‘General Chemistry (Inorganic Chemistry) ’Course Syllabus
Course Category：大类基础课程（Major Basic）
Total Hours：72 Hours Credit：4 学分
Lecture Hours：90 Hours
Instructors：Professor Chuanjiang Hu
Textbooks：普通化学原理与应用(第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(第3版影印版)/“Inorganic Chemistry” Miessler and Tarr，
“Instant Notes in Inorganic Chemistry” P. A. Pox
“Basic Inorganic Chemistry” Albert F. Cotton
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
Hours：half week，2.5 hours
1-1 The Scientific Method
1-2 Properties of Matter
Matter is anything that occupies space and displays the properties of mass
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
Hours：half week，2.5 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
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
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
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
Hours：1 week，5 hours
3-1 Types of Chemical Compounds and Their Formulas
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
3-3 Composition of Chemical Compounds
3-4 Oxidation States: A Useful Tool in Describing Chemical Compounds
3-5 Naming Compounds: Inorganic Compounds
Ionic compoundsMolecular compoundsOxoacids and their saltsProblem:
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 ReactionsHours：1 week，5 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 ReactantThe 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 SolutionsHours：1 week，5 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 GasesHours：1 week，4 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 ThermochemistryHours：1 week，5 hours
7-1 Getting Started: Some Terminology
7-3 Heats of Reaction and Calorimetry
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 valueChapter 8 Atmosphere and HydrogenHours：half week，2.5 hours8-1 Atomosphere8-6 HydrogenProblem:What are the application of hydrogen gases?Chapter 9 Electrons in AtomsHours：1.5 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 OrbitalsSchrödinger'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
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 PropertiesHours：1 week，5 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
Distinguishing luster (shine)
Non-lustrous, various colors
Malleable and ductile (flexible) as solids
Brittle, hard or soft
Conduct heat and electricity
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 ElementsProblem:
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 IHours：2 week，10 hours
11-1 Lewis Theory: An OverviewLewis Symbols and the Octet RuleValence 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 electronsallows 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 IntroductionCovalent BondingG.N. Lewis reasoned that an atom might attain a noble gas electron configuration by sharing electronsA chemical bond formed by sharing a pair of electrons is called a covalent bond11-3 Polar Covalent Bonds and Electrostatic Potential MapsBond PolarityBond polarity is a useful concept for describing the sharing of electrons between atomsA 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)
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)Problem:
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
Hours：1 week，5 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
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
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 ForcesHours：1 week，5 hours
13-1 Intermolecular Forces
Intermolecular Forces are generally much weaker than covalent bonds Attractive forces between neutral molecules
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
Ionic Crystal Structures
13-7 Energy Changes in the Formation of Ionic Crystals
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 PropertiesHours：1 week，5 hours
14-1 Types of Solutions:
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
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
(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 KineticsHours：1 week，5 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
Transition State Theory
15-9 The Effect of Temperature on Reaction Rates
15-10 Reaction Mechanisms
The Steady-State Approximation
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 EquilibriumHours：1 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.Problem:
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 BasesHours：1 week，5 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 EquilibriaHours：1 week，5 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?
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