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Unit B: Atomic Structure & Nuclear Chemistry Atoms: Building Blocks of Elements • An atom is the smallest particle of an element that has the characteristics of that element. The Atom Proton Electron Neutron Subatomic Particles • Proton+ has a positive charge and is located in the nucleus of an atom • Neutron0 has no charge and is also located in the nucleus • Electron- has a negative charge and travels at the speed of light around the outside of the atom The Nucleus • Is positively charged because it contains protons (p+) and neutrons (n0) • Is much smaller than the whole atom yet contains most of its mass This is what the atomic model would look like drawn to scale History of the Atom • In ~400 BC, Democritus proposed that everything in the universe was composed of small, indivisible objects called “atomos” John Dalton • In the 1800’s, John Dalton performed a series of experiments combining elements that showed that matter actually did consist of small particles, but he couldn’t predict their structure • He used this to come up with his Atomic Theory Dalton’s Atomic Theory 1) chemical elements are made of atoms. 2) the atoms of an element are identical in their masses 3) atoms of different elements have different masses 4) atoms only combine in small, whole number ratios (there are no ½ atoms) 5) atoms can be neither created nor destroyed Evaluate • Which aspects of Dalton’s Atomic Theory were correct? Which were not? J.J. Thomson • In 1897, another English scientist, J.J. Thomson, discovered the electron • Based upon his experiments, he proposed the “Plum Pudding” model of the atom, with negative electrons floating in a sea of positive matter Ernest Rutherford • In 1911, Ernest Rutherford performed experiments that led him to discover that electrons orbit a positively charged nucleus • He later discovered it contained protons and neutrons His hypothesis was incorrect…and he learned something! Rutherford “planetary” model Niels Bohr • Danish scientist Niels Bohr came up with the final model of the atom • He discovered that electrons can only orbit the nucleus at certain specific distances, which he called “orbitals” or “energy levels” Energy Levels • Bohr discovered that different orbitals required different amounts of energy for an electron to reach them • Electrons could not exist “between” levels, but they could leap from one level to the next if charged with enough energy • Each level could only hold a certain number of electrons (2, 8, 18, etc) The Bohr Model Present – Quantum Mechanical Model (Heisenberg/Schrodinger) • Bohr and other scientists (Heisenberg) updated the model in light of quantum mechanics • Rather than treating electrons as having specific locations and circular orbits, they are represented by probability clouds where electrons are most likely to be found Not THAT Heisenberg… This one. History of Atomic Theory Summary Scientist/Philosopher Discovery/Contribution Democritus Atoms are “uncuttable” John Dalton Atomic Theory; elements are composed of one type of atom JJ Thomson Electron Ernest Rutherford Nucleus Neils Bohr Energy Levels Werner Heisenberg/Schrodinger Electron Clouds This is what they look like Elements • Atoms of different elements differ in their number of protons • All known 118 elements in the universe are found on the Periodic Table of Elements Chemical Symbols • Each element has its own unique symbol which always begins with a CAPITAL letter Atomic Number • Is the number of protons in an atom • In a neutral atom, it is also the number of electrons Atomic Mass • Is the average weight of an atom based upon the number of protons and neutrons in the nucleus • Measured in atomic mass units (amu) • Tells you the mass number of MOST atoms of an element Mass Number • Is the total number of protons and neutrons in an atom. • To find the typical mass number of an element, round its atomic mass to the nearest whole number • The mass number for Hydrogen is 1 amu Simplified for notes • Atomic # = p+ • e- = p+(if neutral) • Mass # = p+ + no • # of neutrons = mass # - atomic # Examples • • • • • • Element name: Carbon Atomic Number: 6 Mass Number: 12 # Protons: 6 # Electrons: 6 # Neutrons: 12-6 = 6 Another... • Element name: Gold • Atomic Number: 79 • Mass Number: 197 • # Protons: 79 • # Electrons: 79 • # Neutrons: 197-79 = 118 Isotopes • -are atoms of an element that have the same number of protons but different numbers of neutrons • E.g.: • -Carbon-12 has 6 protons, 6 neutrons • -Carbon 13 has 6 protons, 7 neutrons Representing Isotopes This is the isotope needed for nuclear fission NUCLEAR CHEMISTRY Marie Curie • In 1896, Marie Curie demonstrated that the fogging of film plates exposed to uranium was caused by rays emitted by the atoms • She called this radioactivity and became the first woman to receive a Nobel Prize in 1903 Radioactive Isotopes • Are atoms with an imbalance of protons to neutrons and an unstable nucleus • To become stable, they slowly give off excess particles in the form of radiation • Any atom that is giving off radiation is said to be radioactive Radiation • The penetrating rays and particles emitted by a radioactive source is known as radiation • Radiation that is harmful to the body is known as ionizing radiation • It can take 3 forms: alpha, beta, and gamma Alpha Radiation (α) • Consists of Helium nuclei that have been emitted from a radioactive source • Because it is just the nucleus (no electrons), these particles have a positive charge Beta Radiation (β) • Consists of fast moving electrons formed by the decomposition of a neutron • The neutron breaks into a proton, which stays in the nucleus, and an electron which is emitted Gamma Radiation • Is high-energy electromagnetic radiation (light) given off by a radioisotope alongside alpha or beta particles • Most damaging type of radiation Gamma ray bursts like this one could wipe out all life on Earth at any moment. Property Alpha Radiation Beta Radiation Gamma Radiation Particle Helium Nucleus Electron High-energy electromagnetic radiation Symbol α, 42He β, 0-1e γ Charge Positive (+2) Negative (-1) 0 Common Source Radium-226 Carbon-14 Cobalt-60 Penetrating Power Low (0.05 mm body tissue) Moderate (4 mm Very High body tissue) (penetrates body easily) Shielding Paper, clothing Metal foil Lead, concrete Radiation Penetration Half-life • Radioactive isotopes decay at very specific, unique rates • The time required for one-half of the nuclei of a sample to decay is called its half-life Half Life Problems • Each time a sample goes through 1 half life, half of the sample decays and half remains. • Ex: Let’s say a sample has a half life of 5 years and we start with 8g… 8g remain Year 0 (0 half lives) 4g remain 2g remain 1g remains Year 5 (1 half life) Year 10 (2 half lives) Year 15 (3 half lives) Note: will the sample ever reach 0? No! Graphing Half Lives E=mc2 • In the early 1900’s, Albert Einstein showed that energy and matter were equivalent with his famous equation, E=mc2 • This lead later scientists to hypothesize that tremendous amounts of energy could be produced by tiny amounts of matter Nuclear Fission • The splitting of an atomic nucleus is known as nuclear fission • The split atoms then release more neutrons, leading to a chain reaction Nuclear Chain Reactions Controlled (nuclear power plant) Uncontrolled (atomic bomb) Nuclear Fusion • Occurs when nuclei combine to produce a bigger nucleus • For example, 2 Hydrogen nuclei can combine to form Helium • The sun is powered by nuclear fusion Tsar Bomba (Hydrogen bomb) – the biggest bomb ever A nuclear fission reaction sets off a second nuclear reaction, this one an uncontrolled fusion reaction, resulting in incredible amounts of energy COUNTING ATOMS Avg Atomic Mass • Is the weighted average of all of the isotopes of an element • To find the mass #, multiply the % abundance of each isotope by the isotope’s mass and then add the products together Avg Atomic Mass of Carbon • 0.989 x 12 = 11.868 • 0.011 x 13 = 0.143 • 0.000001 x 14 = 0.000014 Atomic Mass = 11.868 + 0.143 + 0.000014 = 12.01 amu Avogadro's Number (the Mole) • A unit used for counting atoms • Similar to a dozen, except instead of 12, it’s 602 billion trillion 602,000,000,000,000,000,000,000 • 6.02 X 1023 (in scientific notation) • This number is named in honor of Amedeo _________ (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the same number of molecules present in a given volume (Avogadro’s Number) The mole defined… • The mole is defined as the number of carbon12 atoms in exactly 12g of carbon • This number turns out to be Avogadro’s Number, 6.02 x 1023 (scientists have actually calculated this down to 8 significant figures, but we will only hold you to the 3 ) Just How Big is a Mole? • Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles. • If you had Avogadro's number of unpopped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. • If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole. • In other words: a LOT. Everybody Has Avogadro’s Number! But Where Did it Come From? • It was NOT just picked! It was MEASURED. • One of the better methods of measuring this number was the Millikan Oil Drop Experiment • Has been consistently supported by repeated experiments Why do scientists use the mole? • Used for counting extremely large numbers, such as the number of atoms in 18g of water • Similar to why bakers use dozens Counting units Suppose we invented a new collection unit called a rapp. One rapp contains 8 objects. 1. How many paper clips in 1 rapp? a) 1 b) 4 c) 8 2. How many oranges in 2.0 rapp? a) 4 b) 8 c) 16 3. How many rapps contain 40 gummy bears? a) 5 b) 10 c) 20 The Mole • 1 dozen cookies = 12 cookies • 1 mole of cookies = 6.02 X 1023 cookies • 1 dozen cars = 12 cars • 1 mole of cars = 6.02 X 1023 cars • 1 dozen Al atoms = 12 Al atoms • 1 mole of Al atoms = 6.02 X 1023 atoms Note that the NUMBER is always the same, but the MASS is very different! (a mole of cars will probably weigh more than a mole of cookies!) Mole is abbreviated mol (gee, that’s a lot quicker to write, huh?) A Mole of Particles Contains 6.02 x 1023 particles 1 mole C = 6.02 x 1023 C atoms 1 mole H2O = 6.02 x 1023 H2O molecules 1 mole NaCl = 6.02 x 1023 NaCl “molecules (technically, ionic compounds are ratios not molecules so they are called formula units) 6.02 x 1023 Na+ ions and 6.02 x 1023 Cl– ions Avogadro’s Number as Conversion Factor 6.02 x 1023 particles 1 mole or 1 mole 6.02 x 1023 particles Note that a particle could be anything since the mole is a counting unit Converting from moles to particles • If I have 2 moles of pennies, how many pennies do I have? 2 moles X 6.02 x 1023 pennies = 12.04 x 1023 of pennies mole pennies pennies = 1.20 x 1024 pennies Particles to moles • If I have 1.8 x 1024 pencils, how many moles of pencils do I have? 1.8 x 1024 pencils x 1 mole of pencils 6.02 x 1023 pencils = 3 moles of pencils Learning Check 1. Number of atoms in 0.500 mole of Al a) 500 Al atoms b) 6.02 x 1023 Al atoms c) 3.01 x 1023 Al atoms 2.Number of moles of S in 1.8 x 1024 S atoms a) 1.0 mole S atoms b) 3.0 mole S atoms c) 1.1 x 1048 mole S atoms Part 2: Molar Mass • Explain what molar mass is • Distinguish between gram molecular mass and gram formula mass • Find the molar mass of elements and compounds Molar Mass • The mass (in grams) of 6.02 x 1023 particles (1 mole) of a substance • Relative in value to the avg atomic mass: 1 C atom = 1 mole of C atoms 1 Mg atom atoms 12 amu = 12 g = 24 amu 1 mole of Mg = 24 g Same number, different unit! The Unit • We express molar mass using the unit g/mol(grams per mole) • Similar to density (g/mL) Atomic Mass and Molar Mass Atomic Mass • the mass of one ATOM • Measured in amu (atomic mass units) Molar Mass • The mass of one MOLE of atoms (6.02 x 1023 atoms) • Measured in grams (g) per mole The atomic mass and molar mass are the same number…the only thing that is different is the unit! Formula/Atomic/Molecular Mass • Gram atomic mass (gam) is the mass of 1 mole of atoms • Gram Molecular Mass (gmm) is the mass of 1 mole of molecules (covalent only!) • Gram Formula Mass (gfm) is the mass of 1 mole of ionic compounds These are all examples of the same thing: molar mass! Learning Check! Find the molar mass (usually we round to the hundredths place) A. 1 mole of Br atoms = 79.90 g/mole B. 1 mole of Sn atoms = 118.71 g/mole Molar Mass of Molecules and Compounds Mass in grams of 1 mole equal numerically to the sum of the atomic masses 1 mole of CaCl2 = 110.98 g/mol 1 mole Ca x 40.08 g/mol + 2 moles Cl x 35.45 g/mol 1 mole of N2O4 = 110.98 g/mol CaCl2 = 92.02 g/mol Learning Check! A. Molar Mass of K2O = ? Grams/mole 94. 20 g/mol B. Molar Mass of antacid Al(OH)3 = ? Grams/mole 78.01 g/mol Challenge Prozac, C17H18F3NO, is a widely used antidepressant that inhibits the uptake of serotonin by the brain. Find its molar mass. Pt 3: Calculations with Molar Mass molar mass Grams Moles • Moles to Grams/Grams to Moles • Particles to Grams/Grams to Particles • Volume Converting Moles and Grams Aluminum is often used for the structure of light-weight bicycle frames. How many grams of Al are in 3.00 moles of Al? 3.00 moles Al ? g Al 1. Molar mass of Al 1 mole Al = 26.98 g Al 2. Conversion factors for Al 26.98g Al 1 mol Al or 1 mol Al 26.98 g Al 3. Setup 3.00 moles Al x 26.98 g Al 1 mole Al Answer = 80.94 g Al Example 2 After leaving his bike out all night, Mr. Pinson collected a 36g sample of Iron (II) Oxide from his wheel. How many moles of rust were formed? 1. First, find the molar mass of FeO Fe 1 55.85 O 1 16.00 = 71.85 g/mol 2. Write the known quantity 36g FeO x 1 mole FeO 71.85g FeO 3. Multiply by the conversion factor. = 0.50 Moles FeO