Chapter 2. Atoms, Molecules, and Ions
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1 Chapter 2 Atoms, Molecules, and Ions
2 Section 2.1 The Early History of Chemistry Early History of Chemistry Greeks were the first to attempt to explain why chemical changes occur. Alchemy dominated for 2000 years. Several elements discovered. Mineral acids prepared. Robert Boyle was the first chemist. Performed quantitative experiments. Developed first experimental definition of an element. Copyright Cengage Learning. All rights reserved 2
3 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. Law of definite proportion (Proust): A given compound always contains exactly the same proportion of elements by mass. Copyright Cengage Learning. All rights reserved 3
4 Section 2.2 Fundamental Chemical Laws Three Important Laws (continued) Law of multiple proportions (Dalton): When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. Copyright Cengage Learning. All rights reserved 4
5 Section 2.3 Dalton s Atomic Theory Dalton s Atomic Theory (1808) Each element is made up of tiny particles called atoms. Copyright Cengage Learning. All rights reserved 5
6 Section 2.3 Dalton s Atomic Theory Dalton s Atomic Theory (continued) The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. Copyright Cengage Learning. All rights reserved 6
7 Section 2.3 Dalton s Atomic Theory Dalton s Atomic Theory (continued) 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 7
8 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. Copyright Cengage Learning. All rights reserved 8
9 Section 2.3 Dalton s Atomic Theory Gay-Lussac and Avogadro ( ) Gay Lussac Measured (under same conditions of T and P) the volumes of gases that reacted with each other. 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. Copyright Cengage Learning. All rights reserved 9
10 Section 2.3 Dalton s Atomic Theory Representing Gay Lussac s Results Copyright Cengage Learning. All rights reserved 10
11 Section 2.3 Dalton s Atomic Theory Representing Gay Lussac s Results Copyright Cengage Learning. All rights reserved 11
12 Section 2.3 Dalton s Atomic Theory Representing Gay Lussac s Results Copyright Cengage Learning. All rights reserved 12
13 Section 2.3 Dalton s Atomic Theory Representing Gay Lussac s Results Copyright Cengage Learning. All rights reserved 13
14 Section 2.4 Early Experiments to Characterize the Atom J. J. Thomson ( ) Postulated the existence of negatively charged particles, that we now call electrons, using cathode-ray tubes. Determined the charge-to-mass ratio of an electron. The atom must also contain positive particles that balance exactly the negative charge carried by electrons. Copyright Cengage Learning. All rights reserved 14
15 Section 2.4 Early Experiments to Characterize the Atom Cathode-Ray Tube Copyright Cengage Learning. All rights reserved 15
16 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 ( kg). Copyright Cengage Learning. All rights reserved 16
17 Section 2.4 Early Experiments to Characterize the Atom Millikan Oil Drop Experiment To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright Cengage Learning. All rights reserved 17
18 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 18
19 Section 2.4 Early Experiments to Characterize the Atom Ernest Rutherford (1911) Explained the nuclear atom. The atom has a dense center of positive charge called the nucleus. Electrons travel around the nucleus at a large distance relative to the nucleus. Copyright Cengage Learning. All rights reserved 19
20 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 20
21 Section 2.5 The Modern View of Atomic Structure: An Introduction The nucleus is: Small compared with the overall size of the atom. Extremely dense; accounts for almost all of the atom s mass. Copyright Cengage Learning. All rights reserved 21
22 Section 2.5 The Modern View of Atomic Structure: An Introduction Nuclear Atom Viewed in Cross Section Copyright Cengage Learning. All rights reserved
23 Section 2.5 The Modern View of Atomic Structure: An Introduction Isotopes Atoms with the same number of protons but different numbers of neutrons. Show almost identical chemical properties; chemistry of atom is due to its electrons. In nature most elements contain mixtures of isotopes. Copyright Cengage Learning. All rights reserved 23
24 Section 2.5 The Modern View of Atomic Structure: An Introduction Two Isotopes of Sodium Copyright Cengage Learning. All rights reserved
25 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 25
26 Section 2.5 The Modern View of Atomic Structure: An Introduction EXERCISE! 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 26
27 Section 2.6 Molecules and Ions Chemical Bonds Covalent Bonds Bonds form between atoms by sharing electrons. Resulting collection of atoms is called a molecule. Copyright Cengage Learning. All rights reserved 27
28 Section 2.6 Molecules and Ions Covalent Bonding To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright Cengage Learning. All rights reserved 28
29 Section 2.6 Molecules and Ions Chemical Bonds Ionic Bonds Bonds form due to force of attraction between oppositely charged ions. Ion atom or group of atoms that has a net positive or negative charge. Cation positive ion; lost electron(s). Anion negative ion; gained electron(s). Copyright Cengage Learning. All rights reserved 29
30 Section 2.6 Molecules and Ions Molecular vs Ionic Compounds To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright Cengage Learning. All rights reserved 30
31 Section 2.6 Molecules and Ions EXERCISE! 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 31
32 Section 2.6 Molecules and Ions CONCEPT CHECK! Which of the following statements regarding Dalton s atomic theory are still believed to be true? I. Elements are made of tiny particles called atoms. II. All atoms of a given element are identical. III. A given compound always has the same relative numbers and types of atoms. IV. Atoms are indestructible. Copyright Cengage Learning. All rights reserved 32
33 Section 2.7 An Introduction to the Periodic Table The Periodic Table Metals vs. Nonmetals Groups or Families elements in the same vertical columns; have similar chemical properties Periods horizontal rows of elements Copyright Cengage Learning. All rights reserved 33
34 Section 2.7 An Introduction to the Periodic Table The Periodic Table
35 Section 2.7 An Introduction to the Periodic Table Groups or Families Table of common charges formed when creating ionic compounds. Group or Family Charge Alkali Metals (1A) 1+ Alkaline Earth Metals (2A) 2+ Halogens (7A) 1 Noble Gases (8A) 0 Copyright Cengage Learning. All rights reserved 35
36 Section 2.8 Naming Simple Compounds Naming Compounds Binary Compounds Composed of two elements Ionic and covalent compounds included Binary Ionic Compounds Metal nonmetal Binary Covalent Compounds Nonmetal nonmetal Copyright Cengage Learning. All rights reserved 36
37 Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type I) 1. The cation is always named first and the anion second. 2. A monatomic cation takes its name from the name of the parent element. 3. A monatomic anion is named by taking the root of the element name and adding ide. Copyright Cengage Learning. All rights reserved 37
38 Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type I) Examples: KCl Potassium chloride MgBr 2 Magnesium bromide CaO Calcium oxide Copyright Cengage Learning. All rights reserved 38
39 Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type II) Metals in these compounds form more than one type of positive ion. Charge on the metal ion must be specified. Roman numeral indicates the charge of the metal cation. Transition metal cations usually require a Roman numeral. Elements that form only one cation do not need to be identified by a roman numeral. Copyright Cengage Learning. All rights reserved 39
40 Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type II) Examples: CuBr Copper(I) bromide FeS Iron(II) sulfide PbO 2 Lead(IV) oxide Copyright Cengage Learning. All rights reserved 40
41 Section 2.8 Naming Simple Compounds Polyatomic Ions Must be memorized (see Table 2.5 on pg. 65 in text). Examples of compounds containing polyatomic ions: NaOH Sodium hydroxide Mg(NO 3 ) 2 Magnesium nitrate (NH 4 ) 2 SO 4 Ammonium sulfate Copyright Cengage Learning. All rights reserved 41
42 Section 2.8 Naming Simple Compounds Formation of Ionic Compounds To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright Cengage Learning. All rights reserved 42
43 Section 2.8 Naming Simple Compounds Binary Covalent Compounds (Type III) Formed between two nonmetals. 1. The first element in the formula is named first, using the full element name. 2. The second element is named as if it were an anion. 3. Prefixes are used to denote the numbers of atoms present. 4. The prefix mono- is never used for naming the first element. Copyright Cengage Learning. All rights reserved 43
44 Section 2.8 Naming Simple Compounds Prefixes Used to Indicate Number in Chemical Names Copyright Cengage Learning. All rights reserved 44
45 Section 2.8 Naming Simple Compounds Binary Covalent Compounds (Type III) Examples: CO 2 Carbon dioxide SF 6 Sulfur hexafluoride N 2 O 4 Dinitrogen tetroxide Copyright Cengage Learning. All rights reserved 45
46 Section 2.8 Naming Simple Compounds Flowchart for Naming Binary Compounds Copyright Cengage Learning. All rights reserved 46
47 Section 2.8 Naming Simple Compounds Overall Strategy for Naming Chemical Compounds Copyright Cengage Learning. All rights reserved 47
48 Section 2.8 Naming Simple Compounds Acids Acids can be recognized by the hydrogen that appears first in the formula HCl. Molecule with one or more H + ions attached to an anion. Copyright Cengage Learning. All rights reserved 48
49 Section 2.8 Naming Simple Compounds Acids If the anion does not contain oxygen, the acid is named with the prefix hydro and the suffix ic. Examples: HCl HCN H 2 S Hydrochloric acid Hydrocyanic acid Hydrosulfuric acid Copyright Cengage Learning. All rights reserved 49
50 Section 2.8 Naming Simple Compounds Acids If the anion does contain oxygen: The suffix ic is added to the root name if the anion name ends in ate. Examples: HNO 3 H 2 SO 4 HC 2 H 3 O 2 Nitric acid Sulfuric acid Acetic acid Copyright Cengage Learning. All rights reserved 50
51 Section 2.8 Naming Simple Compounds Acids If the anion does contain oxygen: The suffix ous is added to the root name if the anion name ends in ite. Examples: HNO 2 Nitrous acid H 2 SO 3 Sulfurous acid HClO 2 Chlorous acid Copyright Cengage Learning. All rights reserved 51
52 Section 2.8 Naming Simple Compounds Flowchart for Naming Acids Copyright Cengage Learning. All rights reserved
53 Section 2.8 Naming Simple Compounds EXERCISE! Which of the following compounds is named incorrectly? a) KNO 3 potassium nitrate b) TiO 2 titanium(ii) oxide c) Sn(OH) 4 tin(iv) hydroxide d) PBr 5 phosphorus pentabromide e) CaCrO 4 calcium chromate Copyright Cengage Learning. All rights reserved 53
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