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Chemistry : The Study of Matter and Its Changes

by: Brady, James E; Senese, Fred

Chemistry : The Study of Matter and Its Changes

by: Brady, James E; Senese, Fred

ISBN 13:
9780471215172
ISBN 10:
0471215171
Edition: 4th
Format: Hardcover
Copyright: 02/01/2004
Publisher: Wiley
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Summary

Work more effectively and gauge your progress as you go along! This Study Guide is designed to accompany Brady & Senese' s "Chemistry: Matter & Its Changes, 4th Edition," It contains learning objectives, chapter reviews, examples, and self-test questions. Brady & Senese's Chemistry: Matter & Its Changes, 4th Edition, is a reader-friendly textbook that makes the content accessible without sacrificing either breadth or depth of coverage. The text's informal writing style, emphasis on problem solving, and state-of-the-art media package make this book an ideal fit for readers of various backgrounds and abilities. The 4th edition welcomes new co-author Fred Senese, the architect of the most visited general chemistry website. Together Jim Brady and Fred Senese offer accurate, lucid, and interesting explanations of the basic concepts of chemistry, as well as comprehensive coverage and aid to readers in developing problem solving skills.

1 Atoms and Elements: The Building Blocks of Chemistry.

1.1 Chemistry is important for anyone studying the sciences.

1.2 The scientific method helps us build models of nature.

1.3 Properties of materials can be classified in different ways.

1.4 Materials are described by their properties.

1.5 Atoms of an element have properties in common.

1.6 Atoms are composed of subatomic particles.

1.7 The periodic table is used to organize and correlate facts.

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THINKING IT THROUGH.

2 Compounds and Chemical Reactions.

2.1 Elements combine to form compounds.

2.2 Chemical equations describe what happens in chemical reactions.

2.3 Energy is an important part of chemical change.

2.4 Molecular compounds contain neutral particles called molecules.

2.5 Naming molecular compounds follows a system.

2.6 Ionic compounds are composed of charged particles called ions.

2.7 The formulas of many ionic compounds can be predicted.

2.8 Naming ionic compounds also follows a system.

2.9 Molecular and ionic compounds have characteristic properties.

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3 Measurement.

3.1 Measurements are quantitative observations.

3.2 Measurements always include units.

3.3 Measurements always contain some uncertainty.

3.4 Measurements are written using the significant figures convention.

3.5 Units can be converted using the factor-label method.

3.6 Density is a useful intensive property.

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4 The Mole: Connecting the Macroscopic and Molecular Worlds.

4.1 Use large-scale measurements to count tiny objects.

4.2 The mole conveniently links mass to number of atoms or molecules.

4.3 Chemical formulas relate amounts of substances in a compound.

4.4 Chemical formulas can be determined from experimental mass measurements.

4.5 Chemical equations link amounts of substances in a reaction.

4.6 Chemical equations cannot create or destroy atoms.

4.7 The reactant in shortest supply limits the amount of product.

4.8 The predicted amount of product is not always obtained experimentally.

TOOLS YOU HAVE LEARNED.

THINKING IT THROUGH.

TEST OF FACTS AND CONCEPTS Chapters 1–4.

5 Reactions Between Ions in Aqueous Solutions.

5.1 Special terminology applies to solutions.

5.2 Ionic compounds conduct electricity when dissolved in water.

5.3 Equations for ionic reactions can be written in different ways.

5.4 Reactions that produce precipitates can be predicted.

5.5 Acids and bases are classes of compounds with special properties.

5.6 Naming acids and bases follows a system.

5.7 Acids and bases are classified as strong or weak.

5.8 Neutralization occurs when acids and bases react.

5.9 Gases are formed in some metathesis reactions.

5.10 Predicting metathesis reactions—A summary.

5.11 The composition of a solution is described by its concentration.

5.12 Molarity is used for problems in solution stoichiometry.

5.13 Chemical analysis and titration are applications of solution stoichiometry.

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6 Oxidation–Reduction Reactions.

6.1 Oxidation–reduction reactions involve electron transfer.

6.2 The ion–electron method creates balanced net ionic equations for redox reactions.

6.3 Metals are oxidized when they react with acids.

6.4 A more active metal will displace a less active one from its compounds.

6.5 Molecular oxygen is a powerful oxidizing agent.

6.6 Redox reactions follow the same stoichiometric principles as other reactions.

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7 Energy and Chemical Change: Breaking and Making Bonds.

7.1 Energy is the ability to do work and supply heat.

7.2 Internal energy is the total energy of an object’s molecules.

7.3 Heat can be determined by measuring temperature changes.

7.4 Energy is absorbed or released when chemical bonds are broken or formed.

7.5 Heats of reaction are measured at constant volume or at constant pressure.

7.6 Thermochemical equations are chemical equations that quantitatively include heat.

7.7 Thermochemical equations can be combined because enthalpy is a state function.

7.8 Tabulated standard heats of reaction can be used to predict any heat of reaction using

Hess’s law.

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TEST OF FACTS AND CONCEPTS Chapters 5–7.

8 The Quantum Mechanical Atom.

8.1 Electromagnetic radiation can be described as a wave or as a stream of photons.

8.2 Atomic line spectra are experimental evidence that electrons in atoms have

quantized energies.

8.3 Electron waves in atoms are called orbitals.

8.4 Electron spin affects the distribution of electrons among orbitals in atoms.

8.5 The ground state electron configuration is the lowest energy distribution of electrons among orbitals.

8.6 Electron configurations explain the structure of the periodic table.

8.7 Nodes in atomic orbitals affect their energies and their shapes.

8.8 Atomic properties correlate with an atom’s electron configuration.

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9 Chemical Bonding: General Concepts.

9.1 Electron transfer leads to the formation of ionic compounds.

9.2 Lewis symbols help keep track of valence electrons.

9.3 Covalent bonds are formed by electron sharing.

9.4 Carbon compounds illustrate the variety of structures possible with covalent bonds.

9.5 Covalent bonds can have partial charges at opposite ends.

9.6 The reactivities of metals and nonmetals can be related to their electronegativities.

9.7 Drawing Lewis structures is a necessary skill.

9.8 Formal charges help select correct Lewis structures.

9.9 Resonance applies when a single Lewis structure fails.

9.10 Both electrons in a coordinate covalent bond come from the same atom.

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10 Chemical Bonding and Molecular Structure.

10.1 Molecular shapes are built from five basic arrangements.

10.2 Molecular shapes are predicted using the VSEPR model.

10.3 Polar molecules are asymmetric.

10.4 Valence bond theory explains bonding as an overlap of orbitals.

10.5 Hybrid orbitals are used to explain experimental molecular geometries.

10.6 Hybrid orbitals can be used to explain multiple bonds.

10.7 Molecular orbital theory explains bonding as constructive interference of atomic orbitals.

10.8 Molecular orbital theory uses delocalized orbitals to describe molecules with resonance structures.

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TEST OF FACTS AND CONCEPTS Chapters 8–10.

11 Properties of Gases.

11.1 Familiar properties of gases can be explained at the molecular level.

11.2 Pressure is a measured property of gases.

11.3 The gas laws summarize experimental observations.

11.4 The ideal gas law relates P, V, T, and the number of moles of gas, n.

11.5 In a mixture each gas exerts its own partial pressure.

11.6 Gas volumes are used in solving stoichiometry problems.

11.7 Effusion and diffusion in gases lead to Graham’s law.

11.8 The kinetic-molecular theory explains the gas laws.

11.9 Real gases don’t obey the ideal gas law perfectly.

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12 Intermolecular Attractions and the Properties of Liquids and Solids.

12.1 Gases, liquids, and solids differ because intermolecular forces depend on the distances between molecules.

12.2 Intermolecular attractions involve electrical charges.

12.3 Intermolecular forces and tightness of packing affect the properties of liquids and solids.

12.4 Changes of state lead to dynamic equilibria 520

12.5 Vapor pressures of liquids and solids are controlled by temperature and intermolecular attractions 523

12.6 Boiling occurs when a liquid’s vapor pressure equals atmospheric pressure.

12.7 Energy changes occur during changes of state.

12.8 Changes in a dynamic equilibrium can be analyzed using Le Châtelier’s principle.

12.9 Phase diagrams graphically represent pressure–temperature relationships.

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13 Structures, Properties, and Applications of Solids.

13.1 Crystalline solids have an ordered internal structure.

13.2 X-ray diffraction is used to study crystal structures.

13.3 Physical properties are related to crystal types.

13.4 Band theory explains the electronic structures of solids.

13.5 Polymers are composed of many repeating molecular units.

13.6 Liquid crystals have properties of both liquids and crystals.

13.7 Modern ceramics have applications far beyond porcelain and pottery.

13.8 Nanotechnology deals with controlling structure at the molecular level.

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14 Solutions.

14.1 Substances mix spontaneously when there is no energy barrier to mixing.

14.2 Enthalpy of solution comes from unbalanced intermolecular attractions.

14.3 A substance’s solubility changes with temperature.

14.4 Gases become more soluble at higher pressures.

14.5 Molarity changes with temperature; molality, mass percentages, and mole fractions do not.

14.6 Substances have lower vapor pressures in solution.

14.7 Solutions have lower freezing points and higher boiling points than pure solvents.

14.8 Osmosis is flow of material through a semipermeable membrane due to unequal concentrations.

14.9 Ionic solutes affect colligative properties differently than nonionic solutes.

14.10 Colloids are molecular assemblies of submicrometer dimensions, suspended in a solvent.

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TEST OF FACTS AND CONCEPTS Chapters 11–14.

15 Kinetics:The Study of Rates of Reaction.

15.1 The rate of a reaction is the change in reactant or product concentrations with time.

15.2 Five factors affect reaction rates.

15.3 Rates of reaction are measured by monitoring change in concentration over time.

15.4 Rate laws give reaction rate as a function of reactant concentrations.

15.5 Integrated rate laws give concentration as a function of time.

15.6 Reaction rate theories explain experimental rate laws in terms of molecular collisions.

15.7 Activation energies are measured by fitting experimental data to the Arrhenius equation.

15.8 Experimental rate laws can be used to support or reject proposed mechanisms for a reaction.

15.9 Catalysts change reaction rates by providing alternative paths between reactants and products.

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16 Chemical Equilibrium—General Concepts.

16.1 Dynamic equilibrium is achieved when the rates of two opposing processes are equal.

16.2 Closed systems reach the same equilibrium concentrations whether we start with reactants or with products.

16.3 A law relating equilibrium concentrations can be derived from the balanced chemical equation for a reaction.

16.4 Equilibrium laws for gaseous reactions can be written in terms of concentrations or pressures.

16.5 A large K means a product-rich equilibrium mixture; a small K means a reactant-rich mixture at equilibrium.

16.6 A simple expression relates KP and Kc.

16.7 Heterogeneous equilibria involve reaction mixtures with more than one phase.

16.8 When a system at equilibrium is stressed, it reacts to relieve the stress.

16.9 Equilibrium concentrations can be used to predict equilibrium constants, and vice versa.

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17 Acids and Bases: A Second Look.

17.1 Brønsted–Lowry acids and bases exchange protons.

17.2 Strengths of Brønsted acids and bases follow periodic trends.

17.3 Lewis acids and bases involve coordinate covalent bonds.

17.4 Elements and their oxides demonstrate acid–base properties.

17.5 pH is a measure of the acidity of a solution.

17.6 Strong acids and bases are fully dissociated in solution.

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18 Equilibria in Solutions of Weak Acids and Bases.

18.1 Ionization constants can be defined for weak acids and bases.

18.2 Calculations can involve finding or using Ka and Kb.

18.3 Salt solutions are not neutral if the ions are weak acids or bases.

18.4 Simplifications fail for some equilibrium calculations.

18.5 Buffers enable the control of pH.

18.6 Polyprotic acids ionize in two or more steps.

18.7 Salts of polyprotic acids give basic solutions.

18.8 Acid–base titrations have sharp changes in pH at the equivalence point.

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19 Solubility and Simultaneous Equilibria.

19.1 An insoluble salt is in equilibrium with the solution around it.

19.2 Solubility equilibria of metal oxides and sulfides involve reactions with water.

19.3 Metal ions can be separated by selective precipitation.

19.4 Complex ions participate in equilibria in aqueous solutions.

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TEST OF FACTS AND CONCEPTS Chapters 15–19.

20 Thermodynamics.

20.1 Internal energy can be transferred as heat or work, but it cannot be created or destroyed.

20.2 A spontaneous change is a change that continues without outside intervention.

20.3 Spontaneous processes tend to proceed from states of low probability to states of higher probability.

20.4 All spontaneous processes increase the total entropy of the universe.

20.5 The third law of thermodynamics makes experimental measurement of absolute entropies possible.

20.6 The standard free energy change, G°, is G at standard conditions.

20.7 G is the maximum amount of work that can be done by a process.

20.8 G is zero when a system is at equilibrium.

20.9 Equilibrium constants can be estimated from standard free energy changes.

20.10 Bond energies can be estimated from reaction enthalpy changes.

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21 Electrochemistry.

21.1 Galvanic cells use redox reactions to generate electricity.

21.2 Cell potentials can be related to reduction potentials.

21.3 Standard reduction potentials can predict spontaneous reactions.

21.4 Cell potentials are related to free energy changes.

21.5 Concentrations in a galvanic cell affect the cell potential.

21.6 Batteries are practical examples of galvanic cells.

21.7 Electrolysis uses electrical energy to cause chemical reactions.

21.8 Stoichiometry of electrochemical reactions involves electric current and time.

21.9 Electrolysis has many industrial applications.

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22 Nuclear Reactions and Their Role in Chemistry.

22.1 Mass and energy are conserved in all of their forms.

22.2 The energy required to break a nucleus into separate nucleons is called the nuclear binding energy.

22.3 Radioactivity is an emission of particles and/or electromagnetic radiation by unstable atomic nuclei.

22.4 Stable isotopes fall within the "band of stability" on a plot based on numbers of protons and neutrons.

22.5 Transmutation is the change of one isotope into another.

22.6 How is radiation measured?

22.7 Radionuclides have many medical and analytical applications.

22.8 Nuclear fission is the breakup of a nucleus into two fragments of comparable size after capture of a slow neutron.

TOOLS YOU HAVE LEARNED.

TEST OF FACTS AND CONCEPTS Chapters 20–22.

23 Metallurgy and the Properties of Metals and Metal Complexes.

23.1 Metals are prepared from compounds by reduction.

23.2 Metallurgy is the science and technology of metals.

23.3 Metal compounds exhibit varying degrees of covalent bonding.

23.4 Complex ions are formed by many metals.

23.5 The nomenclature of metal complexes follows an extension of the rules developed earlier.

23.6 Coordination number and structure are often related.

23.7 Isomers of coordination complexes are compounds with the same formula but different structures.

23.8 Bonding in transition metal complexes involves d orbitals.

23.9 Metal ions serve critical functions in biological systems.

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24 Some Chemical Properties of the Nonmetals and Metalloids.

24.1 Metalloids and nonmetals are found as free elements and in compounds.

24.2 The free elements have structures of varying complexity.

24.3 Hydrogen forms compounds with most nonmetals and metalloids.

24.4 Catenation occurs when atoms of the same element bond to each other.

24.5 Oxygen combines with almost all nonmetals and metalloids.

24.6 Nonmetals form a variety of oxoacids and oxoanions.

24.7 Halogen compounds are formed by most nonmetals and metalloids.

25 Organic Compounds and Biochemicals.

25.1 Organic chemistry is the study of carbon compounds.

25.2 Hydrocarbons consist of only C and H atoms.

25.3 Alcohols and ethers are organic derivatives of water.

25.4 Amines are organic derivatives of ammonia.

25.5 Organic compounds with carbonyl groups include aldehydes, ketones, and carboxylic acids.

25.6 Most biochemicals are organic compounds.

25.7 Carbohydrates include sugars, starch, and cellulose.

25.8 Lipids comprise a family of water-insoluble compounds.

25.9 Proteins are almost entirely polymers of amino acids.

25.10 Nucleic acids carry our genetic information.

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Appendices.

A. Electron Configurations of the Elements.

B. Answers to Practice Exercises and Selected Review Exercises.

C. Tables of Selected Data .

Glossary.

Photo Credits.

Index.

Amazon no longer offers textbook rentals. We do!

Chemistry : The Study of Matter and Its Changes

Chemistry : The Study of Matter and Its Changes

9780471215172

0471215171

Summary

Author Biography

Table of Contents

Supplemental Materials

Write a Review

Amazon no longer offers textbook rentals. We do!

Chemistry : The Study of Matter and Its Changes

Chemistry : The Study of Matter and Its Changes

9780471215172

0471215171

Summary

Author Biography

Table of Contents

Supplemental Materials

Write a Review

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