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CHAPTER #2 - Atoms, Molecules, and Ions 2.1 The Early History of Chemistry Ancient Greeks - thought matter was composed to 4 substances – earth, air, fire, water (_________ vs. _________) Alchemy - ______________________________ (discovered many elements; learned to prepare mineral acids) Metallurgy - extraction of metals from ore Robert Boyle - (1661 - Skeptical Chymist) - first quantitative experiments; current concept of “element” Georg Stahl – suggested “phlogiston” flowed out of burning material Joseph Priestley - (1733-1804) “discovered” oxygen (not phlogiston) 2.2 - 2.3 Fundamental Chemical Laws/ Dalton’s Atomic Theory NATURAL LAWS: 1. Conservation of Mass - In any chemical reaction, mass is neither created nor destroyed 2. Constant Composition – a given compound always contains exactly the same proportion of elements by mass. 3. Multiple Proportions - When two elements form a series of compounds, the ratios of the masses of the 2nd element that combine with 1 gram of the 1st element can always be reduced to small whole numbers. Mass Oxygen w/ 1 g C Cpd I 1.33 g Cpd II 2.66 g Mass Nitrogen w/ 1 g O Cpd I 1.750 g Cpd II 0.8750 g Cpd III 0.4375 g John Dalton (1808) “Father of Atomic Theory” Essentials of his theory. . . 1. An element is composed of tiny particles called atoms. All atoms of a given element show the same chemical properties. 2. Atoms of different elements have different properties. In an ordinary chemical reaction, no atom of any element disappears or is changed into an atom of another element. 3. Compounds are formed when atoms of two or more elements combine. In a given compound, the relative numbers of atoms of each kind are definite and constant. In general, these relative numbers can be expressed as integers or simple fractions. IN GENERAL Elements consist of tiny particles called _________, which retain their identity in ____________________. In a compound, atoms of two or more elements are combined in a fixed ratio of ___________________. Gay-Lussac - (1809) Combining Volumes of Gases • Performed experiments in which he measured (under same T & P) the ___________________ that reacted with each other. Amadeo Avogadro - (1811) “Avogadro’s Hypothesis” • Equal volumes of different gases contain the ______________________ at the same temperature and pressure. 2.4 Early Experiments to Characterize the Atom JJ Thomson (1856 - 1940): English physicist ____________ Experiment (late 1890’s) Important Observations Inferences This lead to the _____________ Model Robert Milikan (1868 - 1953) _____________ Experiment (1909) Important Observations Inferences Sir James Chadwick (1891-1974): confirmed existence of ____________ Henri Bequerel (1896) French Chemist essentially discovered _________________ (alpha, beta, and gamma particles discovered later) Ernest Rutherford (1911) ________________ experiment Important Observations Inferences 2.5 - 2.6 Modern View of Atomic Structure/ Molecules and Ions COMPONENTS Proton Neutron Electron Nuclear Symbol Relative Mass _____ _____ _____ Relative Charge _____ _____ _____ A ZX Location _______ _______ _______ Actual Mass Actual Charge 1.6727 x 10-24 g 1.6022 x 10-19 C 1.6750 x 10-24 g __0____ -1.6022 x 10-19 C 9.1095 x 10-28 g (X = element symbol) • Z = Atomic Number = number of protons and is UNIQUE • A = Mass Number = sum of neutrons and protons is NOT unique • Atoms and Molecules are electrically neutral ---> # protons = # electrons Ions - Have an electrical charge due to gain/loss of electrons cations anions isoelectric – K+, Ca2+, P3-, S2-, Cl-, Ar Isotopes Isotope Mg-24 Mg-25 Mg-26 Mass (amu) 23.9850 24.9858 25.9826 Percent (Abundance) 78.99% 10.00% 11.01% Determine the relative average atomic weight of Magnesium. COMPLETE THE FOLLOWING TABLE: 31 # electrons # protons # neutrons 15P _____ _____ _____ 40 Ca2+ _____ _____ _____ 74 As3_____ _____ _____ 38 F2 _____ _____ _____ The Periodic Table Naming – Ionic Cpds, Molecular Cpds, & Acids – see next H3O+ _____ _____ _____