Chapter 2 Atoms, Molecules, and Ions

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

Section 2.1 The Early History of Chemistry  Robert Boyle was the first “chemist”.  Performed quantitative experiments.  Developed first experimental definition of an element.

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

Section 2.2 Fundamental Chemical Laws

Section 2.2 Fundamental Chemical Laws  Law of definite proportions (Proust):  A given compound always contains exactly the same proportion of elements by mass.

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

Section 2.2 Fundamental Chemical Laws

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

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

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

Section 2.3 Dalton’s Atomic Theory Representing Gay—Lussac’s Results Copyright © Cengage Learning. All rights reserved 12

Section 2.3 Dalton’s Atomic Theory Representing Gay—Lussac’s Results Copyright © Cengage Learning. All rights reserved 13

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  x 10 23

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

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

Section 2.4 Early Experiments to Characterize the Atom Cathode-Ray Tube Copyright © Cengage Learning. All rights reserved 17

Section 2.4 Early Experiments to Characterize the Atom Thomson’s Experiment

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 × kg). Copyright © Cengage Learning. All rights reserved 19

Section 2.4 Early Experiments to Characterize the Atom Copyright © Cengage Learning. All rights reserved 20

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

Section 2.4 Early Experiments to Characterize the Atom Becquerel’s Experiment

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

Section 2.4 Early Experiments to Characterize the Atom  Rutherford model- nucleus contains all positive charge, most of atom empty space

Section 2.4 Early Experiments to Characterize the Atom Rutherford’s Experiment

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

Section 2.4 Early Experiments to Characterize the Atom Bohr Model

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

Section 2.4 Early Experiments to Characterize the Atom Electron Cloud Model

Section 2.4 Early Experiments to Characterize the Atom History of Atomic Chemistry

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

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

Section 2.5 The Modern View of Atomic Structure: An Introduction Nuclear Atom Viewed in Cross Section Copyright © Cengage Learning. All rights reserved

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

Section 2.5 The Modern View of Atomic Structure: An Introduction

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

Section 2.5 The Modern View of Atomic Structure: An Introduction

Section 2.5 The Modern View of Atomic Structure: An Introduction Two Isotopes of Sodium Copyright © Cengage Learning. All rights reserved

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

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!

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!

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 = ( x 12.00) + ( x 13.00) = amu

Section 2.7 An Introduction to the Periodic Table The Periodic Table

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

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

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)

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

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

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

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

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

Section 2.7 An Introduction to the Periodic Table The Periodic Table