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Chapter 3: The Atom and the Mole (with nuclear) The investigation and understanding of the atom is what chemistry is all about! Topics rearranged from your text, pages 62-117. We come here to be philosophers, and I hope you will always remember that whenever a result happens, especially if it be new, you should say, “What is the cause? Why does it occur?” and you will, in the course of time, find out the reason. -Michael Faraday Bires, 2010 The Mole The “mole” represents a number of things….like a dozen. How many things is a mole? 6.022137 x 1023… we use 6.02 x1023. This is Avogadro’s number – named for a lawyer, Amadoe Avogadro, that studied molecular gasses as a hobby. When you have three moles of atoms, you have (3 x 6.02x1023 =) 1.81x1024 atoms total. Bires, 2010 Recall: Parts of the atom -1 (subatomic particles) +1 0 Proton charge = +1, mass = 1 Neutron charge = 0, mass = 1 Electron charge = -1, mass = 0 In a normal, neutral, unreacted atom, the number of electrons Ions have more or equals the number of protons. less electrons than protons Bires, 2010 They have a charge Atomic History Greek philosopher Democritus (400BC) – coined the term atomon which means “that which cannot be divided.” John Dalton (1803) a colorblind chemist. – Among his interests, Dalton was very interested in a scientific explanation for his colorblindness the behavior of gasses. In his A New System of Chemical Philosophy, Dalton published five principles of matter. Bires, 2010 Dalton’s Top Five All matter is made of indestructible and indivisible atoms. – (atoms are hard, unbreakable, the smallest thing there is) Atoms of a given element have identical physical and chemical properties. – (all atoms of X will behave the same anywhere) Different atoms have different properties. – (X behaves differently than Y) Atoms combine in whole-number ratios to form compounds. – (two H’s and one O = Water (H2O) Atoms cannot be divided, created or destroyed, – (just rearranged in chemical reactions). Bires, 2010 The Laws: Constant Composition – Ratios of atoms in a compound is constant for that compound. H 2O hydrogen-oxygen atomic ratio = 2:1 Conservation of Mass – Mass is not created or destroyed in a chemical reaction. Multiple Proportions: – Since atoms bond in small, whole number ratios to form compounds, the ratio of their mass ratios are small whole numbers. CO Oxygen-carbon mass ratio = 1.33 x2 Bires, 2010 CO2 Oxygen-carbon mass ratio = 2.66 Conservation of Mass Bires, 2010 Multiple Proportions Bires, 2010 The Cathode Ray Tube The cathode ray tube – A new invention suggested the presence of charges – areas of positive and negative…charge. – This suggested that atoms must be divisible, and Dalton’s theory had to be modified. Electrostatics J. J. Thomson (1897) – English Physicist proposed that the atom is a sphere of positive charge with small areas of negative charge. – This theory become known as the “plum pudding” model after an English “dessert” of purple bread and raisins. Bires, 2010 Millikan’s Oil Thompson used electrostatics experiments to determine the electron’s charge-to-mass ratio. Robert Millikan’s (1909) – oil-drop experiment allowed the charge of a single electron to be determined: 1.60 x 10-19 C. Scientists calculated the mass of an electron to be 1/2000 of the mass of a proton! Bires, 2010 Ernest Rutherford Ernest Rutherford (1910) – New Zealander Physicist, while studying radioactive elements, found that radioactive alpha particles deflected when fired at a very thin gold foil. The Gold Foil Experiment – the atom was not a hard sphere but – was mostly space, with a small concentration of positively-charged mass (the nucleus). Link Bires, 2010 to experiment Niels Bohr The Bohr Model – A Danish physicist (and student of Rutherford) rebuilt the model of the atom placing the electrons in energy levels. Bohr was one of the founders of quantum chemistry: – energy can be taken in and given off in small packets or quanta of specific size. – When a specific amount of energy was added to an atom, an electron could jump into a higher energy level. No more…no less! Bires, 2010 Adding the Neutrons James Chadwick (1932) – British physicist, proved there was too much mass in the nucleus – Suggested the existence of massive, neutral particles in the nucleus. (neutrons) Bires, 2010 The Modern Model Dalton’s atom electron Thompson’s electrons neutron Rutherford’s space and nucleus proton Bohr’s energy levels (not to scale) Bires, 2010 Chadwick’s neutrons Elements 112 known elements – 92 of which are naturally occurring. – 93 through 112: transuranium. Each has an atomic symbol. Atomic number 8 O OXYGEN 15.9994 – is number of protons Atomic mass – is the total mass of the protons plus the neutrons. Notice that the atomic mass is not a round number, even though protons and neutrons each have a mass of 1. This is due to natural abundance. Bires, 2010 Natural Abundance - Isotopes Isotopes: – Each element may have several isotopes – Isotopes differ in the number of neutrons. Example: – the element carbon has 6 protons, but it could have 5, 6, 7, or 8 neutrons, to form Carbon-11, Carbon12, Carbon-13, and Carbon-14. 11C, 12C, 13C, 14C In nature, there is a mix of different natural isotopes. We use this mix to calculate average atomic mass… Bires, 2010 Calculating Average Atomic Mass isotope 1 : relative abundance (%) x isotope mass (amu) isotope 2 : relative abundance (%) x isotope mass (amu) isotope 3 : relative abundance (%) x isotope mass (amu) Sum of the products = average atomic mass Example: – The isotopes of element “Bob” are found below: – Bob-18.0, 25.0% 0.25x18 0.60x19 0.15x20 – Bob-19.0, 60.0% 18 . 9 amu – Bob-20.0, 15.0% – What is the average atomic mass of naturally occurring Bob? Bires, 2010 1 amu = 1.66x10-27kg Isotopes – atoms of the same _______ – different number of _______ Review … Ions – atoms of the same _______ – different number of _______ Allotropes – forms of the same _______ – bonded in different _______ Quanta / Quantum – Packets of energy of _______ size Atomic Mass – Is the _______ of all _______ found in nature. Bires, 2010 Molar Mass Molar mass – expressed in grams per mole (g/mol) – mass of one mole of a substance. – link between the atom and the gram. No more AMU: AMU Molar mass (we can measure) The average atomic mass of carbon is 12.01. What is the mass of a mole of carbon atoms? What is the molar mass of Copper, Cu? Cu 63.5 g / mol What is the molar mass of Nitrogen, N2? Nx2 14.0 x 2 28.0 g / mol Bires, 2010 Molar Mass Practice Determine the Molar Mass of the following elements and compounds: Ca 40.41g / mol Ca Clx 2 35.5 x 2 71.0 g / mol Cl2 CaCl2 Ca Clx 2 40.1 35.5 x 2 111.g / mol H2O Hx 2 O 1x 2 16 18 g / mol Ba(OH)2 Ba (O H ) x 2 137 (16 1) x 2 171g / mol FeSO4 Fe S Ox 4 55.8 32.1 16 x 4 152 g / mol Al2(SO4)3 Alx 2 ( S Ox 4) x3 27 x 2 (32.1 16 x 4) x3 342 g / mol Bires, 2010 Mole-to-Mass Conversions Sodium has an atomic mass of 23 g/mol. How many moles do you have in 115 grams? Use a T-chart! 115 grams 23g / mol 5.0 moles How many grams are equal to 3.5 moles of CaCl2? 3.5moles 111g / mol 389 grams What is the mass of 0.46 moles of SiO2? .46mol 60.1g / mol 27.7 grams SiO2 Bires, 2010 Mole-Mass Conversion Practice Complete the following mole-to-mass conversions: Mass in grams of 2.25 moles of iron, Fe? 126 grams Fe Use your periodic table to find molar mass Mass in grams of 0.375 moles of potassium, K? 14.7 grams K Number of moles in 5.00 grams of calcium, Ca? 0.125 moles Ca Number of moles in 3.60x10-10 grams of gold, Au? 1.83x10-12 mol Au Bires, 2010 End of Chapter 3 Mole-Atoms Conversions Phew! Mole = 6.02x1023 things, how many atoms are in: 23 24 3 . 0 x 6 . 02 x 10 1 . 8 x 10 atoms 3.0 moles of silver, Ag? 21 0.010 moles of copper, Cu? 6.0 x10 atoms How many moles do you have in: 24 24 2 . 4 x 10 2.4x10 atoms of helium, He? 4.0 moles 6.02 x1023 3.0x1023 atoms of lithium, Li? 0.50 moles How many moles do you have in 222 grams of 222 grams / 63.55 g / mol 3.49 moles copper? How many atoms in 127.1 grams of copper? 127.1grams / 63.55g / mol 2.00 moles 6.02 x1023 1.20 x1024 atoms Bires, 2010 Isotopes – Nuclides - Radioactivity Nuclides mass – the nucleus of an isotope Place the mass above the charge as seen here. Nuclides undergo decay: charge – transformation into different nuclides – Balanced nuclear reactions – “Radioactive” – Half Life: time to decay ½ (mass) of a sample Bires, 2010 Images from ChemZone Alpha Decay Alpha Decay mass – a helium nucleus is released. Alpha particles: – move very slowly – because of their size, can be blocked with a few pages of paper or human skin – cause ionization (damaging!) – are positively charged Bires, 2010 charge This is a Nuclear Equation Alpha Decay occurs in all elements with atomic number above 83. Images from ChemZone Beta Decay Beta Decay – An electron is ejected from the nucleus Beta particles – move fast – can penetrate thick lowdensity materials – but can be blocked with concrete and metals Beta Decay occurs when – are negatively charged a nucleus has a high neutron-proton ratio. Bires, 2010 Images from ChemZone Gamma Decay Gamma Decay No mass – High energy photons (gamma rays) are given off. Gamma rays – given off as the “spare change” during other radioactive decays…. – extremely penetrating and powerful. Several inches of lead is required to slow these particles down to a stop. – Don’t get included in nuclear equations. No charge Summary of three basic particle decays Bires, 2010 Images from ChemZone Nuclear Equations Practice Sodium-24 undergoes alpha decay 24 11 Complete the following nuclear equations: Na ... 24 11 Na He ... 4 2 24 11 (help on click) Na He F 4 2 20 9 Iodine-131 undergoes beta decay I e Xe 131 53 0 1 131 54 Tungsten-190 undergoes alpha decay W He Hf 190 74 4 2 186 72 Uranium-238 undergoes alpha decay, then two beta decays (3 steps) U 24He 234 90Th 238 92 Bires, 2010 Th 10 e 234 91Pa 234 90 234 91 Pa 10 e 234 92 U Nuclear Fission Nuclear fission: – splitting of large, unstable atoms – releases large amounts of energy Critical mass (or critical density) – amount of fissionable fuel needed before reaction will begin. Once fission begins, it is Uncontrolled, nuclear fission proceeds to difficult to stop. completion with great speed. Nuclear Weapons: – two half-spheres of fissionable material are compressed together with conventional explosives, creating the critical mass. In order to harness nuclear fission to create useable electricity, we slow down the process with control rods… Bires, 2010 Nuclear Fusion Nuclear Fusion: – Joining of smaller nuclei to form larger nuclei. – Releases far more energy that nuclear fission. – Easier to control than fission. The sun’s (stars) energy comes from the fusion of hydrogen atoms into helium atoms. The H-bomb is a fusion weapon. Fusion: Bires, 2010 – The power supply of the future? – Why don’t we use it now? 2 1 H H He 2 1 4 2 E=mc2 Einstein: Energy and mass are interchangeable. – E = Energy – m = mass – c = speed of light ( 3 x 108 m/s ) – Very small amounts of mass create large amounts of energy! We use the formula ΔE= Δmc2 to build new, artificial elements in supercolliders (particle accelerators.) Fermilab cyclotron, Argonne National Laboratory, Chicago Bires, 2010