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Chapter 2 Atoms, Molecules, and Ions
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Section 2.1 The Early History of Chemistry Copyright © Cengage Learning. All rights reserved 2 Greeks were the first to attempt to explain why chemical changes occur. Democritus- atom- “indivisible” Alchemy dominated for 2000 years. Several elements discovered. Mineral acids prepared. Early History of Chemistry
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Section 2.1 The Early History of Chemistry Robert Boyle was the first “chemist”. Performed quantitative experiments. Developed first experimental definition of an element.
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Section 2.2 Fundamental Chemical Laws Three Important Laws Law of conservation of mass (Lavoisier): Mass is neither created nor destroyed in a chemical reaction. Copyright © Cengage Learning. All rights reserved4
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Section 2.2 Fundamental Chemical Laws
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Section 2.2 Fundamental Chemical Laws Law of definite proportions (Proust): A given compound always contains exactly the same proportion of elements by mass.
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Section 2.2 Fundamental Chemical Laws Three Important Laws (continued) Law of multiple proportions (Dalton): When two elements combine with each other to form two or more compounds, the ratios of the masses of one element that combines with the fixed mass of the other are simple whole numbers. Copyright © Cengage Learning. All rights reserved7
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Section 2.2 Fundamental Chemical Laws
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Section 2.3 Dalton’s Atomic Theory Dalton’s Atomic Theory (1808) Each element is made up of tiny particles called atoms. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms. Copyright © Cengage Learning. All rights reserved 9
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Section 2.3 Dalton’s Atomic Theory Dalton’s Atomic Theory (continued) Chemical reactions involve reorganization of the atoms—changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. His model was called the Billiard Ball model Solid sphere Copyright © Cengage Learning. All rights reserved 10
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Section 2.3 Dalton’s Atomic Theory Gay-Lussac and Avogadro (1809—1811) Joseph Louis Gay-Lussac Measured (under same conditions of T and P) the volumes of gases that reacted with each other. The law of combining volumes states that, when gases react together to form other gases, all volumes are measured at the same temperature and pressure. “The ratio between the volumes of the reactant gases and the products can be expressed in simple whole numbers.” 11
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Section 2.3 Dalton’s Atomic Theory Representing Gay—Lussac’s Results Copyright © Cengage Learning. All rights reserved 12
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Section 2.3 Dalton’s Atomic Theory Representing Gay—Lussac’s Results Copyright © Cengage Learning. All rights reserved 13
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Section 2.3 Dalton’s Atomic Theory Avogadro’s Hypothesis At the same T and P, equal volumes of different gases contain the same number of particles. Volume of a gas is determined by the number, not the size, of molecules. Avogadro’s Number 6.022 x 10 23
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Section 2.4 Early Experiments to Characterize the Atom J. J. Thomson (1898—1903) Sealed glass tube with metal disks at each end, air removed One end charged +, the other – Beam appears between disks, could be moved with charged metal plates (repelled by -, attracted by +) Conclusion- particles must be negative because they were attracted to positive plate, came from inside atom, and had much smaller mass 15
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Section 2.4 Early Experiments to Characterize the Atom ATOMS ARE MADE OF SMALLER PARTICLES!!! Thomson’s atomic model- Plum Pudding Model Positively charged blob with negative particles embedded
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Section 2.4 Early Experiments to Characterize the Atom Cathode-Ray Tube Copyright © Cengage Learning. All rights reserved 17
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Section 2.4 Early Experiments to Characterize the Atom Thomson’s Experiment
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Section 2.4 Early Experiments to Characterize the Atom Robert Millikan (1909) Performed experiments involving charged oil drops. Determined the magnitude of the charge on a single electron. Calculated the mass of the electron (9.11 × 10 -31 kg). Copyright © Cengage Learning. All rights reserved 19
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Section 2.4 Early Experiments to Characterize the Atom Copyright © Cengage Learning. All rights reserved 20
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Section 2.4 Early Experiments to Characterize the Atom Henri Becquerel (1896) Discovered radioactivity by observing the spontaneous emission of radiation by uranium. Three types of radioactive emission exist: Gamma rays (ϒ) – high energy light Beta particles (β) – a high speed electron Alpha particles (α) – a particle with a 2+ charge Copyright © Cengage Learning. All rights reserved 21
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Section 2.4 Early Experiments to Characterize the Atom Becquerel’s Experiment
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Section 2.4 Early Experiments to Characterize the Atom Ernest Rutherford (1911) Uranium emits positive particles (alpha-α) Predicted that atoms of gold would not change the path of an alpha (α) particle Gold Foil Experiment Aimed narrow beam of alpha particles at a piece of gold foil surrounded by a screen Expected alpha particles to go straight through, but they were deflected Must be central positive area in atom NUCLEUS Copyright © Cengage Learning. All rights reserved 23
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Section 2.4 Early Experiments to Characterize the Atom Rutherford model- nucleus contains all positive charge, most of atom empty space
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Section 2.4 Early Experiments to Characterize the Atom Rutherford’s Experiment
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Section 2.4 Early Experiments to Characterize the Atom Niels Bohr Refined Rutherford’s model Looked at electrons and their arrangement Electrons exist in specific energy levels Like rungs of a ladder May only reside in each level- ground state Electrons can move from one energy level to another Gain energy as they move up a level- excited state As they return to regular level, they give off energy Light- fireworks
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Section 2.4 Early Experiments to Characterize the Atom Bohr Model
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Section 2.4 Early Experiments to Characterize the Atom Electron Cloud Model Refined Bohr’s model Electron Cloud- visual model of most likely locations of an electron Electron cloud model explains location of electrons Example- fan blades- cannot see individual blades when in motion Makes atom look like sphere under scanning tunneling microscope
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Section 2.4 Early Experiments to Characterize the Atom Electron Cloud Model
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Section 2.4 Early Experiments to Characterize the Atom History of Atomic Chemistry
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Section 2.5 The Modern View of Atomic Structure: An Introduction The atom contains: Electrons – found outside the nucleus; negatively charged. Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge. Neutrons – found in the nucleus; no charge; virtually same mass as a proton. Copyright © Cengage Learning. All rights reserved 31
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Section 2.5 The Modern View of Atomic Structure: An Introduction Two regions Nucleus contains protons and neutrons all positive charges almost all of the mass very small part of the atom Electron cloud contains electrons all negative charge most of the size of the atom involved in bonding Copyright © Cengage Learning. All rights reserved 32
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Section 2.5 The Modern View of Atomic Structure: An Introduction Nuclear Atom Viewed in Cross Section Copyright © Cengage Learning. All rights reserved
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Section 2.5 The Modern View of Atomic Structure: An Introduction Atomic Number Number of protons- identifies element In a neutral atom protons = electrons Atomic Mass (Mass Number) Number of protons + number of neutrons Atomic Symbol Notation Indicates particles in nucleus
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Section 2.5 The Modern View of Atomic Structure: An Introduction
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Section 2.5 The Modern View of Atomic Structure: An Introduction Isotopes Atoms of the same element that have different mass numbers (numbers of neutrons) but the same number of protons Show almost identical chemical properties; chemistry of atom is due to its electrons. In nature most elements contain mixtures of isotopes. This is why atomic mass on table is a decimal- weighted average of all isotopes Copyright © Cengage Learning. All rights reserved 36
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Section 2.5 The Modern View of Atomic Structure: An Introduction
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Section 2.5 The Modern View of Atomic Structure: An Introduction Two Isotopes of Sodium Copyright © Cengage Learning. All rights reserved
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Section 2.5 The Modern View of Atomic Structure: An Introduction Isotopes are identified by: Atomic Number (Z) – number of protons Mass Number (A) – number of protons plus number of neutrons Copyright © Cengage Learning. All rights reserved 39
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Section 2.5 The Modern View of Atomic Structure: An Introduction A certain isotope X contains 23 protons and 28 neutrons. What is the mass number of this isotope? Identify the element. Mass Number = 51 Vanadium Copyright © Cengage Learning. All rights reserved 40 EXERCISE!
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Section 2.6 Molecules and Ions A certain isotope X + contains 54 electrons and 78 neutrons. What is the mass number of this isotope? 133 Copyright © Cengage Learning. All rights reserved 41 EXERCISE!
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Section 2.5 The Modern View of Atomic Structure: An Introduction Calculating Average Atomic Mass Atomic Mass = (% isotope 1 (converted to decimal) x mass isotope 1) + (% isotope 2 (converted to decimal) x mass isotope 2) + …… Carbon-12 makes up 98.93% of all the carbon atoms, while carbon-13 is about 1.07% abundant. Calculate the average atomic mass of carbon. Atomic mass = (0.9893 x 12.00) + (0.0107 x 13.00) = 12.01 amu
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Section 2.7 An Introduction to the Periodic Table The Periodic Table
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Section 2.7 An Introduction to the Periodic Table The Periodic Table Metals vs. Nonmetals Metals Majority of elements Good conductors of heat/electricity Solids at room temperature (except Mercury) Ductile, malleable Vary in reactivity Transition Metals (Gr 3-12) Bridge between sides of table Copyright © Cengage Learning. All rights reserved 44
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Section 2.7 An Introduction to the Periodic Table Nonmetals Opposite to metals Poor conductors, brittle Gases at room temp (except bromine, iodine, sulfur, selenium, phosphorus, and carbon) Vary in reactivity Metalloids Properties of metals and nonmetals Variable conductivity B, Si, Ge, As, Sb, Te, At
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Section 2.7 An Introduction to the Periodic Table Periods – horizontal rows of elements 7 periods on table Increasing atomic number Groups or Families elements in the same vertical columns have similar chemical properties arranged by the number of electrons in the outermost energy level (valence electrons)
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Section 2.7 An Introduction to the Periodic Table Alkali Metals Group 1A Metals with a single valence electron Very reactive- only found in nature as compounds Extremely reactive with water Reactivity increases from top to bottom Francium is most reactive metal Most are stored in oil or in argon gas to protect them from violent reactions
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Section 2.7 An Introduction to the Periodic Table Alkaline Earth Metals Group 2A 2 valence electrons Harder than metals in Group 1 Variable reactivity with water Magnesium- assists in photosynthesis, lightweight metal mixtures Calcium- bones, teeth, chalk, limestone, coral, pearls
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Section 2.7 An Introduction to the Periodic Table Transition Metals Groups 1B-10B (or 3-12) All metals Bridge between 2A and 3A (or 2 and 13) Varying number of valence electrons Varying reactivity Gold, Silver, Copper, and other valuable metals here
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Section 2.7 An Introduction to the Periodic Table The Halogens Group 7A 7 valence electrons Fluorine, Chlorine, Bromine, Iodine- nonmetals Fluorine is most active nonmetal toothpaste, Teflon Chlorine is used in bleach and to kill bacteria in pools and drinking water Iodine keeps the thyroid gland functioning (metabolism) Astatine- metalloid
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Section 2.7 An Introduction to the Periodic Table The Noble Gases Group 8A 8 valence electrons (except He, which has 2) Outermost energy level considered “full”, so they do not react Used to help other things remain unreactive, such as silicon in computer chips and alkali metals Used in light bulbs to help filaments last longer, and in “neon” lights Helium, Neon, Argon, Krypton, Xenon, Radon- all nonmetals
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Section 2.7 An Introduction to the Periodic Table The Periodic Table
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