 1 mol Li (18.2)(6.022 x 1023)/6.94 = 1.58 x 1024 atoms Li