Chapter 2: Atoms and the Periodic Table Book: 2.1 & 2.2 I. Progression of Atomic Theory 1. Democritus and Leucippus (5th century B.C.) Matter is composed of small, indivisible particles called atomos (“uncuttable”) 2. Aristotle and Plato They didn’t agree with Democritus and the idea was rejected for 2,000 years 3. John Dalton (1766 – 1844) English schoolteacher and quaker Began teaching at age 12 Kept a diary of meteorological observations (>200,000 entries) Laid in state in Manchester >40,000 people showed up Dalton’s Postulates Each element is composed of extremely small particles called atoms All atoms of a given element are identical (not accepted) Atoms of an element are not changed into a different type of element by chemical reactions; atoms are neither created nor destroyed in chemical reactions (not accepted) Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms 4. J. J. Thomson (1856 – 1940) Won the Nobel Prize for physics in 1906 (his son won the same in 1937) Discovered the electron in 1897 (first part of the atom that is discovered) Proved in cathode ray experiments that there are particles smaller than an atom Discovered that electrons are 2,000 times lighter than hydrogen so he proved that there are tiny particles inside an atom that are negatively charged Thomson’s Plum Pudding Model of the Atom Overall, atoms are neutral But atoms are made up of positively charged matter with negatively charged electrons embedded throughout 5. R. A. Millikan (1868-1953) American physicist Conducted the oil drop experiment Awarded the Nobel Prize (1911) for determining the charge on an individual electron Thomson + Milliken = the ability to calculate the mass of a single electron (9.10x10-28 g) 6. Ernest Rutherford (1871 – 1937) Graduate research student for J.J. Thomson Won the Nobel Prize in Physics (1908) Used the gold foil test to develop a new model of the atom Concluded that atoms have a nucleus A nucleus is very small, very dense, and positively charged Rutherford’s Gold Foil Experiment Sample inside of lead box with a single hole Radioactive sample emits beam of alpha particles Most alpha particles go straight through the gold, showing that gold atoms were mostly empty space Occasionally an alpha (positively charged) particle would beam off, proving the presence of a small nucleus The Neutron mass of electrons + mass of protons + mass of neutrons = atomic mass James Chadwick proved the existence of neutrons (1932) Bombarded a sheet of Be with alpha particles which gives off neutrons as a byproduct He proves they were neutral but had a mass similar to protons II. The Parts of the Atom Charge Mass (amu) Protons +1 1.0073 ~1 Neutron 0 1.0087 ~1 Electron -1 5.534 x 10-4 The Nucleus Contains the protons and neutrons Tiny (nucleus is ~10-4 A, and the whole atom is 1-5 A) If an atom was the size of a football stadium, the nucleus would be a blueberry sitting on the field Extremely dense (1013-1014 g/mL) A matchbox full of nuclei would weigh over 2.5 billion tons Protons and Neutrons: 1. Make up almost none of the volume 2. Make up about half of its mass (each) 3. Involved in nuclear reactions only Electrons: 1. Make up most of the volume of the atom 2. Make up almost none of the mass 3. Very involved in the atom’s reactions All atoms of an element have the same number of protons (ex: all carbon atoms have 6 protons) If an atom is neutral, the # protons = # electrons # neutrons fluctuates III. Abbreviations A ZX Atomic number (Z) = subscript = # of protons Mass number (A) = superscript + # protons + # neutrons 52 24Cr ________ protons ________ neutrons ________ electrons 23 11Na + ________ protons ________ neutrons ________ electrons 19 9F ________ protons ________ neutrons ________ electrons IV. Isotopes Isotopes are atoms of the same element that have a different # of neutrons Atomic Mass (g/mole) Atomic Structure Relative Abundance 24 Mg 23.98504 12p, 12e, 12n 78.99% 25 Mg 24.98584 12p, 12e, 13n 10.00% 26 Mg 25.98259 12p, 12e, 14n 11.01% Q: What is the average atomic mass of Mg? (23.98504)(0.7899) + (24.98584)(0.10) + (25.98259)(0.1101) = 24.3050 g/mole V. Dmitri Mendeleev (1834 – 1907) Proposed the modern periodic table Table is organized by increasing atomic number (# of protons) Columns are called groups Elements in a group often exhibit similar properties Rows signify an additional valence shell 3 Parts: Metals, nonmetals, metalloids Group Names of the Periodic Table Group 1A: Alkali Metals Group 2A: Alkaline Earth Metals Group 7A: Halogens Group 8A: Noble Gases Groups 3B – 2B: Transition Metals Atoms 57-70: Lanthanides Atoms 89-102: Actinides Amedeo Avogadro (1776 – 1856) Italian savant Resolved the differences between atoms, molecules, and gases Most noted for contributions in the theory of molarity, molecular weight, and Avogadro’s Law The Mole A unit of counting, like the dozen 1 dozen = 12 somethings 1 gross of pencils = 144 pencils (a dozen dozens) 1 ream of paper = 500 sheets of paper 1 mole = 6.02 x 1023 atoms/molecules/ions/whatevers (Avogadro’s number) Q: How many atoms are in 5.18g of P? Q: A pure titanium cube has an edge length of 2.78 inches. How many Ti atoms does it contain? d = 4.50 g/cm3. Q: Two samples of carbon tetrachloride (CCl4) were decomposed into their constituent elements. One sample produced 38.9g of C and 448g of Cl. The other sample produced 14.8g of C and 134g of Cl. Are these results consistent with the law of definite proportions?
