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Chapter 2 Atoms, Molecules and Ions • 2.1 The Atomic Theory of Matter • 2.2 The Discovery of Atomic Structure • 2.3 The Modern View of Atomic Structure • 2.4 The Periodic Table • 2.5 Molecules and Molecular Compounds • 2.6 Ions and Ionic Compounds • 2.7 Naming Inorganic Compounds 2.1 The Atomic Theory of Matter • • • • • Greeks: 4 “elements” Democritus and Leucippus - atoms Aristotle- elements Alchemy 1660 - Robert Boyle- experimental definition of element. • Lavoisier- Father of modern chemistry. 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. © 2009, Prentice-Hall, Dalton's Postulates 1. Each element is composed of extremely small particles called atoms. © 2009, Prentice-Hall, Dalton's Postulates 2. All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. © 2009, Prentice-Hall, Dalton's Postulates 3. Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. Law of Conservation of Mass 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. © 2009, Prentice-Hall, Dalton’s Postulates 4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. Law of Constant Composition (also known as Proust’s Law or the law of definite proportions) It states that the elemental composition of a pure substance never varies. © 2009, Prentice-Hall, Inc. • Modern Version of the Experiment: Using electricity, water is electrolyzed into its components, hydrogen and oxygen. • Results: Water is always found to contain hydrogen and oxygen in a mass ratio of 1:8 (volume ratio of 2:1) no matter how it is formed or where it is found. 2.2 The Discovery of Atomic Structure The Electron Voltage source In 1897 J.J. Thomson discovered that passing an electric current through an evacuated sealed glass tube, makes a stream of particles move from the negative to the positive end. © 2009, Prentice-Hall, Inc. Thomson’s cathode ray tube experiment Voltage source + - By adding an electric field, he found that the moving pieces were negative. Thompson measured the charge/mass ratio of the electron to be 1.76 108 coulombs/g. Plum Pudding Model The atom was thought to be a positive sphere of matter with negative electrons imbedded in it. © 2009, Prentice-Hall, Robert Milikan’s oil drop experiment Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other. Milikan sprayed oil drops above above a charged plate containing a small hole. As the oil drops fall through the hole, they are given a negative charge. When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate. Using this experiment, Millikan determined the charge on the electron to be 1.60 10-19 C. Henri Becquerel discovers radioactivity Radioactivity is the spontaneous emission of radiation by an atom. Also studied by Marie and Pierre Curie. Ernest Rutherford noted three spots on the detector: a spot in the direction of the positive plate ( particles) a spot which is not affected by the electric field ( rays) a spot in the direction of the negative plate ( particles) Rutherford’s gold foil experiment A source of α-particles was placed at the mouth of a circular detector. The α -particles were shot through a piece of gold foil. Rutherford’s nuclear atom Since some particles were deflected at large angles, but most went straight through, Thompson’s model could not be correct. Evidence Theory 1) Some alpha particles deflected greatly presence of very small, dense positive nucleus (with the electrons around the outside) 2) Most alpha particles went straight through most of the volume of the atom is empty space. © 2009, Prentice-Hall, Inc. Niels Bohr • Electrons travel in orbit around nucleus • Idea of electron transitions between energy levels, emitting/absorbing photons Bohr’s energy levels • The energy levels in an atom are sort of like steps of a ladder. • The more energy an electron has, the farther away from the nucleus it usually will be. • The energy levels are not evenly spaced. They get closer together as you travel farther away. • To move from one “rung” to another requires a “quantum” of energy. Bohr Atomic Model The difference between continuous and quantized energy levels. continuous energy levels quantized energy levels James Chadwick discovers the neutron A thin sheet of beryllium was bombarded with alpha particles, and a very high energy radiation was emitted. The radiation was electrically neutral, had no charge but a slightly larger mass than a proton. Heisenberg’s Uncertainty Principle The Modern View Quantum Mechanical Model Quantum Mechanical Model 2.3 The Modern View of Atomic Structure Subatomic particles We define: mass of 12C = exactly 12 amu. Using atomic mass units: 1 amu = 1.66054 x 10-24 g 1 g = 6.02214 x 1023 amu Symbols of Elements Elements are symbolized by one or two letters. First letter is always capital. Second letter (if there is one) is lower case. © 2009, Prentice-Hall, Inc. Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z) © 2009, Prentice-Hall, Inc. Mass Number The mass of an atom in atomic mass units (amu) is the total number of nucleons (protons + neutrons) The mass number (A) By convention, for element X, we write ZAX. © 2009, Prentice-Hall, Inc. Isotopes • Isotopes are atoms of the same element with different masses. • In terms of subatomic particles, isotopes have different numbers of neutrons. 11 C 6 12 C 6 13 C 6 14 C 6 © 2009, Prentice-Hall, Inc. Average Atomic 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. • Average atomic mass = Σ (isotope mass)( relative abundance) © 2009, Prentice-Hall, Inc. Atomic Mass Calculation 1 Isotope Atomic Mass Relative Abundance 28Si 27.976 92.2297% 29Si 28.976 4.6832% 30Si 29.973 3.0872% • Using the isotope information for Silicon. Find the average atomic mass. Atomic Mass Calculation 2 Silver consists of two isotopes 107Ag and 109Ag. Its average atomic mass is 107.87. Calculate the percentage of each isotope in naturally occurring silver. (Assume that the masses are 107.00 and 109.00 respectively.) Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer. © 2009, Prentice-Hall, 2.4 The Periodic Table Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. © 2009, Prentice-Hall, Inc. Periodic Table Metals are on the left side of the chart. • • • • Good conductors Lose e- to form cation Malleable Ductlie © 2009, Prentice-Hall, Periodic Table Metalloids border the stair-step line • Intermediate properties © 2009, Prentice-Hall, Periodic Table Nonmetals are on the right side of the periodic table (except H). • Poor conductors • Gain e- to form anion • Brittle © 2009, Prentice-Hall, Alkali Metals Very reactive 1+ Alkaline Earth Metals Reactive 2+ Transition metals Form colored compounds Variable charges Inner Transition Metals Lanthanides or rare earth Actinides are all radioactive Z ≥ 92 are manmade Chalcogens Oxygen family 6 valence electrons Halogens Very reactive Only group containing elements in s, l, g @STP Form acids when bonded with H Noble Gases Low reactivity Full octet Colorless, odorless gas 2.5 Molecules and Molecular Compounds Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Molecular compounds are composed of molecules and almost always contain only nonmetals. © 2009, Prentice-Hall, Allotropes Different forms of the same element. Tin: White (metallic tin), Gray Tin Oxygen: Oxygen gas, Ozone Carbon: Graphite, Diamond, “Bucky Balls” Diatomic Molecules These seven elements occur naturally as molecules containing two atoms. © 2009, Prentice-Hall, Inc. Molecular vs. Empirical formula • Empirical formulas give the lowest wholenumber ratio of atoms of each element in a compound. NO2 • Molecular formulas give the exact number of atoms of each element in a compound. N2O4 © 2009, Prentice-Hall, Inc. Picturing molecules © 2009, Prentice-Hall, 2.6 Ions and Ionic Compounds • Ions – charged particle monoatomic or polyatomic • Cations – positively charged particle • Anions – negatively charged particle Background: Predicting Charges Common Cation Charges Common Anions © 2009, Prentice-Hall, Inc. Ionic Compounds Ionic “bond” is the attraction between oppositely charged ions (generally metals and nonmetals) Called salts Examples: Table salt (NaCl), baking soda (NaHCO3), Epsom salts (MgSO4) © 2009, Prentice-Hall, Writing Formulas Because compounds are electrically neutral, one can determine the formula of a compound this way: – The charge on the cation becomes the subscript on the anion. – The charge on the anion becomes the subscript on the cation. – IONIC ONLY: If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor. © 2009, Prentice-Hall, Inc. 2.7 Naming Inorganic Compounds IONIC Naming Cations • Monoatomic cations have same name as metal Na+ Al3+ • If metals can form > 1 charge use roman numerals Fe2+ Fe3+ • Nonmetal polyatomic cations –ium NH4 + H3O + • Naming Anions • Monoatomic anions drop ending add -ide HO2• Oxyanions –ate or -ite NO3- NO2- • Adding H+ to oxyanions prefix = hydrogen or dihydrogen CO3HCO3H2CO3- Ionic Compounds & Formulas Name the compound Write the formula KCl Mg3N2 Na2SO4 (NH4)2CO3 CuO Cu2O FePO4 Sodium fluoride Calcium nitrate Barium nitrite Lead (II) hydroxide Managanese (IV) sulfide Tin (II) sulfite Sodium nitride Names and Formulas of Acids Name these acids HCN HNO3 HNO2 HClO4 HClO3 H2SO3 HCl HBr HI Writing formulas for acids • If it has hydro- in the name it has no oxygen (Anion ends in –ide) • No hydro in name, anion ends in -ate or -ite • Write anion and add enough H to balance the charges. hydrofluoric acid carbonic acid phosphorous acid dichromic acid hydrophosphoric acid perchloric acid Names and Formulas of Binary Molecular Compounds (covalent) • 2 non metals • Cation followed by anion (-ide) • No roman numerals • Prefixes (no mono on cation) mono, di, tri, tetra, penta, hexa, septa, octa, nona, deca Sulfur dioxide diflourine monoxide nitrogen trichloride diphosphorus pentoxide CO2 CO CCl4 N2O4 XeF6 P2O10 N4O4 Hydrates • Hydrates are compounds that contain discrete water molecules as part of the crystal lattice structure. • CuSO4•5H2O is called copper(II)sulfate pentahydrate. • In the formula you put a dot and then write the number of molecules of water, using the Greek prefixes. • Calcium chloride dihydrate • Cr(NO3)3 6H2O