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Atomic Theory Chapter 3 As early as 400 BC scientists have believed in an atomic theory thanks to Democritus. n Atoms were the building blocks of matter. n 2000 years later we can see the atom! n Atoms History of the Atom n John Dalton—English teacher in 1808 developed five principles for his theory 4. Atoms of different elements combine in simple, whole­number ratios to form compounds 5. In chemical reactions, atoms are combined, separated, or rearranged but never created, destroyed, or changed. 1. All matter is composed of extremely small particles called atoms, which cannot be subdivided, created, or destroyed. 2. Atoms of a given element are identical in their physical and chemical properties. 3. Atoms of different elements are different in their physical and chemical properties. Structure & Models of Atoms n n n Law of definite proportions—two samples of a given compound are made of the same elements in exactly the same proportions by mass regardless of the sizes or sources of the samples. Law of conservation of mass—the mass of the reactants equals the mass of the products. Law of multiple proportions—if two or more different compounds are composed of the same two elements, the ratio of the masses of the second element is always a ratio of small whole numbers n n In the mid­1800s: subatomic particles! J.J. Thomson: studying electricity and discovered the electron. n Glass vacuum tube (cathode­ray tube ) and n applied voltage at both ends (electrodes). He noticed a glowing beam! The anode is the positive end and the cathode is the negative end.
1 magnet bent the beam Electrons are negatively charged and have mass. n Thomson’s model: “plum pudding” n n n Electrons embedded inside a dense mass. n n Ernest Rutherford In his lab at McGill University, 1903 1871­1937 n Rutherford disproved Thomson’s model in 1909. Gold Foil experiment directed alpha particles toward the foil. Some particles were deflected straight back, some at angles, and some went straight through. Rutherford hypothesized there was a dense mass called the nucleus that has to be larger than the alpha particle. Rutherford’s model
n n n n n Rutherford calculated the size of the nucleus to be 1/10 000 of the radius of an entire atom. Protons make up the positively charged nucleus. Chadwick found the neutrons in the nucleus—neutral charged Protons and neutrons have equal masses Coulomb’s Law—the closer two charges are, the more they attract. Like charges repel and opposites attract! 2 Neils Bohr Building a model helps us understand what is happening at the microscopic level. Rutherford’s model suggested that electrons revolve around the nucleus in circular or elliptical orbits. Bohr’s model says that electrons can only be a certain distance from the nucleus. The lowest energy level is by the nucleus. Bohr concluded that electrons in an energy level cannot give off energy. 1913 Studied under Rutherford at the Victoria University in Manchester. n Bohr refined Rutherford’s idea by adding that electrons were in orbits, like planets orbiting the sun. Each orbit can only contain a set number of electrons. n Steps to draw a Bohr Model Schrödinger's Quantum Theory n n n n n Determine the number of p, n, e (you must show this) n The modern­day model says that electrons are in orbitals—regions around a nucleus that corresponds to specific energy levels. n Orbitals are areas where electrons are likely to be found. n Draw nucleus–p&n densely packed (same as Rutherford model) n Electrons in circular energy levels around the outside of nucleus: up to 2 electrons in the first energy level up to 8 electrons in the second energy level n to 18 electrons in the third energy level n 32 electrons in the fourth energy level n 50 electrons in the fifth energy level n n n Atomic Number—the number of protons in the nucleus. n n n n Atomic numbers are always whole numbers If the atom is stable, that is, if it doesn’t have a charge, it is also the number of electrons. This identifies the element. n http://www.youtube.com/watch?v=otPvUTNstS4
Mass number—the SUM of the numbers of neutrons and protons. n Neutrons can vary from atom to atom but still have the same number of protons. These are called isotopes of each other. 3 The Atoms Family ­ Atomic Math Challenge Atomic Number Symbol Name Atomic Mass electrons protons Atomic number equals the number of ____________ or ________________. neutrons protons Atomic mass equals the number of ______________ + _______________. Isotopes n Atoms that have the same number of protons but different number of neutrons therefore having a different mass number! 1st Verse: They’re tiny and they’re teeny, Much smaller than a beany, They never can be seeny, The Atoms Family. Chorus 3rd Verse: Neutrons can be found, Where protons hang around; Electrons they surround The Atoms Family. Chorus 2nd Verse: Together they make gases, And liquids like molasses, And all the solid masses, The Atoms Family Chorus Chorus: They are so small. (Snap, snap) They’re round like a ball. (Snap, snap) They make up the air. They’re everywhere. Can’t see them at all. (Snap, snap) Element’s Symbols Each element has a name and a symbol n The atomic number and mass number are written with the element’s symbol. The atomic # appears on the lower left of the symbol and the mass # appears on the upper left of the symbol. n EX: Li Be S
n 4 Average Atomic Mass n Q: Why are a roll of pennies and dimes different masses when they each contain 50 coins? n Q: If you were given a mixed roll of coins, what information would you need to determine the mass of the roll? n n n Example n The masses on the PT are averages of atomic masses (amu—atomic mass unit) These masses take into account the relative abundance of each isotope—like your grade point average. General formula: Average atomic mass = (mass of X * X%) + (mass of Y * Y%) Solution Chlorine exists as chlorine­35, which has a mass of 34.969 amu and makes up 75.8% of chlorine atoms. The rest of naturally occurring chlorine is chlorine­37, with a mass of 36.996 amu. What is the average atomic mass of chlorine? Remember: Mass # = protons + neutrons Charge = protons – electrons Atomic # = number of protons AAM = (34.969 amu * .758) + (36.996 amu * .242) n AAM = 26.51 amu + 8.95 amu n AAM = 35.46 amu n Calculating Subatomic Particles
Mass # = protons + neutrons n Charge = protons – electrons n Atomic # = number of protons n ion – any atom with charge n Symbol Charge Atomic # Mass # Protons Neutrons Electrons 101 258 101 +1 40 ­2 Au +2 8 21 19 55 40 10 126 5