ATOMS AND ELEMENTS. Democritus 400 B.C. Atomic Theory of Matter. Dalton s Postulates (1803) Page 1

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1 ATOMS AND ELEMENTS Democritus 400 BC Believed that matter was composed of invisible particles of matter he called atoms According to Democritus, atoms could not be broken into smaller particles Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton Using two scientific laws discovered in the late 1700 s, Dalton built his atomic theory 1 Law of Conservation of Mass Antoine Lavoisier ( ) The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place 2 Law of Constant Composition Joseph Proust ( ) Also known as the law of definite proportions The elemental composition of a pure substance never varies Dalton s Postulates (1803) Dalton s Atomic Theory not only explained the law of conservation of mass and law of constant composition as they applied to the atom and their compounds, but also predicted the law of multiple proportions law of multiple proportions: If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers Page 1

2 1 Two different compounds are formed by the elements carbon and oxygen The first compound contains 429% by mass carbon and 571% by mass oxygen The second compound contains 273% by mass carbon and 727% by mass oxygen Show that the data are consistent with the Law of Multiple Proportions J J Thomson (1897) Using a cathode ray tube, he determined the charge-to-mass ratio for the electron as: 176 x 10 8 C/g The Atom, circa 1900: Milikan Oil Drop Experiment (1909) Plum pudding model, put forward by Thomson Positive sphere of matter with negative electrons imbedded in it Using voltage and change in the rate of fall of charged oil drops, he was able to determine the charge on each drop From Thompson s charge to mass ratio, Milikan determined the charge and mass of an electron Corrected to modern instrumentation: Mass of an electron x g Charge of an electron x C Radioactivity One of the pieces of evidence for the fact that atoms are made of smaller particles came from the work of Henri Becquerel and Marie Curie In 1892, Henri Becquerel discovered the spontaneous loss of nuclear energy from uranium salts This lead Marie Curie to the discovery of radioactivity, the spontaneous disintegration of some elements into smaller pieces Page 2

3 Ernest Rutherford Rutherford s Gold Foil Experiment (1910) Discovered alpha, beta and gamma radiation Results: 1 No holes 2 α- particles deflected at specific angles some backwards Rutherford (~1911) Nuclear Model Combination of Millikan s Findings and the Au Foil Experiment Lead to Rutherfords Model heavy central (+) nucleus e - about nucleus sea of e - James Chadwick (1932) Further developed the atomic model by theorizing that alpha and beta radiation results from the decomposition of a neutral particle found in the nucleus, the neutron H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4 So, what is a neutron? Possibly, a proton and electron held together by smaller particles called neutrinos (no charge) Neutrinos could account for the mass difference between a proton + electron and a neutron along with Electromagnetic Radiation The Bohr Model (1913) central (+) nucleus Planetary Model e - in allowed orbits Page 3

4 From Bohr s model, atoms could then be described as gaining charges As we have seen, electrons have a charge of x C Likewise, protons have a charge of x C for convenience, charges are reported as whole multiples of their charges or +1, -1; known as an atoms electronic charge Therefore, an atom that has lost an electron would have one more proton than electron and would have a net x C charge, or +1 charge Atoms with uneven numbers of electrons and protons are called ions anion (-) cation (+) Modern View of the Atom (mid 1920s) Heisenberg, debroglie, Schroediner e - in regions defined by math functions Quantum Mechanical Model Atomic Mass Because the masses of atoms are so small, the units of grams is much too large to be used conveniently Therefore, the Atomic Mass Unit (amu) is used The amu is defined by assigning a mass of exactly 12 to and atom of the carbon 12 isotope: 1 amu = x g We will revisit this shortly ATOMIC COMPOSITION Protons + electrical charge mass = x g relative mass = (amu) Electrons negative electrical charge mass = x g relative mass = amu Neutrons no electrical charge mass = x g mass = amu ATOM COMPOSITION The atom is mostly empty space protons and neutrons in the nucleus the number of electrons is equal to the number of protons electrons in space around the nucleus extremely small One teaspoon of water has 3 times as many atoms as the Atlantic Ocean has teaspoons of water Page 4

5 Atomic symbols Nuclear symbol - describes the number of particles in the nucleus of an atom Atomic # (Z) number of protons in the nucleus Mass # (A) A Z X total number of protons and neutrons in the nucleus Example: How many protons, neutrons and electrons are in the following atom? Na Na Na 23 Hyphen notation Isotopes Atoms of the same element (same Z) but different mass number (A) Boron-10 ( 10 B) has 5 p and 5 n Boron-11 ( 11 B) has 5 p and 6 n Masses of Isotopes determined with a mass spectrometer 11 B 10 B Bone scans with radioactive technetium-99 Average Mass Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations Average mass is calculated from the isotopes of an element weighted by their relative abundances Boron is 20% 10 B and 80% 11 B That is, 11 B is 80 percent abundant on earth For boron, its atomic weight = 020 (10 amu) (11 amu) = 108 amu The average atomic mass is found by calculating the mass in grams of each isotope in relationship to its naturally occurring abundance Example: A sample of chlorine gas is 7553% 35 Cl 2447% 37 Cl Page 5

6 Atomic Mass Average Atomic Mass (3545) Remember, the amu is defined by assigning a mass of exactly 12 to and atom of the carbon 12 isotope Therefore, 1amu = 1/12 the mass of a Carbon-12 isotope, or 1/12 the mass of 6 neutrons and 6 protons (electrons are negligible) Try it and see if you get the same number as below: 1 amu = x g What is the mass in amu of 1 Gram of matter? Average Atomic Mass vs Average Atomic Weight The atomic weight describes the average mass of the naturally occurring isotopes multiplied by their relative percentages in atomic mass units The average atomic mass describes the average mass of the naturally occurring isotopes multiplied by their relative percentages in grams Periodic Table Dmitri Mendeleev developed the modern periodic table He argued that element properties are periodic functions of their atomic masses We now know that element properties are periodic functions of their ATOMIC NUMBERS Groups vs Periods Regions of the Periodic Table and Element Abundance Hydrogen Shuttle main engines use H 2 and O 2 The Hindenburg crash, May 1939 Page 6

7 Group 1A: Alkali Metals Group 2A: Alkaline Earth Metals Magnesium Reaction of potassium + H 2 O Cutting sodium metal Magnesium oxide Group 3A: B, Al, Ga, In, Tl Group 4A: C, Si, Ge, Sn, Pb Quartz, SiO 2 Aluminum Boron halides BF 3 & BI 3 Diamond Group 5A: N, P, As, Sb, Bi Group 6A: O, S, Se, Te, Po Ammonia, NH 3 White and red phosphorus Sulfuric acid dripping from snot-tite in cave in Mexico Sulfur from a volcano Page 7

8 Group 7A: F, Cl, Br, I, At Group 8A: He, Ne, Ar, Kr, Xe, Rn Lighter than air balloons Neon signs XeOF 4 Transition Elements Colors of Transition Metal Compounds Lanthanides and actinides Iron in air gives iron(iii) oxide Iron Cobalt Nickel Copper Zinc Page 8

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