(Ebook PDF) Chemistry in Quantitative Language Fundamentals of General Chemistry Calculations 2nd Edition by Christopher Oriakhi 0192638041 9780192638045 full chapters

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ISBN-10 :  0192638041 

ISBN-13 :  9780192638045

Author : Christopher O. Oriakhi 

Problem-solving is one of the most challenging aspects students encounter in general chemistry courses, leading to frustration and failure. Consequently, many students become less motivated to take additional chemistry courses after the first year. This book tackles this issue head on and provides innovative, intuitive, and systematic strategies to tackle any type of calculations encountered in chemistry. The material begins with the basic theories, equations, and concepts of the underlying chemistry, followed by worked examples with carefully explained step-by-step solutions to showcase the ways in which the problems can be presented. The second edition contains additional problems at the end of each chapter with varying degrees of difficulty, and many of the original examples have been revised.

Chemistry in Quantitative Language: Fundamentals of General Chemistry Calculations 2nd Table of contents:

1: Atomic Structure and Isotopes
1.1 Atomic Theory
1.1.1 The law of conservation of mass
1.1.2 The law of definite proportion
1.1.3 The law of multiple proportions
1.1.4 The law of reciprocal proportions
1.2 The Structure of the Atom
1.2.1 Atomic number (Z)
1.2.2 Mass number (A)
1.2.3 Ions
1.3 Isotopes
1.4 Relative Atomic Mass
1.4.1 Calculating atomic masses
1.5 Problems
2: Formula and Molecular Mass
2.1 Formula Mass
2.2 Molecular Mass
2.3 Molar Mass
2.4 Problems
3: The Mole and Avogadro’s Number
3.1 The Mole and Avogadro’s Number (NA)
3.2 The Mole and Molar Mass
3.3 Calculating the Number of Moles
3.4 Problems
4: Formulas of Compounds and Percent Composition
4.1 Percent Composition
4.2 Types of Chemical Formula
4.2.1 Empirical formula
4.2.2 Steps for determining empirical formula
4.3 Empirical Formula from Combustion Analysis
4.4 Molecular Formula
4.4.1 Determination of molecular formula
4.5 Problems
5: Chemical Formulas and Nomenclature
5.1 General Background
5.1.1 Elements
5.1.2 Some basic definitions
5.2 Chemical Formula
5.3 Oxidation Numbers
5.3.1 Rules for assigning ONs
5.3.2 ONs in formulas
5.4 Writing the Formulas of Compounds
5.5 Nomenclature of Inorganic Compounds
5.6 Problems
6: Chemical Equations
6.1 Writing Chemical Equations
6.1.1 General rules for writing chemical equations
6.2 Balancing Chemical Equations
6.2.1 Guidelines for balancing a chemical equation
6.3 Types of Chemical Reactions
6.3.1 Combination or synthesis
6.3.2 Decomposition
6.3.3 Displacement
6.3.4 Double decomposition or metathesis
6.3.5 Neutralization
6.4 Problems
7: Stoichiometry
7.1 Reaction Stoichiometry
7.2 Information From a Balanced Equation
7.3 Types of Stoichiometric Problems
7.3.1 Solving stoichiometric problems
7.3.2 Mole-to-mole stoichiometric problems
7.3.3 Mass-to-mole stoichiometry problems
7.3.4 Mass-to-mass stoichiometry problems
7.3.5 Mass-to-volume stoichiometry problems
7.3.6 Volume-to-volume stoichiometry problems
7.4 Limiting Reagents
7.4.1 Limiting reagent calculations
7.5 Reaction Yields: Theoretical, Actual, and Percent Yields
7.6 Problems
8: Structure of the Atom
8.1 Electronic Structure of the Atom
8.2 Electromagnetic Radiation
8.3 The Nature of Matter and Quantum Theory
8.3.1 Photoelectric effect
8.4 The Hydrogen Atom
8.4.1 The Bohr model
8.4.2 Emission and absorption spectra
8.5 The Quantum-Mechanical Description of the Hydrogen Atom
8.5.1 The wave nature of the electron
8.5.2 The Heisenberg uncertainty principle
8.6 Quantum Mechanics and Atomic Orbitals
8.6.1 Orbitals and quantum numbers
8.6.1.1 Principal quantum number (n)
8.6.1.2 Angular-momentum (or azimuthal) quantum number (�)
8.6.1.3 Magnetic (or orbital) quantum number (ml)
8.6.1.4 Spin quantum number (ms)
8.7 Electronic Configuration of Multielectron Atoms
8.7.1 Aufbau principle
8.7.1.1 Shorthand notation for electron configuration
8.7.2 Pauli exclusion principle
8.7.3 Hund’s principle (or rule)
8.7.4 Summary of building-up principles
8.7.5 Orbital diagrams
8.8 Problems
9: Chemical Bonding 1: Basic Concepts
9.1 Introduction: Types of Chemical Bonds
9.1.1 Types of chemical bonds
9.2 Lewis Dot Symbols
9.2.1 Octet rule
9.2.2 Duet rule
9.3 Ionic Bonding: Formation of Ionic Compounds
9.3.1 Lattice energies and the strength of the ionic bond
9.3.2 Calculating lattice energies of ionic solids
9.3.3 Lewis structure of ionic compounds
9.4 Covalent Bonding: Lewis Structures for Molecules
9.5 Covalent Bonding: Writing Lewis Structures
9.5.1 Rules for writing Lewis structures
9.6 Resonance and Formal Charge
9.6.1 Resonance
9.6.2 Rules for writing resonance structures
9.6.3 Formal charge
9.6.4 Rules for assigning formal charge to an atom in a molecule
9.6.5 Using formal charge to determine molecular structure
9.7 Exceptions to the Octet Rule
9.8 Polar Covalent Bonds: Bond Polarity and Electronegativity
9.8.1 Determining polarity of a molecule
9.8.2 Dipole moment and percent ionic character
9.8.3 Dipole moment (μ)
9.8.4 Percent ionic character of a covalent polar bond
9.9 Problems
10: Chemical Bonding 2: Modern Theories of Chemical Bonding
10.1 VSPER Theory: Molecular Geometry and the Shapes of Molecules
10.1.1 VSEPR theory
10.1.2 Assumptions of the VSEPR theory
10.2 VSEPR Theory: Predicting Electron Group Geometry and Molecular Shape with the VSEPR Model
10.2.1 Electron group geometries
10.2.2 VSEPR theory: the AXE system
10.3 VSEPR Theory: Predicting Molecular Shape and Polarity
10.3.1 Steps for predicting molecular polarity
10.4 Valence Bond Theory
10.4.1 Postulates of valence bond theory
10.5 Valence Bond Theory: Types of Overlap
10.5.1 Sigma (σ) bond
10.5.2 Pi (π) bond
10.5.3 The strength of sigma and pi bonds
10.6 Hybridization
10.6.1 Characteristics of hybrid orbitals
10.6.2 Types of hybrid orbitals
10.6.2.1 sp3 hybridization—tetrahedral structure
10.6.2.2 sp2 hybridization—trigonal planar structure
10.6.2.3 sp hybridization
10.6.3 Hybridization and multiple bonds
10.6.3.1 Ethene, C2H4
10.6.3.2 Ethyne (acetylene), C2H2
10.6.4 Hybridization of elements involving d-orbitals
10.6.4.1 sp3d hybridization
10.6.4.2 sp3d2 hybridization
10.6.5 Predicting the hybrid orbitals used by an atom in bonding
10.7 Limitations of Valence Bond Theory
10.8 Molecular Orbital Theory
10.8.1 Linear combination of atomic orbitals (LCAO)
10.8.2 Molecular orbitals for simple diatomic molecules (H2 and He2)
10.8.3 Bonding in Li2 molecules
10.8.4 Molecular orbital energy-level diagram for homonuclear diatomic molecules
10.8.5 Molecular orbitals for heteronuclear diatomic molecules
10.8.6 MO electronic configuration and properties of the molecule
10.9 Problems
11: Gas Laws
11.1 Standard Temperature and Pressure
11.2 Boyle’s Law: Volume vs Pressure
11.3 Charles’s Law: Volume vs Temperature
11.4 The Combined Gas Law
11.5 Gay-Lussac’s Law and Reactions Involving Gases
11.6 Avogadro’s Law
11.7 The Ideal Gas Law
11.8 Density and Molecular Mass of a Gas
11.9 Molar Volume of an Ideal Gas
11.10 Dalton’s Law of Partial Pressure
11.11 Partial Pressure and Mole Fraction
11.12 Real Gases and Deviation from the Gas Laws
11.13 Graham’s Law of Diffusion
11.14 Problems
12: Liquids and Solids
12.1 The Liquid State
12.1.1 Properties of liquids
12.2 Vapor Pressure and the Clausius–Clapeyron Equatio
12.3 The Solid State
12.3.1 Types of solids
12.3.2 Crystal lattices
12.3.3 Unit cells
12.4 The Crystal System
12.4.1 Close-packed structure
12.4.2 Cubic unit cells
12.4.2.1 Guidelines for determining the number of atoms in a unit cell
12.4.2.2 Simple or primitive cubic unit cell
12.4.2.3 Body-centered cubic unit cell
12.4.2.4 Face-centered cubic unit cell
12.4.3 Coordination number
12.5 Calculations Involving Unit Cell Dimensions
12.6 Ionic Crystal Structure
12.6.1 The sodium chloride (NaCl), or “rock-salt” structur
12.6.2 The cesium chloride (CsCl) structure
12.6.3 The zinc blende (ZnS) structure
12.7 The Radius Ratio Rule for Ionic Compounds
12.8 Determination of Crystal Structure by X-Ray Diffraction
12.9 Problems
13: Solution Chemistry
13.1 Solution and Solubility
13.1.1 Some definitions
13.2 Concentration of Solutions
13.2.1 Percent by mass
13.2.2 Parts per million (ppm) and parts per billion (ppb)
13.2.3 Percent by volume
13.2.4 Molarity
13.2.5 Normality
13.2.6 Mole fraction
13.2.7 Molality
13.2.8 Dilute solutions
13.3 Solving Solubility Problems
13.3.1 Solubility in grams per 100 g of solvent
13.3.2 Solubility in moles per liter of solvent
13.4 Effect of Temperature on Solubility
13.5 Solubility Curves
13.6 Effect of Pressure on Solubility
13.7 Problems
14: Volumetric Analysis
14.1 Introduction
14.2 Applications of Titration
14.2.1 Acid-base titrations
14.2.2 A molar solution
14.2.3 Standard solutions
14.2.4 Standardization
14.3 Calculations Involving Acid-Base Titration
14.3.1 Calculation involving mass and percentage of substance titrated
14.3.2 Calculations involving molarity, mass concentration, solubility, and percentage purity from a
14.4 Back Titrations
14.5 Kjeldahl Nitrogen Determination
14.6 Problems
15: Ideal Solutions and Colligative Properties
15.1 Colligative Properties
15.2 Vapor Pressure and Raoult’s Law
15.2.1 Vapor pressure
15.2.2 Raoult’s law
15.2.3 Ideal solutions with two or more volatile components
15.3 Elevation of Boiling Point
15.4 Depression of Freezing Point
15.5 Osmosis and Osmotic Pressure
15.6 Problems
16: Chemical Kinetics
16.1 Rates of Reaction
16.2 Measurement of Reaction Rates
16.2.1 Instantaneous rate
16.3 Reaction Rates and Stoichiometry
16.4 Collision Theory of Reaction Rates
16.4.1 Factors affecting reaction rates
16.5 Rate Laws and the Order of Reactions
16.6 Experimental Determination of Rate Law Using Initial Rates
16.6.1 Alternate method
16.6.2 Determining the value of x
16.6.3 Determining the value of y
16.6.4 Determining the value of z
16.7 The Integrated Rate Equation
16.7.1 First-order reactions
16.7.2 Graphing first-order data
16.7.3 Second-order reactions
16.8 Half-Life of a Reaction
16.9 Reaction Rates and Temperature: The Arrhenius Equation
16.10 Problems
17: Chemical Equilibrium
17.1 Reversible and Irreversible Reactions
17.2 The Equilibrium Constant
17.2.1 Equilibrium constant in terms of pressure
17.2.2 Relationship between Kp and Kc
17.3 The Reaction Quotient
17.4 Predicting the Direction of Reaction
17.5 Position of Equilibrium
17.6 Homogeneous vs Heterogeneous Equilibria
17.7 Calculating Equilibrium Constants
17.7.1 Calculating Kc or Kp from known equilibrium a
17.7.2 Calculating K from initial concentration and one equilibrium concentration
17.8 Calculating Equilibrium Concentrations from K
17.8.1 A faster method for solving equilibrium problems
17.8.2 When to use the approximation method
17.9 Qualitative Treatment of Equilibrium: Le Chatelier’s Principle
17.9.1 Factors affecting a chemical reaction at eq
17.9.1.1 Changes in concentration
17.9.1.2 Changes in pressure
17.9.1.3 Changes in temperature
17.9.2 Addition of a catalyst
17.10 Problems
18: Ionic Equilibria and pH
18.1 The Ionization of Water
18.2 Definition of Acidity and Basicity
18.2.1 Ionic product of water
18.3 The pH of a Solution
18.4 The pOH of a Solution
18.5 The Acid Ionization Constant, Ka
18.5.1 Definition of pKa
18.6 Calculating pH and Equilibrium Concentrations in Solutions of Weak Acid
18.6.1 When to use the approximation method
18.7 Percent Dissociation of Weak Acids
18.8 The Base Dissociation Constant, Kb
18.9 Relationship Between Ka and Kb
18.10 Salt Hydrolysis: Acid–Basis Properties of Salts
18.10.1 The hydrolysis constant, Kh
18.10.2 Relationship between Kh and Kw
18.11 The Common Ion Effect
18.12 Buffers and pH of Buffer Solutions
18.12.1 How does a buffer work?
18.12.2 Buffer capacity and pH
18.12.3 The Henderson–Hasselbalch equation
18.13 Polyprotic Acids and Bases
18.14 More Acid–Base Titration
18.14.1 Acid–base indicators
18.14.2 Perception of color change of indicators
18.15 pH Titration Curves
18.15.1 Titration of strong acid against a strong base
18.15.2 Titration of weak acid against a strong base
18.16 Problems
18.16.1 pH, pOH, and percent ionization
18.16.2 Salt hydrolysis
18.16.3 Common ion effect
18.16.4 Buffers
18.16.5 Titration
19: Solubility and Complex-Ion Equilibria
19.1 Solubility Equilibria
19.2 The Solubility Product Principle
19.3 Determining Ksp from Molar Solubility
19.4 Calculating Molar Solubility from Ksp
19.5 Ksp and Precipitation
19.6 Complex-Ion Equilibria
19.6.1 Formation of complex ions
19.7 Problems
20: Thermochemistry
20.1 Introduction
20.2 Calorimetry and Heat Capacity
20.3 Enthalpy
20.3.1 Calculating �H of reaction
20.4 Hess’s Law of Heat Summation
20.4.1 Hints for using Hess’s law
20.5 Lattice Energy and the Born–Haber Cycle
20.6 Bond Energies and Enthalpy
20.7 Problems
21: Chemical Thermodynamics
21.1 Definition of Terms
21.2 The First Law of Thermodynamics
21.3 Expansion Work
21.4 Entropy
21.5 The Second Law of Thermodynamics
21.6 Calculation of Entropy Changes in Chemical Reactions
21.7 Free Energy
21.8 The Standard Free Energy Change
21.8.1 Calculating the standard free energy change
21.9 Enthalpy and Entropy Changes During a Phase Change
21.10 Free Energy and the Equilibrium Constant
21.11 Variation of �G0 and Equilibrium Constant with Temperature
21.11.1 Relationship between �G0 and K at different temperatures
21.12 Problems
22: Oxidation and Reduction Reactions
22.1 Introduction
22.2 Oxidation and Reduction in Terms of Electron Transfer
22.3 Oxidation Numbers (ON)
22.3.1 Rules for assigning oxidation numbers
22.3.2 Oxidation numbers in formulas
22.3.3 Oxidation number and nomenclature
22.4 Oxidation and Reduction in Terms of Oxidation Number
22.5 Disproportionation Reactions
22.6 Oxidizing and Reducing Agents
22.6.1 Identifying oxidizing and reducing agents
22.7 Half-Cell Reactions
22.8 Balancing Redox Equations
22.8.1 The oxidation-number method
22.8.2 The half-reaction method
22.9 Oxidation-Reduction Titration
22.9.1 Calculations involving redox titration
22.9.1.1 The conversion factor method
22.9.1.2 The formula method
22.10 Problems
23: Fundamentals of Electrochemistry
23.1 Galvanic Cells
23.2 The Cell Potential
23.3 Standard Electrode Potential
23.3.1 Standard reduction potential
23.4 The Electrochemical Series (ECS)
23.5 Applications of Electrode Potentia
23.6 Cell Diagrams
23.7 Calculating E0cell from Electrode Potential
23.8 Relationship of the Standard Electrode Potential, the Gibbs Free Energy, and the Equilibrium Co
23.8.1 Conditions for spontaneous change in redox reactions
23.9 Dependence of Cell Potential on Concentration (the Nernst Equation)
23.10 Electrolysis
23.11 Faraday’s Laws of Electrolysis
23.11.1 First law of electrolysis
23.11.2 The Faraday constant
23.11.3 Second law of electrolysis
23.12 Problems
24: Radioactivity and Nuclear Reactions
24.1 Definitions
24.2 Radioactive Decay and Nuclear Equations
24.2.1 Alpha emission 42α
24.2.2 Beta emission ( 0−1β)
24.2.3 Gamma radiation (γ )
24.2.4 Positron emission (01e)
24.2.5 Electron capture ( 0−1e)
24.3 Nuclear Transmutations
24.4 Rates of Radioactive Decay and Half-Life
24.4.1 Half-life
24.5 Energy of Nuclear Reactions
24.5.1 Mass defect
24.5.2 Binding energy
24.6 Problems

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