Chapter 2. Atoms, Molecules, and Ions
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 2.4 Cathode-ray Tube with Perpendicular 2.2 The Discovery of Atomic Structure
Magnetic and Electric Fields
Figure 2.5 Millikan’s Oil-drop Experiment 2.2 The Discovery of Atomic Structure
Figure 2.8 The Behavior of Alpha (), Beta (), and 2.2 The Discovery of Atomic Structure
Animations: Section:
Multiple Proportions 2.1 The Atomic Theory of Matter
Millikan Oil Drop Experiment 2.2 The Discovery of Atomic Structure
Separation of Alpha, Beta, and Gamma Rays 2.2 The Discovery of Atomic Structure
Rutherford Experiment: Nuclear Atom 2.2 The Discovery of Atomic Structure
Activities: Section:
3-D Models: Section:
Hydrogen 2.6 Molecules and Molecular Compounds
Oxygen 2.6 Molecules and Molecular Compounds
Chlorine 2.6 Molecules and Molecular Compounds
Atoms, Molecules, and Ions
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Nitrogen Dioxide 2.6 Molecules and Molecular Compounds
Iodine Pentafluoride 2.6 Molecules and Molecular Compounds
Sodium Chloride (1 1 Unit Cell) 2.7 Ions and Ionic Compounds
Nitrite Ion 2.8 Naming Inorganic Compounds
Acetone 2.9 Some Simple Organic Compounds
Hydroxylamine 2.9 Some Simple Organic Compounds
Other Resources
Further Readings: Section:
Analogical Demonstration 2.1 The Atomic Theory of Matter
A Millikan Oil Drop Analogy 2.2 The Discovery of Atomic Structure
Marie Curie’s Doctoral Thesis: Prelude to a 2.2 The Discovery of Atomic Structure
Nobel Prize
Bowling Balls and Beads: A Concrete Analogy 2.2 The Discovery of Atomic Structure
to the Rutherford Experiment
The Curie-Becquerel Story 2.2 The Discovery of Atomic Structure
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Nomenclature Made Practical: Student Discovery 2.8 Naming Inorganic Compounds
of the Nomenclature
Live Demonstrations: Section:
Turning Plastic into Gold: An Analogy to 2.2 The Discovery of Atomic Structure
Demonstrate Rutherford Gold Foil Experiment
Dramatizing Isotopes: Deuterated Ice Cubes Sink 2.3 The Modern View of Atomic Structure
Atoms, Molecules, and Ions
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Chapter 2. Atoms, Molecules, and Ions
Common Student Misconceptions
• Students have problems with the concept of amu.
• Beginning students often do not see the difference between empirical and molecular formulas.
Teaching Tips
• It is critical that students learn the names and formulas of common and polyatomic ions as soon as
Lecture Outline
2.1 The Atomic Theory of Matter
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• Greek Philosophers: Can matter be subdivided into fundamental particles?
• Democritus (460–370 BC): All matter can be divided into indivisible atomos.
• Dalton: proposed atomic theory with the following postulates:
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FUTURE REFERENCES
• The law of conservation of mass (matter) falls under the First Law of Thermodynamics discussed
in Chapter 5.
2.2 The Discovery of Atomic Structure
• By 1850 scientists knew that atoms consisted of charged particles.
Cathode Rays and Electrons
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• Cathode rays were first discovered in the mid-1800s from studies of electrical discharge through
partially evacuated tubes (cathode-ray tubes, or CRTs).
• Computer terminals were once popularly referred to as CRTs (cathode-ray tubes).
• Cathode rays = radiation produced when high voltage is applied across the tube.
• The voltage causes negative particles to move from the negative electrode (cathode) to the positive
electrode (anode).
Radioactivity
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• Radioactivity is the spontaneous emission of radiation.
• Consider the following experiment:
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• A radioactive substance is placed in a lead shield containing a small hole so that a beam of
radiation is emitted from the shield.
• The radiation is passed between two electrically charged plates and detected.
• Three spots are observed on the detector:
1. a spot deflected in the direction of the positive plate,
2. a spot that is not affected by the electric field, and
The Nuclear Atom
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• The plum pudding model is an early picture of the atom.
• The Thomson model pictures the atom as a sphere with small electrons embedded in a positively
charged mass.
• Rutherford carried out the following “gold foil” experiment:
• A source of -particles was placed at the mouth of a circular detector.
• The -particles were shot through a piece of gold foil.
FUTURE REFERENCES
• Radioactivity will be further discussed in Chapter 21.
2.3 The Modern View of Atomic Structure
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• The atom consists of positive, negative and neutral entities (protons, electrons and neutrons).
• Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the
atom is due to the nucleus.
• Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.
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Z
A
X
• The quantity 1.602 10–19 C is called the electronic charge.
• The charge on an electron is –1.602 10–19 C; the charge on a proton is +1.602 10–19 C;
neutrons are uncharged.
Atomic Numbers, Mass Numbers, and Isotopes
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• Atomic number (Z) = number of protons in the nucleus.
• Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons).
• By convention, for element X, we write .
• Thus, isotopes have the same Z but different A.
• Isotopes of a specific element differ in the number of neutrons.
FUTURE REFERENCES
• The concept of an isotope (specifically 12C) will be useful when defining the mole in Chapter 3.
• Because the atomic number signifies the number of electrons in an atom, it will be commonly
used to write electron configurations of atoms (Chapter 6), draw Lewis structures (Chapter 8),
and understand molecular orbitals (Chapter 9).
• Radioactive decay will be further discussed in Chapter 14 as an example of first order kinetics.
2.4 Atomic Weights
The Atomic Mass Scale
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• Consider 100 g of water:
• Upon decomposition 11.1 g of hydrogen and 88.9 g of oxygen are produced.
• The mass ratio of O to H in water is 88.9/11.1 = 8.
• Therefore, the mass of O is 2 8 = 16 times the mass of H.
• If H has a mass of 1, then O has a relative mass of 16.
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• Atomic mass units (amu) are convenient units to use when dealing with extremely small masses
of individual atoms.
• 1 amu = 1.66054 10–24 g and 1 g = 6.02214 1023 amu
• By definition, the mass of 12C is exactly 12 amu.
Average Atomic Masses
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• We average the masses of isotopes to give average atomic masses.
The Mass Spectrometer
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• A mass spectrometer is an instrument that allows for direct and accurate determination of atomic
(and molecular) weights.
• The sample is charged as soon as it enters the spectrometer.
FUTURE REFERENCES
• Being able to locate atomic weights on the periodic table will be crucial in calculating molar
masses in Chapter 3 and beyond.
2.5 The Periodic Table
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• The periodic table is used to organize the elements in a meaningful way.
• As a consequence of this organization, there are periodic properties associated with the periodic table.
• Rows in the periodic table are called periods.
• Columns in the periodic table are called groups.
• Several numbering conventions are used (i.e., groups may be numbered from 1 to 18, or from 1A
to 8A and 1B to 8B).
• Some of the groups in the periodic table are given special names.
• These names indicate the similarities between group members.
• Examples:
• Group 1A: alkali metals
• Group 2A: alkaline earth metals
• Group 7A: halogens
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FORWARD REFERENCES
• Additional information that can be associated with the unique location of an element in the
periodic table will be covered in Chapter 6 (electron configurations), Chapter 7 (periodic
properties), Chapter 8 (tendency to form ionic or covalent bonds), and Chapter 16 (relative acid
strength).
2.6 Molecules and Molecular Compounds
• A molecule consists of two or more atoms bound tightly together.
Molecules and Chemical Formulas
• Each molecule has a chemical formula.
• The chemical formula indicates
1. which atoms are found in the molecule, and
Molecular and Empirical Formulas
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• Molecular formulas
• These formulas give the actual numbers and types of atoms in a molecule.
• Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4.
• Empirical formulas
• These formulas give the relative numbers and types of atoms in a molecule (they give the lowest
whole-number ratio of atoms in a molecule).
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• Examples: H2O, CO2, CO, CH4, HO, CH2.
Picturing Molecules
• Molecules occupy three-dimensional space.
• However, we often represent them in two dimensions.
FORWARD REFERENCES
• More detailed discussion of bonding in molecules and molecular shapes will take place in
Chapters 8 and 9, respectively.
2.7 Ions and Ionic Compounds
• If electrons are added to or removed from a neutral atom, an ion is formed.
• When an atom or molecule loses electrons it becomes positively charged.
Predicting Ionic Charges
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• An atom or molecule can lose more than one electron.
• Many atoms gain or lose enough electrons to have the same number of electrons as the nearest noble
gas (group 8A).
Ionic Compounds
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• A great deal of chemistry involves the transfer of electrons between species.
• Example:
• To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na+.
• The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then
becomes an anion: Cl–.
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• In general, ionic compounds are combinations of metals and nonmetals, whereas molecular
compounds are composed of nonmetals only.
• There are exceptions; notably (NH4)2SO4 and other ammonium salts are ionic.
• Writing empirical formulas for ionic compounds:
• You need to know the ions of which it is composed.
Chemistry and Life: Elements Required by Living Organisms
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• Of the known elements, only about 29 are required for life.
• Water accounts for at least 70% of the mass of most cells.
• More than 97% of the mass of most organisms comprises just six elements (O, C, H, N, P and S).
FORWARD REFERENCES
• Formulas (including correct charges) of ions will be important in writing metathesis and net
ionic equations in Chapter 4 (sections 4.2-4.3).
• Periodic trends in ionization energy (in gas phase) as well as ionic radii (in crystals) will be
covered in Chapter 7.
• The nature of bonding between ions and charges of most monoatomic ions will be rationalized in
terms of electron configurations in Chapter 8 (section 8.2).
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2.8 Naming Inorganic Compounds
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• Chemical nomenclature is the naming of substances.
• Common names are traditional names for substances (e.g., water, ammonia).
Names and Formulas of Ionic Compounds
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1. Positive Ions (Cations)
• Cations formed from a metal have the same name as the metal.
• Example: Na+ = sodium ion.
• Ions formed from a single atom are called monoatomic ions.
2. Negative Ions (Anions)
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• Monatomic anions (with only one atom) use the ending -ide.
• Example: Cl– is the chloride ion.
• Some polyatomic anions also use the -ide ending:
• Examples: hydroxide, cyanide, and peroxide ions.
• Polyatomic anions containing oxygen with more than two members in the series are named as follows
(in order of decreasing oxygen):
• per-….-ate example: ClO4– perchlorate
• -ate ClO3– chlorate
• -ite ClO2– chlorite
• hypo-….-ite ClO– hypochlorite
• Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi-
(one H), dihydrogen (two H) etc., to the name as follows:
• CO32– is the carbonate anion.
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3. Ionic Compounds
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• These are named by the cation then the anion.
• Examples:
• CaCl2 = calcium chloride
• (NH4)3PO4 = ammonium phosphate
• KClO4 = potassium perchlorate
Names and Formulas of Acids
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• Acids are substances that yield hydrogen ions when dissolved in water (Arrhenius definition).
Names and Formulas of Binary Molecular Compounds
• Binary molecular compounds have two elements.
• The most metallic element (i.e., the one to the farthest left on the periodic table) is usually written
first. The exception is NH3.
FORWARD REFERENCES
• Nomenclature will be required throughout the textbook.
• Acids will be mentioned again in Chapter 4 and further discussed in Chapters 16 and 17.
2.9 Some Simple Organic Compounds
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• Organic chemistry is the study of carbon-containing compounds.
• Organic compounds are those that contain carbon and hydrogen, often in combination with other
elements.
Alkanes
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Atoms, Molecules, and Ions
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• The names of alkanes end in –ane.
• Examples: methane, ethane, propane, butane.
Some Derivatives of Alkanes
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• When functional groups, specific groups of atoms, are used to replace hydrogen atoms on alkanes,
new classes of organic compounds are obtained.
• Alcohols are obtained by replacing a hydrogen atom of an alkane with an –OH group.
FORWARD REFERENCES
• Simple organic compounds will be used throughout the textbook to illustrate: weak acid behavior
(e.g., acetic acid in Chapters 16 and 17), weak base behavior (e.g., amines in Chapters 16 and 17),
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Further Readings:
1. John J. Fortman, “Analogical Demonstration,” J. Chem. Educ., Vol. 69, 1992, 323–324. This reference
includes demonstrations of the concepts of the conservation of mass in chemical reactions, the Law of
Multiple Proportions, etc.
6. Harold F. Walton, “The Curie–Becquerel Story,” J. Chem. Educ., Vol. 69, 1992, 10–15.
7. William Spindel and Takanobu Ishida, “Isotope Separation,” J. Chem. Educ., Vol. 68, 1991, 312–318.
An article describing methods used to isolate important isotopes.
8. Stephen DeMeo, “Revisiting Molar Mass, Atomic Mass, and Mass Number: Organizing, Integrating,
and Sequencing Fundamental Chemical Concepts,” J. Chem. Educ., Vol. 83, 2006, 617–620.
9. Josefina Arce de Sanabia, “Relative Atomic Mass and the Mole: A Concrete Analogy to Help Students
Understand These Abstract Concepts,” J. Chem. Educ, Vol. 70, 1993, 233–234.
13. Steven I. Dutch, “Periodic Tables of Elemental Abundance,” J. Chem. Educ., Vol. 76, 1999, 356–358.
14. Werner Fischer, “A Second Note on the Term ‘Chalcogen’,” J. Chem. Educ., Vol. 78, 2001, 1333.
15. Marshall W. Cronyn, “The Proper Place for Hydrogen in the Periodic Table,” J. Chem. Educ., Vol.
80, 2003, 947–950.
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18. Michael C. Wirtz, Joan Kaufmann, and Gary Hawley, “Nomenclature Made Practical: Student
Discovery of the Nomenclature Rules,” J. Chem. Educ., Vol. 83, 2006, 595–598.
19. Steven J. Hawkes, “A Mnemonic for Oxy-Anions,” J. Chem. Educ., Vol. 67, 1990, 149.
Live Demonstrations:
1. Arthur B. Ellis, Edward A Adler, and Frederick H. Juergens, “Dramatizing Isotopes: Deuterated Ice
Cubes Sink,” J. Chem. Educ., Vol. 67, 1990, 159–160. Differences in density of H2O(l) and D2O(s) are
used to demonstrate the effects of isotopic substitution.