chapter 2 atoms, molecules, and ions. chapter 2 table of contents return to toc copyright © cengage...
TRANSCRIPT
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Chapter 2
Atoms, Molecules,
and Ions
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Chapter 2
Table of Contents
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2.1 The Early History of Chemistry
2.2 Fundamental Chemical Laws
2.3 Dalton’s Atomic Theory
2.4 Early Experiments to Characterize the Atom
2.5 The Modern View of Atomic Structure: An Introduction
2.6 Molecules and Ions
2.7 An Introduction to the Periodic Table
2.8 Naming Simple Compounds
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Section 2.2
Fundamental Chemical Laws
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• Law of conservation of mass (Lavoisier): Mass is neither created nor destroyed.
• Law of definite proportion (Proust): A given compound always contains exactly the same
proportion of elements by mass.
Three Important Laws
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Section 2.2
Fundamental Chemical Laws
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• 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.
Three Important Laws (continued)
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Section 2.3
Dalton’s Atomic Theory
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• Each element is made up of tiny particles called atoms.
Dalton’s Atomic Theory (1808)
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Section 2.3
Dalton’s Atomic Theory
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• The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.
Dalton’s Atomic Theory (continued)
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Section 2.3
Dalton’s Atomic Theory
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• 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.
Dalton’s Atomic Theory (continued)
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Section 2.3
Dalton’s Atomic Theory
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• 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.
Dalton’s Atomic Theory (continued)
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Section 2.3
Dalton’s Atomic Theory
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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.
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Section 2.4
Early Experiments to Characterize the Atom
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• Postulated the existence of 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 particles that we now call electrons.
J. J. Thomson (1898—1903)
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Section 2.4
Early Experiments to Characterize the Atom
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Cathode-Ray Tube
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Section 2.4
Early Experiments to Characterize the Atom
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(Uranium
compound) 2.2
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Section 2.4
Early Experiments to Characterize the Atom
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• Performed experiments involving charged oil drops.
• Determined the magnitude of the charge on a single electron.
• Calculated the mass of the electron.
Robert Millikan (1909)
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Section 2.4
Early Experiments to Characterize the Atom
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Millikan Oil Drop Experiment
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Section 2.4
Early Experiments to Characterize the Atom
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• Explained the nuclear atom.
• Atom has a dense center of positive charge called the nucleus.
• Electrons travel around the nucleus at a relatively large distance.
Ernest Rutherford (1911)
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Section 2.4
Early Experiments to Characterize the Atom
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Rutherford’s Gold Foil Experiment
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Section 2.5
The Modern View of Atomic Structure: An Introduction
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• 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.
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Section 2.5
The Modern View of Atomic Structure: An Introduction
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• The nucleus is: Small compared with the overall size of the
atom. Extremely dense; accounts for almost all of
the atom’s mass.
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Section 2.5
The Modern View of Atomic Structure: An Introduction
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Nuclear Atom Viewed in Cross Section
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Section 2.5
The Modern View of Atomic Structure: An Introduction
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• 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.
Isotopes
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Section 2.5
The Modern View of Atomic Structure: An Introduction
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Two Isotopes of Sodium
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Section 2.5
The Modern View of Atomic Structure: An Introduction
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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
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Section 2.6
Molecules and Ions
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• Covalent Bonds Bonds form between atoms by sharing
electrons. Resulting collection of atoms is called a
molecule.
Chemical Bonds
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Section 2.6
Molecules and Ions
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Covalent Bonding
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Section 2.6
Molecules and Ions
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• 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).
Chemical Bonds
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Section 2.6
Molecules and Ions
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Molecular vs. Ionic Compounds
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Section 2.6
Molecules and Ions
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Exercise
A certain isotope X+ contains 54 electrons and 78 neutrons.
• What is the mass number of this isotope?
133
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Section 2.7
An Introduction to the Periodic Table
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• Metals vs. Nonmetals• Groups or Families – elements in the same
vertical columns; have similar chemical properties
• Periods – horizontal rows of elements
The Periodic Table
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Section 2.7
An Introduction to the Periodic Table
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The Periodic Table
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Section 2.7
An Introduction to the Periodic Table
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• Table of common charges formed when creating ionic compounds.
Groups or Families
Group or Family Charge
Alkali Metals (1A) 1+
Alkaline Earth Metals (2A) 2+
Halogens (7A) 1–
Noble Gases (8A) 0
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Section 2.8
Naming Simple Compounds
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• Binary Compounds Composed of two elements Ionic and covalent compounds included
• Binary Ionic Compounds Metal—nonmetal
• Binary Covalent Compounds Nonmetal—nonmetal
Common Names only: H2O & NH3
Naming Compounds
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Section 2.8
Naming Simple Compounds
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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.
Binary Ionic Compounds (Type I)
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Section 2.8
Naming Simple Compounds
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• Examples:
KCl Potassium chloride
MgBr2 Magnesium bromide
CaO Calcium oxide
Binary Ionic Compounds (Type I)
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Section 2.8
Naming Simple Compounds
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• Metals in these compounds form more than one type of positive charge.
• 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.
Binary Ionic Compounds (Type II)
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Section 2.8
Naming Simple Compounds
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• Examples:
CuBr Copper(I) bromide
FeS Iron(II) sulfide
PbO2 Lead(IV) oxide
Binary Ionic Compounds (Type II)
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Section 2.8
Naming Simple Compounds
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• Must be memorized (see Table 2.5 on pg. 62 in text).
• Examples of compounds containing polyatomic ions:
NaOH Sodium hydroxide
Mg(NO3)2 Magnesium nitrate
(NH4)2SO4 Ammonium sulfate
Polyatomic Ions
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Section 2.8
Naming Simple Compounds
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Formation of Ionic Compounds
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Examples of Older Names of Cations formed from Transition Metals
(memorize these!!)From Zumdahl HD book – NOT in AP book!!!
Use these names, when possible, on the homework worksheet
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Section 2.8
Naming Simple Compounds
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• 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 (-ide).
3. Prefixes are used to denote the numbers of atoms present.
4. The prefix mono- is never used for naming the first element.
Binary Covalent Compounds (Type III)
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Section 2.8
Naming Simple Compounds
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Prefixes Used to Indicate Number in Chemical Names
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Section 2.8
Naming Simple Compounds
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• Examples:
CO2 Carbon dioxide
SF6 Sulfur hexafluoride
N2O4 Dinitrogen tetroxide
Binary Covalent Compounds (Type III)
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Section 2.8
Naming Simple Compounds
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Overall Strategy for Naming Chemical Compounds
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Section 2.8
Naming Simple Compounds
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Flowchart for Naming Binary Compounds
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Section 2.8
Naming Simple Compounds
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• 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.
Acids
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Section 2.8
Naming Simple Compounds
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• If the anion does not contain oxygen, the acid is named with the prefix hydro– and the suffix –ic.
• Must be dissolved in water to be an acid! (aq) or acid required – otherwise it’s a GAS
• Examples:
HCl Hydrochloric acid
HCN Hydrocyanic acid
H2S Hydrosulfuric acid
Acids
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Section 2.8
Naming Simple Compounds
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Copyright © Cengage Learning. All rights reserved 46
• If the anion does contain oxygen: The suffix –ic is added to the root name if
the anion name ends in –ate. Automatically acid – these gases do not
exist• Examples:
HNO3Nitric acid
H2SO4 Sulfuric acid
HC2H3O2 Acetic acid
Acids
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Section 2.8
Naming Simple Compounds
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Copyright © Cengage Learning. All rights reserved 47
• If the anion does contain oxygen: The suffix –ous is added to the root name
if the anion name ends in –ite. Automatically acid – these gases do not
exist• Examples:
HNO2Nitrous acid
H2SO3 Sulfurous acid
HClO2 Chlorous acid
Acids
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Section 2.8
Naming Simple Compounds
Return to TOC2.7
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Section 2.8
Naming Simple Compounds
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Flowchart for Naming Acids
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Section 2.8
Naming Simple Compounds
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Exercise
Which of the following compounds is named incorrectly?
a) KNO3 potassium nitrate
b) TiO2 titanium(II) oxide
c) Sn(OH)4 tin(IV) hydroxide
d) PBr5 phosphorus pentabromide
e) CaCrO4 calcium chromate
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Mixed Practice
1. Dinitrogen monoxide
2. Potassium sulfide
3. Copper (II) nitrate
4. Dichlorine heptoxide
5. Chromium (III) sulfate
6. Ferric sulfite
7. Calcium oxide
8. Barium carbonate
9. Iodine monochloride
1. N2O
2. K2S
3. Cu(NO3)2
4. Cl2O7
5. Cr2(SO4)3
6. Fe2(SO3)3
7. CaO
8. BaCO3
9. ICl
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Mixed Practice
1.1. BaIBaI22
2.2. PP44SS33
3.3. Ca(OH)Ca(OH)22
4.4. FeCOFeCO33
5.5. NaNa22CrCr22OO77
6.6. II22OO55
7.7. Cu(ClOCu(ClO44))22
8.8. CSCS22
9.9. BB22ClCl44
1.1. Barium iodideBarium iodide
2.2. Tetraphosphorus trisulfideTetraphosphorus trisulfide
3.3. Calcium hydroxideCalcium hydroxide
4.4. Iron (II) carbonateIron (II) carbonate
5.5. Sodium dichromateSodium dichromate
6.6. Diiodine pentoxideDiiodine pentoxide
7.7. Cupric perchlorateCupric perchlorate
8.8. Carbon disulfideCarbon disulfide
9.9. Diboron tetrachlorideDiboron tetrachloride