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Chapter 2 Atoms, Molecules, and Ions

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1 Chapter 2 Atoms, Molecules, and Ions

2 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

3 Copyright © Cengage Learning. All rights reserved
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

4 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

5 Dalton’s Atomic Theory
Matter consists of tiny particles called atoms. Atoms are indestructible. In chemical reactions, the atoms rearrange but they do not themselves break apart. In any sample of a pure element, all the atoms are identical in mass and other properties. The atoms of different elements differ in mass and other properties. In a given compound the constituents atoms are always present in the same fixed numerical ratio.

6 Gay-Lussac and Avogadro (1809—1811)
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

7 Representing Gay—Lussac’s Results
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8 Representing Gay—Lussac’s Results
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9 Copyright © Cengage Learning. All rights reserved
J. J. Thomson (1898—1903) 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

10 Copyright © Cengage Learning. All rights reserved
Cathode-Ray Tube Copyright © Cengage Learning. All rights reserved

11 Copyright © Cengage Learning. All rights reserved
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

12 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

13 Copyright © Cengage Learning. All rights reserved
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

14 Copyright © Cengage Learning. All rights reserved
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

15 Florescent Screen Lead block Uranium Gold Foil

16 Copyright © Cengage Learning. All rights reserved
The atom contains: Electrons – found outside the nucleus; negatively charged. Mass = x g Symbol 0-1e Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge. Mass = x10-24 g Symbol 11H+1 Neutrons – found in the nucleus; no charge; virtually same mass as a proton. Mass = x10-24 g Symbol 10n Copyright © Cengage Learning. All rights reserved

17

18 ATOMIC THEORY Atomic number = # of protons.
In a neutral atom the number of protons equals the number of electrons. Mass number = number of protons plus the number of neutrons. Mass number minus the atomic number equals the number of neutrons.

19 Isotopes -- atoms having the same number of protons but a different number of neutrons.
Ions -- atoms that have acquired a charge by gaining or losing electrons. Metals usually lose electrons and nonmetals usually gain electrons. Metallic ions normally have a + charge and nonmetallic ions a negative charge. Cation (+ ion) and anion(- ion)

20 Symbols A X Z 23 Na 11

21 Copyright © Cengage Learning. All rights reserved
Two Isotopes of Sodium Copyright © Cengage Learning. All rights reserved

22 Copyright © Cengage Learning. All rights reserved
Chemical Bonds Covalent Bonds Bonds form between atoms by sharing electrons. Resulting collection of atoms is called a molecule. - Two atoms share a pair of electrons. - Properties -- soft, low melting point, poor conductor of electricity Copyright © Cengage Learning. All rights reserved

23 Covalent Bonding To play movie you must be in Slide Show Mode
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24 Copyright © Cengage Learning. All rights reserved
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

25 IONIC BONDING 1. One atom, usually a metal, loses an electron(s) to form a positive ion. 2. A second atom, usually a nonmetal, gains an electron(s) to form a negative ion. 3. positive and negative ions attract Properties -- hard, brittle, high m.p., conducts electricity in solution or as liquid but not as a solid.

26 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

27 Copyright © Cengage Learning. All rights reserved
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

28 TYPES OF ELEMENTS METALS -- shiny, conduct heat and electricity, ductile (can be pulled into wires), malleable (can be rolled into sheets) NONMETALS -- dull, brittle, poor conductors METALLOIDS -- have properties of both metals and nonmetals. Also known as semimetals. POSITION ON PERIODIC CHART

29 The Periodic Table

30 Groups or Families Table of common charges formed when creating ionic compounds. Group or Family Charge Alkali Metals Alkali Metals (1A) 1+ Alkaline Earth Metals (2A) 2+ Halogens (7A) 1– Noble Gases (8A) Copyright © Cengage Learning. All rights reserved

31 Expanded Periodic Table

32 Copyright © Cengage Learning. All rights reserved
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

33 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

34 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

35 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

36 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

37 Prefixes Used to Indicate Number in Chemical Names
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38 Flowchart for Naming Binary Compounds
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39 Overall Strategy for Naming Chemical Compounds
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40 Copyright © Cengage Learning. All rights reserved
Acids If the anion does not contain oxygen, the acid is named with the prefix hydro– and the suffix –ic. Examples: HCl Hydrochloric acid HCN Hydrocyanic acid H2S Hydrosulfuric acid Copyright © Cengage Learning. All rights reserved

41 Copyright © Cengage Learning. All rights reserved
Acids If the anion does contain oxygen: The suffix –ic is added to the root name if the anion name ends in –ate. Examples: HNO3 Nitric acid H2SO4 Sulfuric acid HC2H3O2 Acetic acid Copyright © Cengage Learning. All rights reserved

42 Copyright © Cengage Learning. All rights reserved
Acids If the anion does contain oxygen: The suffix –ous is added to the root name if the anion name ends in –ite. Examples: HNO2 Nitrous acid H2SO3 Sulfurous acid HClO2 Chlorous acid Copyright © Cengage Learning. All rights reserved

43 Flowchart for Naming Acids
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