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WVU CHEM 115 - Lecture 3

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Atomic Theory: History of the AtomAtomic Theory: experimental observations that led scientists to postulate the existence of the atom (smallest bit of an element).1. Law of Conservation of Mass -During a chemical reaction, mass is conserved. (A. Lavoisier 1743-1794)Ex. 71.8 g iron oxide decomposed to 55.8 g iron and 16.0 g oxygen.2. Law of Definite Proportions - In a given compound, the elements are always combined in the same ratio by mass. (Joseph Proust 1754-1826)Ex. 1.668 g molybdenum disulfide decomposed to 1.000 g Mo and 0.6680 g S.Mass Ratio Mo/S = 1.55.000 g molybdenum disulfide decomposed to 2.996 g Mo and 2.004 g S.Mass Ratio Mo/S = _____Dalton’s Atomic Theory of MatterJohn Dalton (1766-1844) proposed this theory to explain the experimental observations given by the laws of conservation of mass and definite proportions.Postulates:• Matter consists of tiny particles called atoms.• Atoms are indestructible. In chemical reactions, the atoms are rearranged but they do not themselves break apart.• Atoms of the same element are identical in mass and other properties.• Atoms of different elements differ in mass and other properties.• Chemical combination of elements to form compounds occurs. However, in a given compound the atoms of each element are present in a fixed number ratio.Law of Multiple ProportionsWhen two elements can form more than one compound, the mass ratios of the elements in the two compounds occur in small whole number ratios.Example: Consider two different compounds of iron and sulfur.A. Iron(II) disulfide (pyrite or “fool’s gold”):5.00 g pyrite decomposed to 2.67 g S and 2.33 g Fe.Mass Ratio S/Fe = 2.67 g/2.33 g = 1.15 g S/g FeB. Iron(II) sulfide:5.00 g iron(II) sulfide decomposed to 1.82 g S and 3.18 g Fe.Mass Ratio S/Fe = 1.82 g/3.18 g = 0.572 g S/g FeC. Law of Multiple Proportions:Mass Ratio S/Fe in pyrite:iron(II) disulfide = 1.15 g S/g Fe = 2.01 ~ 20.572 g S/g FeIn that Dalton’s Atomic Theory predicted the law of multiple proportions, this helped to give validity and force acceptance of the theory.Subatomic Particles: Particles Within the AtomThree key experiments helped to elucidate the structure of the atom and the nature of the subatomic particles.J.J. Thomson’s Experiments with Cathode Ray Tubes (1897)• Voltage applied across metal plates. Cathode ray obtained independent of metal used to make plates.• Cathode ray traveled from negative to positive.• Magnet deflected cathode ray. Used amount of deflection and magnetic field strength to calculate charge to mass ratio of cathode ray.Atom consists of parts, one of which is the electron.Electron is negatively charged.Charge to mass ratio of electron = 1.76x108C/g.R. Milliken’s Experiments with Oil Droplets (1909)• Milliken watched how fast oil droplets fell: calculated mass of oil droplets.• Placed negative charge on oil droplets using X-rays. Applied voltage to plates and suspended oil droplets in midair.• From mass of oil droplets and voltage needed to suspend, calculated charge on each oil droplet. Charge always some whole # multiple of 1.60x10-19 C.Fundamental charge on electron: 1.60x10-19CFrom e- charge and Thomson’s charge to mass ratio found mass of electron: 9.09x10-28gRutherford’s Experiments w/Alpha Particles (positive and 7000xheavier than e-)• Most of alpha particles went straight through metal foil.• 1/20,000 alpha particle deflected at large angles (repelled by something positive).• 1/20,000 alpha particle deflected straight back toward source (hitting something massive).Most of the atom is empty spaceIn center of atom there is a massive, positively charged corecalled the nucleus.Overview of subatomic particles• Atoms are composed of 3 subatomic particles, the proton, the neutron and the electron• Protons & neutrons have nearly the same mass and are located in a small nucleus in the atom• The electrons occupy most of the volume of the atom outside of the nucleus (atoms consist of mostly empty space)• Three Subatomic ParticlesParticle Rel. Charge MassElectron -1 9.109x10-28g (5.486x10-4amu)Proton +1 1.673x10-24g (1.007 amu)Neutron 0 1.675x10-24g (1.009 amu)• In a neutral atom, # protons = # electrons.• Isotopes: atoms of the same element but with different mass– same # protons (and same atomic number)– different # neutronsOverview of subatomic particlesAtomic Notation• An element is a substance whose atoms all contain the identical number of protons, called the atomic number (Z)─ Given as integer # above the element on periodic table.Ex. C Z=?Ca Z=?It is the # protons (or atomic #) that specifies the element.Ex. An element has 15 protons in the nucleus. What element is present?• Isotopes are distinguished by mass number (A):– Mass number, A = (number of protons) + (number of neutrons)– Not given on the periodic table, but can be calculated once # neutrons is known (A = # p + #n) or.. (A = Z + #n) or.. (#n = A – Z)Ex. Find # of subatomic particles present in boron-10 and boron-11.SyAZCarbon-12 (12C) Atomic Mass ScaleRelative atomic masses were not useful until a standard reference point was established. Atomic masses of all elements were referenced to the atomic mass of the most abundant isotope of carbon (12C).Atomic Mass Reference: Carbon-12 or 12C1 atom 12C = 12 amu (exactly)OR1 amu = 1/12 the mass of an atom of 12CDesignation was arbitrary but gave atomic masses close to whole numbers for most elements.Example: If the relative mass of Mo:12C is 7.995, what is the atomic mass ofMo on the12C atomic mass scale?620phosporusboron-10: A=10, Z=5,*p=5,*n=5boron-11: A=11, Z=5,*p=5,*n=6**Atomic masses on 12C atomic mass scale are shown as non-integer numbers below the elements on the periodic table.But…..Why is atomic mass of carbon given as 12.011 amu instead of as 12 amu?Atomic masses shown on periodic table are average atomic masses takinginto account the different isotopes of each element and their percent abundances.Isotopes are atoms of the same element but with a different mass. These isotopesoccur in different percentages in nature (percent abundances or isotopicabundances).Thus, the third postulate of Dalton’s Atomic Theory (Atoms of the sameelement are identical in mass) is NOT strictly true.Calculation of Average Atomic MassesExample: It is found that carbon consists of two naturally occurring isotopes(12C and 13C) with atomic masses and % abundances given below.


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