Course Syllabus for English-Taught Majors

General Chemistry (Inorganic Chemistry) ’Course Syllabus


Course Code09041031

Course Category:大类基础课程(Major Basic



Total Hours72 Hours         Credit4 学分

Lecture Hours90 Hours         

InstructorsProfessor Chuanjiang Hu

Textbooks普通化学原理与应用(8版影印版)(General Chemistry:Principles and Modern Applications) 作者:(美国)彼德勒  


“Inorganic Chemistry”, Housecroft, C. E.; Sharpe, A. G.

无机化学(3版影印版)/“Inorganic Chemistry” Miessler and Tarr

“Instant Notes in Inorganic Chemistry” P. A. Pox

“Basic Inorganic Chemistry” Albert F. Cotton





Teaching Aim

    Since this freshman level chemistry course is offered to students with a wide variety of majors, it is a unique mixture of inorganic and physical chemistry with very little analytical chemistry. It will include the following topics from these areas of chemistry over a period of two semesters: Inorganic Chemistry : atomic structure, molecular structure, valence theory, atomic and molecular orbital theory, periodicity, Acids and Bases, models of ionic and covalent bonding, and bonding properties of matter, Metals, Nonmetals, and Metalloids,. Physical Chemistry: Quantum Theory and Electronic Structure of Atoms, Gases and Kinetic Molecular Theory, Chemical Thermodynamics, Chemical Kinetics, Chemical equilibrium, Buffers and Titration Curves.


Chapter 1             Matter

   Hourshalf week2.5 hours



1-1  The Scientific Method

1-2  Properties of Matter

Matter is anything that occupies space and displays the properties of mass

and inertia.


1-3  Classification of Matter

Matter is made up of very tiny units called atoms. Each different type of atom is the building block of a different chemical element.

Elements and compounds are called substances.

A mixture of substances can vary in composition and properties from one sample to another.


1-4 Measurement of Matter: SI (Metric) Units 8

1-5 Density and Percent Composition: Their Use in Problem Solving 13

1-6 Uncertainties in Scientific Measurements 18

1-7 Significant Figures 19


In your own words, define or explain the following

terms or symbols: (a) mL; (b) % by mass; (c) °C;

(d) density; (e) element

Chapter 2             Atom and Atomic Theory

      Hourshalf week2.5 hours

2-1 Early Chemical Discoveries and the Atomic Theory

Law of Conservation of Mass

Law of Constant Composition

Dalton s Atomic Theory

Law of multiple proportions


2-2 Electrons and Other Discoveries in Atomic Physics

cathode rays

2-3 The Nuclear Atom

Properties of Protons, Neutrons, and Electrons

2-4 Chemical Elements

All atoms of a particular element have the same atomic number, Z, and,

conversely, all atoms with the same number of protons are atoms of the

same element.


2-5 Atomic Mass

he atomic mass (weight)* of an element is the average of

the isotopic masses, weighted according to the naturally occurring abundances

of the isotopes of the element.

2-6 Introduction to the Periodic Table

Features of the Periodic Table

2-7 The Concept of the Mole and the Avogadro Constant


mole is the amount of a substance that contains the same number of elementary

entities as there are atoms in exactly 12 g of pure carbon-12. The number

of elementary entities (atoms, molecules, and so on) in a mole is the

Avogadro constant,


The Avogadro constant consists of a number, 6.02214179 1023, known as Avogadro s number, and a unit, . The unit signifies that the entities being counted are those present in 1 mole.

The value of Avogadro s number is based on both a definition and a measurement.

Amole of carbon-12 is defined to be 12 g. If the mass of one carbon-12 atom is measured by using a mass spectrometer , the mass would be about g. The ratio of these two masses provides an estimate of Avogadro s number.

NA = 6.02214179 * 1023mol-1

2-8 Using the Mole Concept in Calculations

A solution was prepared by dissolving 2.50 g potassium
chlorate (a substance used in fireworks and flares) in 100.0 mL water at 40 °C. When the solution was cooled to 20 °C, its volume was still found to be
100.0 mL, but some of the potassium chlorate had
crystallized (deposited from the solution as a solid).
At 40 °C, the density of water is and at 20 °C, the potassium chlorate solution had a density of
(a) Estimate, to two significant figures, the mass of potassium perchlorate that crystallized.
(b) Why can’t the answer in (a) be given more precisely?
Chapter 3             Chemical Compounds
   Hours1 week5 hours

3-1 Types of Chemical Compounds and Their Formulas

Molecular Compounds

Ionic Compounds

3-2 The Mole Concept and Chemical Compounds

The molar mass is the mass of one mole of compound one mole

of molecules of a molecular compound and one mole of formula units of an

ionic compound.

3-3 Composition of Chemical Compounds

3-4 Oxidation States: A Useful Tool in Describing Chemical Compounds

3-5 Naming Compounds: Inorganic Compounds

Ionic compounds

Molecular compounds
Oxoacids and their salts

Acetaminophen, an analgesic and antipyretic drug, has a molecular mass of 151.2 u and a mass

percent composition of 63.56% C, 6.00% H, 9.27% N, and 21.17% O. What is the molecular formula of acetaminophen?

Chapter 4             Chemical Reactions
Hours1 week5 hours

4-1 Chemical Reactions and Chemical Equations

4-2 Chemical Equations and Stoichiometry

4-3 Chemical Reactions in Solution

4-4 Determining the Limiting Reactant

The reactant that is completely consumed in a chemical reaction is called the limiting reactant (or limiting reagent) because it determines (or limits) the amount of product formed.



Write a chemical equation to represent the complete combustion of malonic acid, a compound with

34.62% C, 3.88% H, and 61.50% O, by mass.


Chapter 5             Reactions in Aqueous Solutions
Hours1 week5 hours

5-1 The Nature of Aqueous Solutions

A solution is a homogenous mixture of two or more substances, consisting of

The solvent - usually the substance in greater concentration

The other component(s) is (are) called the solute(s) - they are said to be dissolved in the solvent

5-2 Precipitation Reactions

Metathesis reactions that result in an insoluble precipitate are called precipitation reactions

5-3 Acid Base Reactions

5-4 Oxidation Reduction Reactions: Some General Principles

When one substance in a reaction loses electrons, another substance in the reaction must gain them. The oxidation of a reactant is always accompanied by the reduction of another reactant in the reaction, These types of reactions are therefore called oxidation/reduction (or REDOX) reactions


5-5 Balancing Oxidation Reduction Equations

5-6 Oxidizing and Reducing Agents

5-7 Stoichiometry of Reactions in Aqueous Solutions: Titrations



The active component in one type of calcium dietary supplement is calcium carbonate. A 1.2450 g tablet of the supplement is added to 50.00 mL of 0.5000 M HCl and allowed to react. After completion of the reaction,the excess HCl(aq) requires 40.20 mL of 0.2184 M NaOH for its titration to the equivalence point. What is the calcium content of the tablet, expressed in milligrams of Ca2+?

Chapter 6             Gases
Hours1 week4 hours

6-1 Properties of Gases:Gas Pressure

6-2 The Simple Gas Laws

The Pressure-Volume Relationship: Boyle's Law

The Temperature-Volume Relationship: Charles's Law

The Quantity-Volume Relationship: Avogadro's Law

6-3 Combining the Gas Laws: The Ideal Gas Equation and the General Gas Equation

6-4 Applications of the Ideal Gas Equation


6-5 Gases in Chemical Reactions

6-6 Mixtures of Gases

The total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone

6-7 Kinetic-Molecular Theory of Gases

6-8 Gas Properties Relating to the Kinetic-Molecular Theory

The Kinetic-Molecular Theory ("the theory of moving molecules"; Rudolf Clausius, 1857)

Gases consist of large numbers of molecules (or atoms, in the case of the noble gases) that are in continuous, random motion

The volume of all the molecules of the gas is negligible compared to the total volume in which the gas is contained

Attractive and repulsive forces between gas molecules is negligible

The average kinetic energy of the molecules does not change with time (as long as the temperature of the gas remains constant). Energy can be transferred between molecules during collisions (but the collisions are perfectly elastic)

The average kinetic energy of the molecules is proportional to absolute temperature. At any given temperature, the molecules of all gases have the same average kinetic energy. In other words, if I have two gas samples, both at the same temperature, then the average kinetic energy for the collection of gas molecules in one sample is equal to the average kinetic energy for the collection of gas molecules in the other sample.

6-9 Nonideal (Real) Gases


Two evacuated bulbs of equal volume are connected by a tube of negligible volume. One of the bulbs is placed in a constant-temperature bath at 225 K and the other bulb is placed in a constant-temperature bath at 350 K. Exactly 1 mol of an ideal gas is injected into the system. Calculate the final number of moles of gas in each bulb.

Chapter 7             Thermochemistry
Hours1 week5 hours

7-1 Getting Started: Some Terminology

7-2 Heat

7-3 Heats of Reaction and Calorimetry

7-4 Work

7-5 The First Law of Thermodynamics

The energy of an isolated system is constant.


7-6 Heats of Reaction:

7-7 Indirect Determination of DH; Hess s Law

if a reaction is carried out in a series of steps, DH for the reaction will be equal to the sum of the enthalpy changes for the individual steps

7-8 Standard Enthalpies of Formation

7-9 Fuels as Sources of Energy



ABritish thermal unit (Btu) is defined as the quantity of heat required to change the temperature of 1 lb of water by 1° F. Assume the specific heat of water to be independent of temperature. How much heat is required to raise the temperature of the water in a 40 gal water heater from 48 to 145° F in (a) Btu; (b) kcal; (c) k?

he energy released when 1 gram of material is combusted is called its fuel value

Chapter 8 Atmosphere and Hydrogen
Hourshalf week2.5 hours
8-1 Atomosphere
8-6 Hydrogen
What are the application of hydrogen gases?
Chapter 9             Electrons in Atoms
Hours1.5 week 7.5 hours

9-1 Electromagnetic Radiation

9-2 Atomic Spectra

9-3 Quantum Theory

9-4 The Bohr Atom

·         Bohr began with the assumption that electrons were orbiting the nucleus, much like the earth orbits the sun.

·         From classical physics, a charge traveling in a circular path should lose energy by emitting electromagnetic radiation

·         If the "orbiting" electron loses energy, it should end up spiraling into the nucleus (which it does not). Therefore, classical physical laws either don't apply or are inadequate to explain the inner workings of the atom

·         Bohr borrowed the idea of quantized energy from Planck

o    He proposed that only orbits of certain radii, corresponding to defined energies, are "permitted"

o    An electron orbiting in one of these "allowed" orbits:

§  Has a defined energy state

§  Will not radiate energy

§  Will not spiral into the nucleus

If the orbits of the electron are restricted, the energies that the electron can possess are likewise restricted and are defined by the equation:

Where RH is a constant called the Rydberg constant and has the value

2.18 x 10-18 J

'n' is an integer, called the principle quantum number and corresponds to the different allowed orbits for the electron.

9-5 Two Ideas Leading to a New Quantum Mechanics

9-6 Wave Mechanics

9-7 Quantum Numbers and Electron Orbitals

Schrödinger's wave equation incorporates both wave- and particle-like behaviors for the electron. 
Schrödinger's equation results in a series of so called wave functions, represented by the letter j (psi). Although has no actual physical meaning, the value of j2 describes the probability distribution of an electron.


9-8 Interpreting and Representing the Orbitals of the Hydrogen Atom

s,p,d,f orbitals

9-9 Electron Spin: A Fourth Quantum Number

The electron spin quantum number may have a value of +1/2 or -1/2; the value of ms does not depend on any of the other three quantum numbers.

9-10 Multielectron Atoms

Penetration and Shielding


9-11 Electron Configurations

Rules for Assigning Electrons to Orbitals

1. Electrons occupy orbitals in a way that minimizes the energy of the atom.

2. No two electrons in an atom can have all four quantum numbers alike

-the Pauli exclusion principle.

3. When orbitals of identical energy (degenerate orbitals) are available,

electrons initially occupy these orbitals singly.- Hund’s rule,

9-12 Electron Configurations and the Periodic Table

The periodic table is structured so that elements with the same type of valence electron configuration are arranged in columns.



Explain the important distinctions between each pair of terms: (a) frequency and wavelength; (b) ultraviolet and infrared light; (c) continuous and discontinuous spectra; (d) traveling and standing waves; (e) quantum number and orbital; (f) spdf notation and orbital diagram; (g) s block and p block; (h) main group and transition element; (i) the ground state and excited state of a hydrogen atom.


Chapter 10           The Periodic Table and Some Atomic Properties
Hours1 week5 hours

10-1 Classifying the Elements: The Periodic Law and the Periodic Table

1869: Dmitri Mendeleev and Lothar Meyer published schemes for classifying elements

The elements could be ordered according to their atomic weight (i.e. grams/mole for the naturally occuring mixture of isotopic forms) which resulted in periodic characteristics

The accuracy of Mendeleev's predictions for undiscovered elements, based on his periodic table, convinced scientists of its validity


10-2 Metals and Nonmetals and Their Ions

Metallic Elements

Nonmetallic elements

Distinguishing luster (shine)

Non-lustrous, various colors

Malleable and ductile (flexible) as solids

Brittle, hard or soft

Conduct heat and electricity

Poor conductors

Metallic oxides are basic, ionic

Nonmetallic oxides are acidic, compounds

Cations in aqueous solution

Anions, oxyanions in aqueous solution

10-3 Sizes of Atoms and Ions

10-4 Ionization Energy

The ionization energy, I, is the quantity of energy a gaseous atom must absorb to

be able to expel an electron.


10-5 Electron Affinity

Electron affinity, EA, can be defined as the enthalpy change, that occurs when an atom in the gas phase gains an electron.

10-6 Magnetic Properties

Diamagnetic and paramagnetic

10-7 Periodic Properties of the Elements


Explain why the third ionization energy of Li(g) is an easier quantity to calculate than either the first or second ionization energies. Calculate I3 for Li, and express the result in kJ/mol.

Chapter 11           Chemical Bonding I
Hours2 week10 hours

11-1 Lewis Theory: An Overview

Lewis Symbols and the Octet Rule 
Valence electrons reside in the outer shell and are the electrons which are going to be involved in chemical interactions and bonding (valence comes from the Latin valere, "to be strong"). 
Electron-dot symbols (Lewis symbols): 
convenient representation of valence electrons 
allows you to keep track of valence electrons during bond formation 

consists of the chemical symbol for the element plus a dot for each valence electron

Atoms often gain, lose, or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table

Because all noble gasses (except He) have filled s and p valence orbitals (8 electrons), many atoms undergoing reactions also end up with 8 valence electrons. This observation has led to the Octet Rule:

11-2 Covalent Bonding: An Introduction

Covalent Bonding
G.N. Lewis reasoned that an atom might attain a noble gas electron configuration by sharing electrons
A chemical bond formed by sharing a pair of electrons is called a covalent bond
11-3 Polar Covalent Bonds and Electrostatic Potential Maps
Bond Polarity
Bond polarity is a useful concept for describing the sharing of electrons between atoms 
A nonpolar covalent bond is one in which the electrons are shared equally between two atoms 

A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. If this relative attraction is great enough, then the bond is an ionic bond

11-4 Writing Lewis Structures

The general procedure...

1. Sum the valence electrons from all atoms

·         Use the periodic table for reference

·         Add an electron for each indicated negative charge, subtract an electron for each indicated positive charge

2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond

·         You may need some additional evidence to decide bonding interactions

·         If a central atom has various groups bonded to it, it is usually listed first: CO32-, SF4

·         Often atoms are written in the order of their connections: HCN

3. Complete the octets of the atoms bonded to the central atom (H only has two)

4. Place any leftover electrons on the central atom (even if it results in more than an octet)

5. If there are not enough electrons to give the central atom an octet, try multiple bonds (use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds)

11-5 Resonance

Equivalent Lewis structures are called resonance structures, or resonance forms

11-6 Exceptions to the Octet Rule

There are three general ways in which the octet rule breaks down:

1. Molecules with an odd number of electrons

2. Molecules in which an atom has less than an octet

3. Molecules in which an atom has more than an octet

11-7 Shapes of Molecules

These structures can generally be predicted, when A is a nonmetal, using the "valence-shell electron-pair repulsion model (VSEPR)

The best spatial arrangement of the bonding pairs of electrons in the valence orbitals is one in which the repulsions are minimized

11-8 Bond Order and Bond Lengths

11-9 Bond Energies

Bond energy is the enthalpy change required to break a bond (in 1 mole of a gaseous substance)


The following statements are not made as carefully as they might be. Criticize each one.

(a) Lewis structures with formal charges are incorrect.

(b) Triatomic molecules have a planar shape.

(c) Molecules in which there is an electronegativity difference between the bonded atoms are polar.

Chapter 12 Chemical Bonding II

Hours1 week5 hours


12-1 What a Bonding Theory Should Do

12-2 Introduction to the Valence-Bond Method

Combine Lewis' idea of electron pair bonds with electron orbitals (quantum mechanics)

12-3 Hybridization of Atomic Orbitals

Hybrid Orbitals

 the two steps often observed when constructing hybrid orbitals is to 1) promote a valence electron from the ground state configuration to a higher energy orbital, and then 2) hybridize the appropriate valence electron orbitals to achieve the desired valence electron geometry (i.e. the correct number of hybrid orbitals for the appropriate valence electron geometry)

12-4 Multiple Covalent Bonds

Generally speaking:

A single bond is composed of a s bond

A double bond is composed of one s bond and one p bond

A triple bond is composed of one s bond and two p bonds

12-5 Molecular Orbital Theory

Basic Ideas Concerning Molecular Orbitals

Here are some useful ideas about molecular orbitals and how electrons are

assigned to them.

1. The number of molecular orbitals (MOs) formed is equal to the number of

atomic orbitals combined.

2. Of the two MOs formed when two atomic orbitals are combined, one is a

bondingMO at a lower energy than the original atomic orbitals. The other is

an antibonding MO at a higher energy.

3. In ground-state configurations, electrons enter the lowest energy MOs


4. The maximum number of electrons in a given MO is two (Pauli exclusion

principle, page 339).

5. In ground-state configurations, electrons enter MOs of identical energies

singly before they pair up

12-6 Delocalized Electrons: Bonding in the Benzene Molecule


Lewis theory is satisfactory to explain bonding in the ionic compound but it does not readily explain

formation of the ionic compounds potassium superoxide, and potassium peroxide,

(a) Show that molecular orbital theory can provide this explanation.

(b) Write Lewis structures consistent with the molecular orbital explanation.

Chapter 13 Liquids, Solids and Intermolecular Forces
Hours1 week5 hours

13-1 Intermolecular Forces

Intermolecular Forces are generally much weaker than covalent bonds Attractive forces between neutral molecules

Dipole-dipole forces

London dispersion forces

Hydrogen bonding forces

Typically, dipole-dipole and dispersion forces are grouped together and termed van der Waals forces (sometimes the hydrogen bonding forces are also included with this group)

13-2 Some Properties of Liquids

Surface tension is the energy required to increase the surface area of a liquid by a unit amount

Viscosity: The resistance of a liquid to flow is called its viscosity

13-3 Some Properties of Solids

Melting, Melting Point, and Heat of Fusion


13-4 Phase Diagrams

Phases and Phase Transitions

13-5 Network Covalent Solids and Ionic Solids

13-6 Crystal Structures

Crystal Lattices

Ionic Crystal Structures

13-7 Energy Changes in the Formation of Ionic Crystals

Lattice Energy,


In acetic acid vapor, some molecules exist as monomers and some as dimers (see Figure 12-9). If the density of the vapor at 350 K and 1 atm is 3.23g/L, what percentage of the molecules must exist as dimers? Would you expect this percent to increase or decrease with temperature?

Chapter 14 Solutions and Their Physical Properties
Hours1 week5 hours

14-1 Types of Solutions:

Some Terminology

14-2 Solution Concentration

Mass Percent, Volume Percent, and Mass/Volume Percent

Mole Fraction and Mole Percent, Molarity


14-3 Intermolecular Forces and the Solution Process

Enthalpy of Solution

Intermolecular Forces in Mixtures

14-4 Solution Formation and Equilibrium

Solubility as a Function of Temperature

14-5 Solubilities of Gases

Effect of Temperature

Effect of Pressure

14-6 Vapor Pressures of Solutions

Raoult s law states that the partial pressure exerted by solvent vapor

above an ideal solution, is the product of the mole fraction of solvent in

the solution, and the vapor pressure of the pure solvent at the given temperature,


Liquid Vapor Equilibrium: Ideal Solutions

Fractional Distillation

Liquid Vapor Equilibrium: Nonideal Solutions

14-7 Osmotic Pressure

Applying pressure to the sucrose solution slows down the net flow of water

across the membrane into the solution. With a sufficiently high pressure, the

net influx of water can be stopped altogether. The necessary pressure to stop

osmotic flow is called the osmotic pressure of the solution.

14-8 Freezing-Point Depression and Boiling-Point Elevation of Nonelectrolyte Solutions

14-9 Solutions of Electrolytes


Demonstrate that

(a) for a dilute aqueous solution, the numerical value of the molality is essentially equal to that of the molarity.

(b) in a dilute solution, the solute mole fraction is proportional to the molality.

(c) in a dilute aqueous solution, the solute mole fraction is proportional to the molarity.

Chapter 15 Chemical Kinetics
Hours1 week5 hours

15-1 Rate of a Chemical Reaction

15-2 Measuring Reaction Rates

Rate of Reaction Expressed as Concentration

Change over Time

15-3 Effect of Concentration on Reaction Rates: The Rate Law

Method of Initial Rates

15-4 Zero-Order Reactions

15-5 First-Order Reactions

15-6 Second-Order Reactions

15-7 Reaction Kinetics: A Summary

1. To calculate a rate of reaction when the rate law is known, use this expression:

rate of reaction

2. To determine a rate of reaction when the rate law is not given, use

the slope of an appropriate tangent line to the graph of versus t

the expression -D[A]/ Dt, with a short time interval Dt

3. To determine the order of a reaction, use one of the following methods.

Use the method of initial rates if the experimental data are given in the

form of reaction rates at different initial concentrations.

Find the graph of rate data that yields a straight line (Table 14.5).

Test for the constancy of the half-life (good only for first-order).

Substitute rate data into integrated rate laws to find the one that gives a

constant value of k.

4. To find the rate constant k for a reaction, use one of the following methods.

Obtain k from the slope of a straight-line graph.

Substitute concentration time data into the appropriate integrated rate law.

Obtain k from the half-life of the reaction (good only for a first-order


5. To relate reactant concentrations and times, use the appropriate integrated

rate law after first determining k.

15-8 Theoretical Models for Chemical Kinetics

Collision Theory

Transition State Theory

15-9 The Effect of Temperature on Reaction Rates

15-10 Reaction Mechanisms

Equilibrium method

The Steady-State Approximation

15-11 Catalysis

A reaction can generally be made to go faster by increasing the temperature.

Another way to speed up a reaction is to use a catalyst. A catalyst provides an

alternative reaction pathway of lower activation energy.


Show that the following mechanism is consistent with

the rate law established for the iodide hypochlorite

reaction in Exercise 79.

Chapter 16 Principles of Chemical Equilibrium
Hours1 week 5 hours

16-1 Dynamic Equilibrium

a system at equilibrium

two opposing processes take place at equal rates.

16-2 The Equilibrium Constant Expression

16-3 Relationships Involving Equilibrium Constants

16-4 The Magnitude of an Equilibrium Constant

16-5 The Reaction Quotient, Q: Predicting the Direction of Net Change

16-6 Altering Equilibrium Conditions: Le Châtelier s Principle

When an equilibrium system is subjected to a change in temperature,

pressure, or concentration of a reacting species, the system responds by

attaining a new equilibrium that partially offsets the impact of the change.


Briefly describe each of the following ideas or phenomena:

(a) dynamic equilibrium; (b) direction of a net chemical change; (c) Le Châtelier s principle;

(d) effect of a catalyst on equilibrium.

Chapter 17 Acids and Bases
Hours1 week5 hours

17-1 Arrhenius Theory: A Brief Review

17-2 Brønsted Lowry Theory of Acids and Bases

According to their theory, an acid is a proton donor and a base is a proton acceptor.

17-3 Self-Ionization of Water and the pH Scale

17-4 Strong Acids and Strong Bases

17-5 Weak Acids and Weak Bases

ionization of a weak acid, is a reversible reaction that reaches a condition of equilibrium.

17-6 Polyprotic Acids

But some acids have more than one ionizable H atom per molecule. These are polyprotic acids.

17-7 Ions as Acids and Bases

17-8 Molecular Structure and Acid Base Behavior Strengths of Binary Acids

When comparing binary acids of elements in the same row of the periodic

table, acid strength increases as the polarity of the bond increases.

Strengths of Organic Acids

17-9 Lewis Acids and Bases

A Lewis acid is a species (an atom, ion, or molecule) that is an electron-pair

acceptor, and a Lewis base is a species that is an electron-pair donor.


It is possible to write simple equations to relate pH, pK, and molarities (M) of various solutions. Three such equations are shown here.

(a) Derive these three equations, and point out the assumptions involved in the derivations.

(b) Use these equations to determine the pH of 0.10 M CH3COOH(aq), 0.10M NH3(aq) and ).10 M NaCH3COO. Verify that the equations give correct results by determining these pH values in the usual way.

Chapter 18           Additional Aspects of Acid-Base Equilibria   
Hours1 week5 hours

18-1 Common-Ion Effect in Acid Base Equilibria

Solutions of Weak Acids and Strong Acids

18-2 Buffer Solutions

common buffer solutions are

described as combinations of a weak acid and its conjugate base, or

a weak base and its conjugate acid

18-3 Acid Base Indicators

An acid base indicator is a substance whose color depends on the pH of the

solution to which it is added.

18-4 Neutralization Reactions and Titration Curves

18-5 Solutions of Salts of Polyprotic Acids

18-6 Acid Base Equilibrium Calculations: A Summary


A solution is prepared that is 0.150 M CH3COOH and 0.250 M NaHCOO

(a) Show that this is a buffer solution.

(b) Calculate the pH of this buffer solution.

(c) What is the final pH if 1.00 L of 0.100 M HCl is added to 1.00 L of this buffer solution?


考核方式 Assessment Methods:闭卷


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