inorganic chemistry rapid learning series

P a g e | 1

RapidLearningCenter.com Rapid Learning Inc. All Rights Reserved

Inorganic Chemistry Rapid Learning Series

Course Study Guide

© All rights Reserved, Rapid Learning Inc.

P a g e | 2


COURSE TABLE OF CONTENTS

Core Unit #1 – The Inorganic Basics

Tutorial 01: Introduction to Inorganic Chemistry

Different branches of inorganic chemistry

Periodic classification of inorganic compounds

Groups

Periodic variations

Applications of inorganic chemistry:

Catalysis

Surface chemistry

Nanomaterials

Tutorial 02: General Chemistry Review and Periodic Table

The periodic table

Element information

The organization of the periodic table

How properties of an element can be determined from trends of the periodic table

How electronegativity, ionization energy, electron affinity relate to atomic radii

How ionic radii relates to atomic radii

Tutorial 03: Atomic Structures

Basic Structure of Atoms

Protons

Electron

Electron Cloud

Electron Configuration

Noble Gas Configuration

Quantum Numbers

Determining Quantum Numbers

Identifying Quantum Numbers

Tutorial 04: Simple Bonding Theory

Valence Bond Theory

The Octet Rule

Lewis Structures for:

Elements

Covalent Compounds

Polyatomic Ions

Ionic Compounds

P a g e | 3


Valence Shell Electron Pair Repulsion Theory

Electron and Molecular Geometry

Tutorial 05: Chemical Bonding and Molecular Orbitals

Types of Chemical Bonds

Ionic

Covalent

Bond Polarity

Isomers and Resonance

Hybridization

Hybrid Orbitals

sp3 Hybridization

Tutorial 06: Acids and Bases

History

Early theories

Arrhenius' theory of ionization

Arrhenius' definition of acids and bases

Acid-Base Concepts

Arrhenius Concept

Bronsted-Lowry Concept

Frontier Orbitals and Acid-Base Reactions

Hard and Soft Acids and Bases

Solvent Effects

Acid and Base Strength

Tutorial 07: Oxidation and Reduction

Basic Concepts and Definitions

Oxidization numbers

Redox equations

Electrochemical cells

Half cell reactions

Electrochemical Cells and Redox Reactions

Standard reduction potentials

Standard electrochemical cell potentials

Nernst equation

Redox Chemistry in Transition Metal Compounds

Electron counting: The 18 electron rule

Electron counting in the covalent model

Electron counting in the ionic model

Core Unit #2 – Quantum Chemistry

P a g e | 4


Tutorial 08: Principles of Quantum Mechanics

The Photoelectric Effect

Atomic Absorption and Emission of Light

The Bohr Hydrogen Atom

Wave-Particle Duality

Schrodinger Wave Equation

Energy Operator

Wavefunction

The Uncertainty Principle

Tutorial 9: Applications of Quantum Mechanics

Translational Motion

Particle in a Box

Tunnelling

Vibration

Energy Levels

Wavefunctions

Rotation

Two Dimensional

Three Dimensional

Perturbation Theory

Time Independent

Time Dependent

Tutorial 10: Molecular Symmetry and Group Theory

Symmetry Elements of Structures

Symmetry Operations and Elements

Symmetry Classification of Molecules

Character Tables

Matrix Representations

Irreducible Matrix Representations

MO Theory and Spectroscopy

Vanishing Integrals and Orbital Overlap

Vanishing Integrals and Selection Rules

Core Unit #3 – Comparative Chemistry

Tutorial 11: Intermolecular Forces and Introduction to Solids

Intermolecular Forces

London Dispersion Forces

Dipole-Dipole

Hydrogen Bonding

Solids

P a g e | 5


Amorphous Solid

Crystalline Solid

Types of Crystalline Structures

Molecular Solids

Ionic Solids

Phase Changes

Phase Diagrams

Energy of Phase Changes

Tutorial 12: Solid State Chemistry

Atomic Dimension and Molecular Size

The Electromagnetic Spectrum

Wavelength

Bragg’s Law

Bra’vais Lattices

Lattices - cubic, tetragonal

Unit Cell

Coordination Number (Primitive and Body Centered)

Close Packed Arrangements

Phase Diagrams, Density and Packing Arrangements

Bonding in Solids

Interactions in Ionic Solids

Molecular Orbitals

Tutorial 13: Main Group Trends

What are Main Group Elements

Common Group Names

Main Group Trends

Radii

Ionization Energy

Electron Affinity

Electronegativity

Boiling Point

Melting Point

The Physical Properties of Each Element

The Basic Chemistry of Each Element

Tutorial 14: Hydrogen, Groups 1 and 2 Elements

Hydrogen

Hydrogen

Hydrides

Group 1 Elements

Li to Cs

Minerals of Alkali Metals

P a g e | 6


Group 2 Elements

Be, Mg, Ca, Sr. Ba

General Properties

Grignard’s Reagent

Tutorial 15: Groups 13 – 18 Elements

Properties of group 13-18 elements

Properties and structures of the compounds of group 13-18 elements

Structure

Allotropes

Anions and Oxides

Applications of group 13-18 elements

Industrial Production

Effects on the Environment

Tutorial 16: Transition Elements – The d- and f-block Metals

d-block Elements

Group 3 Elements

Group 4 Elements

Group 5 Elements

Group 6 Elements

Group 7-10 Elements

Lanthanides and Actinides: f-block

Lanthanides Reactions and Properties

Actinides Applications

Core Unit #4 – Coordination Chemistry

Tutorial 17: Coordination Chemistry I - Nomenclature, Structures and Isomers


Coordination compounds and nomenclature

Isomers

The octet and 18 electron rules

Crystal Field Theory and Ligand Field Theory

Isomerization in Coordination Compounds

The Octet and 18 Electron Rules

The octet rule.

Exception to the octet rule.

Extension of the octet rule to compounds with d-orbitals: The 18 electron rule.

Crystal Field and Ligand Field Theories

Complexes with Common Coordination Numbers

VSEPR Theory

Tutorial 18: Coordination Chemistry II - Bonding

P a g e | 7


Formation of Coordination Compounds

Stability constants

The chelate effect

The trans effect on ligand substitution

Ligand Types and Bonding

Models of Electronic Structure

Valence bond theory

Crystal field theory

Ligand field theory

Angular Overlap Method

Ligand Positions

Angular Overlap Parameters

Spectrochemical Series

Tutorial 19: Coordination Chemistry III – Electronic Spectra

Absorption of Light

Complementary colors

Beer’s Law

Electron Interaction and Quantum Numbers

Quantum numbers of individual electrons

Russell-Saunders coupling

Quantum numbers of individual microstates

Electronic Spectra of Coordination Compounds

Finding the Ground State Term

Selection Rules

Correlation Diagrams

Examples and Applications

Tanabe-Sugano diagrams and spectra

Determining Δ0 from spectra

Charge transfer spectra

Tutorial 20: Coordination Chemistry IV – Reactions and Mechanisms

Reaction Coordinate Diagrams

Reactants and products

Transition states

Activation energy

Intermediates

Substitution Reactions of Octahedral Complexes

Inert and labile compounds

Dissociative mechanism

Associative mechanism

Substitution Reactions of Square Planar Compounds

Two term rate laws

Associative mechanisms

Trans effect

p-bonding effect

P a g e | 8


Other Reactions of Coordination Compounds

Redox reactions

Ligand reactions

Oxidative addition/reductive elimination

Insertion/elimination

Core Unit #5 – Organometallic Chemistry and Applications

Tutorial 21: Organometallics I – Introduction


Organometallic

Metal Alkyl

Metal Carbonyl

Metal Hydride

Metal Alkyls

Main group metal alkyls

Stability of metal alkyls

Metal alkyl decomposition pathways

Agostic, metallocycle, and bridging alkyls

Synthesis of metal alkyls

Metal Hydrides

Structure and Bonding

Synthesis

Metal Carbonyls

Metal CO Bonding

Bridging Carbonyl Ligands

Synthesis

Reactions

Ligands Related to CO

Metal Phosphine Complexes

Metal-Phosphine Bonding

Steric and Electronic Effects of Phosphine Ligands

Transition Metal Carbenes

Tutorial 22: Organometallics II – Synthesis and Catalysis


Ligand Exchange

Nucleophilic Displacement

Insertion and Elimination

Organometallic Catalysis

Catalytic Cycles

Buchwald-Hartwig Coupling

Heck reaction

Other Organometallic Reactions in Synthesis

Gilman Coupling

Hydroformylation

P a g e | 9


Heterogeneous Catalysts

Tutorial 23: Organometallics III – Main Group Comparison

Isolobal Analogy

Isolobal Fragments

Orbitals and Holes

The Ionic Model

Bonds between Metals

Cluster Compounds

Boron

Higher Order Boranes

Metal Carbonyl Clusters

Tutorial 24: Bioinorganic and Environmental Chemistry Basic Concepts and Definitions

Coordination compounds

Enzyme

Porphyrin

Chemotherapy

Transition Metal Complexes and Enzymes of Biological Importance

Iron

Cobalt

Copper

Coordination Compounds in Medicine

Platinum complexes

Metallocene compounds

Other transition metal complexes

Toxicity of Metal Complexes

Mercury

Lead

Other heavy metals

Chelation therapy

Environmental Inorganic Chemistry

Contamination by metals and metal compounds

Acid rain

Catalytic converters

P a g e | 10


Chapter by Chapter Detailed Content Descriptions

01: Introduction to Inorganic Chemistry

Chapter Summary

This chapter includes a review of the information contained in the periodic table of elements,

as well as introducing applications of inorganic compounds.

Tutorial Features

• Concept map showing inter-connections of concepts.

• Definition slides introduce terms as they are needed.

• Examples given throughout to illustrate how the concepts apply.

• A concise summary is given at the conclusion of the tutorial.

Key Concepts of Inorganic Chemistry

• Main Group Elements and p-block elements

• Transition Metals d-block Elements

• Organometallic Compounds

• Bioinorganic Chemistry

• Lanthanides and Actinides f-block Elements

Chapter Review

The explanation and prediction of the properties of elements and compounds using the

periodic table is the basis of inorganic chemistry.

The transition elements (groups 3-12) are named III B – VIIIB for groups 4-10 and IB, IIB

for groups 11 and 12 (according to the CAS system).

Groups 1-2 and 13-18 belong to the main group of elements. Group 1 and 2 belong to

the s-block in the periodic table (ns1, ns2 electron configuration).

The electron configuration represents the arrangement of electrons of an atom or

molecule. In the periodic table of elements, the first two rows of the transition metal

series in the have the general electron configuration, ns2(n-1)d.

A Coordination bond is formed between a vacant metal orbital (electron acceptor) and a

ligand (electron pair donor).

Organometallic compounds may contain transition or non-transition metals and contain a

bond between carbon and a metal. These compounds are made up of a combination of

inorganic and organic chemistry.

In chemical reactions, a catalyst accelerates the rate of a reaction. Two types of catalysts

are: (1) Homogeneous: Reactant and catalyst are dissolved in a suitable solvent. Easier

to measure and characterize solution species, (2) Heterogeneous: The catalyst is a solid

and in contact with a liquid or gaseous reaction medium.

Metal in Biology: Proteins and Enzymes: Transition metals act as sites for binding and

activation of oxygen. Metals in medicine: Pt anticancer drug, Gd contrast agents.

Environmental Chemistry: Toxicity of Hg and As compounds.

Electron transfer proteins: Iron sulfur proteins, and Cytochromes.

UV/visible spectroscopy: Most transition metal complexes are colored and absorb strongly

in the UV/visible region.

NMR: Multinuclear NMR spectroscopy (13C, 31P, 11B, 19F and 1H) spectroscopy can be

used to get valuable information about electronic and molecular structure.

P a g e | 11


Electron Spin (paramagnetic) Resonance Spectroscopy (ESR or EPR) can be used to get

information regarding the number of unpaired electrons in transition metal complexes.

X-Ray Crystallography: Complete Molecular structure is obtained, including bond lengths,

bond angles.

Chemisorption: Reactant molecule is adsorbed on surface of a substrate (catalyst) through

chemical bonds.

Physisorption: A molecule is bound to the surface via weak van der Walls forces.

P a g e | 12


02: General Chemistry Review and Periodic Table

Chapter Summary

This tutorial will introduce you to the periodic table and to the properties that can be

determined by an elements’ placement on the table, such as electronegativity, ionization

energy, electron affinity and atomic radii. This tutorial will show how to predict the relative

radii of an ion compared to the parent atom.

Tutorial Features





Key Concepts

Organization of the Periodic Table

Information for Each Element

Metals and Non-metals

Matrix Algebra

Atomic Mass

Atomic Radii

Ionization Energy

Electron Affinity

Predicting ion Charge

Chapter Review

Scientists went through many revisions to arrive at the current Periodic Table.

Periodic Table: A tool used by chemists. Organizes the elements and provides

information about them. Elements are organized by increasing order of atomic number.

Element Symbol: If there’s a second letter, it’s lower-case.

Atomic Number: Whole number elements are ordered by this on the periodic table.

Atomic Mass: Number with decimals.

Metals and non-metals have different properties. Metalloids have a blend of the

properties of metals and non-metals.

Mnemonic to memorize the first 20 elements: Happy Henry, The Little Beach Boy, CaN

dO FiNe; Naughty Megan, the Alpine Sister, Pretends to Ski at ClArK Canyon.

Periodicity of the Periodic Table: The predictable pattern by which properties of elements

change across or down the periodic table.

The periodic table is the most frequently used tool by chemists. The elements are

organized by increasing atomic number (which identifies the number of protons, defining

the identity of the atom). Every element has a different number of protons.

Atomic mass: Moving left to right, the number of protons, neutrons and electrons all

increase. More subatomic particles lead to higher mass.

Atomic Mass – the mass in grams for 6.02 x 1023 atoms. Found on the periodic table.

Atomic radius: The radius of an atom is determined by half the distance between two

nuclei when bonded together.

P a g e | 13


Ionization energy is the energy needed to pull off the farthest out electron. Ionization

energy increases across the periodic table and decreases down the table.

Electron affinity is the energy released when an electron is added to an atom. Electron

affinity increases from left to right across a period and decreases from top to bottom down

a group.

An ion forms from the loss or gain of electrons. The charge on some ions can be

determined by their position on the periodic table. When electrons are gained, there are

now more electrons than protons. Therefore, the protons have a weaker “pull” on each of

the electrons.

As a cation is formed, an electron is lost. There is now a greater number of protons than

electrons. Therefore, the protons have a greater pull on the remaining electrons and the

atom decreases in radius. Cations are always smaller than the parent atom.

Anions are formed from the addition of electrons. Anions have a higher number of

electrons than protons. The protons “pull” on each electron are therefore lower, and the

atom becomes larger. Anions are always larger than the neutral atom.

P a g e | 14


03: Atomic Structures

Chapter Summary

This tutorial will review basic atomic structure and showing how to determine the number of

electrons. Once the number of electrons is determined, they can be placed in energy levels,

subshells and orbitals to show electron configurations. Quantum numbers will also be

introduced.

Tutorial Features





Key Concepts

Atomic Structure

The Atom

Proton

Electron

Subshell

Electron Configuration

Spectroscopic Notation

Electron Configurations of Ions

Exceptions to the Aufbau Rule

Quantum Numbers

Determining Quantum Numbers

Chapter Review

Atom - Smallest piece of matter that has the chemical properties of the element.

Atoms contain 3 sub-atomic particles—protons, neutrons and electrons. Many people

confuse neutrons and electrons and think neutrons are negative because it starts with an

“n”—instead, think that “neutrons” are “neutral”

Electron cloud: It is the area outside of the nucleus where the electrons.

Susbhells are a set of orbitals with equal energy that make up an energy level. Orbitals

are the area where an electron is likely to be found.

There are 4 types of subshells, labeled s, p, d or f. The s orbitals are the lowest energy

and f orbitals are the highest energy. Only the s orbitals are low enough energy to be

found in level 1, while level 2 contains both s and p subshells. The third level has s, p,

and d orbitals and the f orbitals are added in the fourth energy level. Each subshell has a

different number of orbitals and each orbital can hold 2 electrons.

Electron Configurations – Shows the grouping and position of electrons in an atom.

Aufbau Principle: Electrons must fill subshells (and orbitals) so that the total energy of

atom is at a minimum.

Subshells are filled from the lowest energy level to increasing energy levels.

Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before

doubling up.

Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different

spins.

P a g e | 15


Spectroscopic notation is a shorthand method of showing electron configurations. Rather

than showing boxes for orbitals and arrows for electrons, the number of electrons is

written as a superscript after the orbital designation.

To write spectroscopic notation: Determine the number of electrons, follow the Aufbau

principle, remembering the number of electrons that each subshell can hold. The sum of

the superscripts gives the total number of electrons.

All groups of the periodic table have the same number of electrons in their highest

subshell. The periodic table is a guide to this. The “s-block” has 2 elements (the s

subshell can hold 2 electrons), the p-block has 6 elements (and the p subshell can hold 6

elements) the d-block has 10 elements (and d subshells can hold 10 electrons) and the f-

block has 14.

Ions: Lose or gain electrons to have a net charge and become more stable. Often, this

means that the outer shell, the valence shell, is full.

Noble gases all have a full valence shell. The noble gas notation uses the noble gas to

represent the inner, core, electrons and just show the configuration of the valence shell.

To write noble gas configuration, determine the number of electrons you need to place,

choose the noble gas closest to that number without going over. Start where that noble

gas left off on the periodic table and begin filling with spectroscopic notation.

The “d” subshells have 5 orbitals, and can hold 10 electrons. According to the Aufbau

principle, Cr would have a full 4s subshell and 4 electrons in the 3d subshell. However, a

half-full or completely full d orbital is more stable with a half-full s orbital is more stable

than a full s-orbital with 4 or 9 electrons in the d subshell.

Therefore, the actual configuration of Cr is 1 in the s and 5 in the d (making it half-full).

Quantum Numbers: A set of 4 numbers that describes the electron’s placement in the

atom.

The 4 quantum numbers are separated by a comma. The first number tells which

principal energy level the electron is in - which shell. The second number tells which

subshell type the electron is located in - s, p, d, or f. The third number tells which orbital

(which box) the electron is in, and the 4th number designates the “spin” of the electron.

Determining Quantum Numbers: The first number is the same as the shell. The second

number follows a coding system - 0 is always for “s”, “1” is always for “p” and so on. The

subshell number must be lower than the principal energy number (for example, there is

no “p” in level 1). The third number uses a number-line system to designate which orbital

an electron is in. The middle orbital is always 0, with negative numbered orbitals to the

left, and positive numbered orbitals to the right. The number line extends from –l to +l.

The last number is also a code system--+ ½ for an up arrow and – ½ for a down arrow.

P a g e | 16


04: Simple Bonding Theory

Chapter Summary

This tutorial will introduce you to valence bonding theory, the octet rule; show you how to

draw Lewis Structures and how to use the Valence Shell Electron Pair Repulsion Theory to

determine electron and molecular geometry.

Tutorial Features





Key Concepts

Valence Bond Theory

The Octet Rule

Lewis Structures for: Elements, Covalent Compounds, Polyatomic Ions, and Ionic

Compounds

Valence Shell Electron Pair Repulsion Theory

Electron and Molecular Geometry

Chapter Review:

Valence Shell: Outermost shell of electrons; the electrons with the highest principal

energy level number; the electrons that form chemical bonds. An Isolated Region Isolated

(System + Surroundings) does not exchange energy or material with the remainder of the

Universe.

Valence Bond Theory: Bonds are formed by overlap of valence orbitals.

Octet Rule: Most atoms are more stable with a full valence shell (which is a noble gas

configuration). A full shell has 8 electrons (“oct-” = 8).

The main groups of the periodic table have # of electrons = main group #.

Lewis structures are 2-dimensional representations of chemical bonds. Electrons are

represented by dots and the shared electrons show bonds. Other common terms for

Lewis Structures are: Electron Dot Structures, Dot Structures, and Lewis Dot Structures

Entropy change is defined in terms of a differential, infinitesimal, change in Entropy.

To draw a Lewis structure for an element, use the element’s symbol to represent the

nucleus and core electrons. Determine the number of valence electrons and draw them

around the symbol—placing one on each side before doubling up.

Drawing Binary Covalent Structures: For compounds with only 2 elements, arrange them

symmetrically. Determine the number of valence electrons for each atom, draw the

valence electrons (do not double up on a side where a bond will be formed). Any electron

written between two atoms can be counted by both atoms. Count the number of

electrons around each atom—when they each have 8 (except Hydrogen, which can only

hold 2), the structure is done.

A pair of electrons not being shared is a lone pair. A pair of electrons being shared

between two atoms is a bonding pair. All must be drawn and shown in the final structure.

Sometimes, in order to have all atoms with full valence shells, a hydrogen must be

bonded in a different location.

Polyatomic Ion: Group of atoms covalently bonded that together have a charge.

Ionic Compound – Metals transfers electrons to non-metals. The resulting ions form an

electrostatic attraction.

Common Exceptions to the Octet Rule: Hydrogen and helium can each only hold 2

electrons (because they only have an s orbital). Boron and beryllium can be stable with

only 6 valence electrons. And any element in the 3rd period or below can have more than

P a g e | 17


8 valence electrons as they have empty “d” orbitals (the d’s are not filled in electron

configurations until the next principle shell has been started). The empty d orbitals can

hold the extra electrons.

VSEPR: “Valence shell electron pair” refers to the pairs of electrons that form bonds or

lone pairs. The “repulsion” refers to the fact that electrons are all negative and like

charges repel each other. Since bonds and lone pairs are all negative, and negative

charges repel each other, bonds and lone pairs arrange themselves to be as far apart in

3D space as possible.

A generic chemical formula for molecular geometry uses “A” as the central atom, “X” as

the atoms bonded to the central atom (“ligands”) and “E” as lone pairs—one E is a pair of

electrons, not a single electron! And just like in chemical formulas, the absence of a

subscript indicates “1”.

Molecular geometry is determined by # of atoms bonded to the central atom.

Lone pairs actually take up slightly more space than a bonded pair as there is no other

nucleus to “control” it. If NH3 strictly followed the tetrahedron electron geometry, it

would have bond angles of 109.5˚. But in experiment, it is found that the bond angles

are 107.3˚. The lone pair on top “spreads out” and “squeezes” the bonded pairs slightly

together. The more lone pairs a molecule has, the greater the distortion.

P a g e | 18


05: Chemical Bonding and Molecular Orbitals

Chapter Summary

There are four types of chemical bonds. Each type of bond has different characteristics

related to what the electrons are doing. The first bond between two atoms is a sigma bond,

each additional bond is a pi bond.

Tutorial Features





Key Concepts

4 types of bonding and the properties associated with those bond types

Valence Bond Theory

Hybridization of orbitals

Sigma and Pi bonds

Molecular Orbital Theory

Bond Order

Paramagnetism and diamagnetism

Chapter Review

Ionic Bond: Formed from the electrostatic attraction between positive and negatively

charged ions due to transfer of electrons - contains both metals and non-metals.

A covalent bond is between non-metals. Rather than transferring electrons to form ions,

electrons are shared.

Polar covalent bonds are between non-metals and are the result of shared electrons as

well. However, the electrons are pulled more by one nucleus than the other and are

therefore shared unevenly.

Metallic Bond: Metal atoms join together and pool their electrons in a large network. Polar Bond – Uneven sharing of electrons due to differences in electronegativity. How is

bond polarity determined? Look up the electronegativity value of each atom and

determine the absolute value of the difference. If the difference is 0.5 – 1.4, the bond is

polar (higher than 1.4 indicates an ionic bond—electrons are pulled so much more by one

atom that they are transferred to that atom and ions are formed).

Isomers are two molecules with the same chemical formula, the same number of each

type of atom, but a different bonding structure.

Resonance: Average of two or more Lewis Structures which differ only in the position of

the electrons between atoms.

Sigma Bond (σ bond) – A bond in which one pair of electrons is shared. Formed from the

overlap of 2 s-orbitals. Pi Bond: 2nd or 3rd bond between two atoms. Formed from parallel p-orbitals.

Hybrid Orbitals: All of the orbitals in an atom involved in sigma bonds hybridize into equal

energy orbitals to form the bonds. Formed from mixing of orbitals.

sp3 hybridization: An atom with 2 sigma bonds would hybridize sp, one with 3 sigma

bonds would use sp2 hybridization and one with 4 sigma bonds would hybridize sp3.

P a g e | 19


Molecular orbital theory (MO Theory): states that instead of atomic orbitals

overlapping to form new bonds (as valence bond theory and hybridization theory states),

the molecule forms new orbitals as a molecule which are then occupied by the electrons.

When molecular orbitals are formed from atomic orbitals, bonding orbitals and

antibonding orbitals are created. Bonding orbitals pull the nuclei together, while

antibonding orbitals push them apart. Two p orbitals can create 2 sigma molecular orbitals—one bonding and one antibonding

2 parallel p orbitals can form 2 pi bond molecular orbitals—one bonding and one

antibonding.

A paramagnetic molecule is one that contains unpaired electrons. These electrons

without another electron of the opposite spin create an overall “spin” of the molecular,

which results in a magnetic property.

Diamagnetic is the opposite—all electrons are paired, resulting in no net spin. Therefore,

diamagnetic molecules have no magnetic properties.

P a g e | 20


06: Acids and Bases

Chapter Summary

This tutorial will cover the history of acid-base theory and the major acid base concepts. It

will also cover the hard and soft theory of acids and bases, solvent effects, and acid-base

strength. Some compounds behave as acids or bases. These can be strong or weak, and

hard or soft. These properties affect solubility and acid-base reactions. Acids may react to

leave behind a conjugate base, or a base may react to form its conjugate acid.

Tutorial Features





Key Concepts

Stabilities of Phases

Arrhenius' theory of ionization

Bronsted's definition of acids and bases

Acid-Base concepts

Hard and soft acids and bases

Solvent effects

Acid and base strength

Chapter Review:

Early Acid-Base Theories: Acids have a sour taste that was originally believed to be

caused by the presence of oxygen. Hydrogen is actually the element in common to many

acids, but the acid concept can be extended beyond hydrogen containing compounds.

In 1810 Humphry Davy proposed that hydrogen is the element common to all acids.

However, many hydrogen containing compounds are not acids. In 1887, Arrhenius

published a paper on ionization, in which, he defined an acid as a proton (H+) donor and a

base as a hydroxide ion (OH-) donor. He failed to consider the role of the solvent in acid-

base chemistry.

Bronsted: “Acids and bases are substances that are capable of splitting off or taking up

hydrogen ions, respectively”. This definition allowed the base concept to be applied to

bases without hydroxide ions, such as ammonia and amines.

Arrhenius considered an acid to be a proton donor and a base as a hydroxide ion donor.

This described the behavior of many common acids and bases.

Like Arrhenius, Bronsted and Lowry considered an acid to be a proton donor. A base is a

proton acceptor. It does not necessarily have to be a hydroxide ion. When an acid transfers a proton, it leaves behind its conjugate base. And when a base

accepts a proton, it forms its conjugate acid.

H3O+ represents a hydrated proton, which is bound to more than one water

molecule. The proton is constantly hopping between groups of water molecules

and exchanging with the protons from water.

Protic solvents can autoionize.

Examples include water, ammonia, sulfuric acid, acetic acid, and alcohols.

Frontier orbitals are at the frontier between occupied and vacant orbitals.

HOMO: Highest occupied molecular orbital. Usually a bonding or nonbonding orbital.

LUMO: Lowest unoccupied molecular orbital. Usually non-bonding or antibonding. Orbitals closely matched in energy form strong bonds. An example is a hydrogen s-orbital

interacting with a carbon sp3 orbital.

P a g e | 21


Degenerate orbitals have the same energy. In the free atom, all 3 p-orbitals are

degenerate. Also in the free atom, all 5 d-orbitals are degenerate. Molecules have lower

symmetry than free atoms. Degeneracy is broken in molecules.

Group theory provides a mathematical treatment of molecular symmetry. Orbitals are

labeled with symbols that reflect their symmetry and degeneracy. The symbol A or B is

given to a singly degenerate orbital, according to its symmetry. Doubly degenerate

orbitals are labeled E, and triply degenerate orbitals are labeled T. Additional superscripts

or subscripts may be used to provide additional symmetry information.

A base (B) may form a hydrogen bonded complex to the acid (HA). This complex may be

stable or an intermediate in proton transfer reactions.

pKa values are a quantitative measure of strength of Bronsted acids and bases. Reactions

of Lewis acids depend more on the polarizability of the cation and anion.

Hard acids are generally small ions, ions with a high positive charge, or complexes with

electron withdrawing groups.

Soft acids have large cations or low positive charge, or contain electron releasing groups. Information about acid-base interactions can be obtained from physical property

measurements. Calorimetry can measure the enthalpy of acid-base reactions.

Thermodynamic: Thermodynamic properties can be measured directly or indirectly by

Hess's Law. For the dissociation of the weak acid HA, Ka can be determined indirectly.

Superacids are defined as acids that are stronger than sulfuric acid. Because of the

leveling effect of water, the Hammett acidity function is used to measure superacid

strength. Fuming sulfuric acid, which contains an excess of SO3, forms several acids that

are stronger than sulfuric acid.

P a g e | 22


07: Oxidation and Reduction

Chapter Summary

Chemical reactions sometimes involve electron transfer. Oxidization is the loss of electrons

and reduction is the gain of electrons. Reduction-oxidization “Redox” reactions can be

described by electrochemical cells. Cell potentials are related to reaction free energy and

equilibrium constants.

Tutorial Features





Key Concepts

Oxidization and Reduction.

Assigning oxidization numbers

Redox equations.

Electrochemical cells

Relationships between cell potential, free energy, and equilibrium constants.

Applications to transition metal complexes.

Chapter Review

Oxidization means the loss of electrons. It originally meant combustion, which consists

of oxidization reactions.

Reduction means gain of electrons. It is the opposite of oxidization.

Oxidization numbers help to keep track of electrons in chemical reactions. They resemble

ionic charges but they do not generally reflect actual atomic charges. By definition, the oxidization number of an element in its uncombined form is zero. The

oxidization number of a monoatomic ion is the charge on the ion. In a neutral molecule,

the sum of all oxidization numbers is zero. In a polyatomic ion, the sum of all oxidization

numbers is equal to the charge on the ion.

A redox reaction can be visualized as taking place in an electrochemical cell. The actual

reaction may take place in a cell, but it doesn’t necessarily have to.

An electrochemical cell consists of compartments with electrodes, where the oxidization

and reduction reactions occur. Oxidization occurs at the anode and reduction occurs at

the cathode. A salt bridge provides a path for electrical current to flow from one half cell

to the other. An electrical load or voltmeter completes the circuit.

Although oxidization and reduction reactions always occur together, it is convenient to

think of the half reactions separately.

Several methods of balancing redox reactions exist. The most fool proof is the half

reaction method. The half reactions are each balanced separately, and then the half

reactions are added together so that there is no net gain or loss of electrons. Slightly

different procedures are used for reactions in acidic or basic solution.

Multiply the balanced half reactions by a coefficient that makes the electrons cancel on

both sides. That is, the number of electrons lost in the oxidization must equal the number

of electrons gained in the reduction. Add the half reactions together and cancel anything

that appears on both sides of the equation.

In organic reactions, electron density decreases in an oxidization and increases in a

reduction. Reduction breaks a bond to an electronegative atom, or forms a C-H bond.

P a g e | 23


Cell potentials measure the potential differences between two electrodes. This

corresponds to the reaction potential. We think in terms of half-cell reactions but there is

no way to measure a half cell potential. We can define one half-cell reaction as having a

potential of zero, and measure all other half cell reactions against it.

Using tabulated standard reduction potentials, the voltage (EMF) of any cell can be

calculated under standard state conditions. E0cell = E0

cathode - E0anode, where the E0 of the

cathode and anode are the standard reduction potentials of the half reactions.

The Nernst equation allows us to calculate cell potentials when the concentrations are not

all 1 M. Recall that the reaction quotient Q is the ratio of product concentrations to

reactant concentrations.

Electron Count in Transition Metal Complexes: Electron counting can be done using either

the covalent or ionic model. Both models will give the same answer. The concepts of redox reactions and oxidization numbers have applications in transition

metal chemistry. Counting electrons is more complex but simple models allow accurate

electron counting.

P a g e | 24


08: Principles of Quantum Mechanics

Chapter Summary

Classical mechanics is unable to explain blackbody radiation or the photoelectric effect.

Quantum Mechanics is able to explain these and other important concerns such as the

structures of atoms and molecules.

Tutorial Features





Key Concepts

Origins of Quantum Mechanics

Dynamics of Microscopic Systems

Quantum Mechanical Principles

Postulates of Quantum Mechanics

Chapter Review

Photoelectric Effect – When white light is shone on a metal substance, electrons may be

emitted.

Light behaves with wave properties. Light is electromagnetic energy, waves consisting of

both electric and magnetic characteristics. Light has wave properties such as wavelength and frequency.

The Photoelectric Effect gives us information that light behaves in ways not explained by

wave theory.

Light is composed of particles of energy. Particles of light are photons which travel at the

speed of light. Light is quantized (it comes in discreet packets of energy).

Bohr Hydrogen Atom: The Bohr hydrogen atom is an improvement on Classical Physics

because it correctly identifies quantized energy states for hydrogen. The Bohr hydrogen

atom gives correct values for the energy levels for hydrogen and all one-electron atoms.

The Bohr hydrogen atom incorrectly places electrons in well-defined orbits around a

nucleus.

Wave-Particle Duality – De Broglie suggested light has both wave and particle properties.

When the wave-particle duality is applied to electrons it is found to be true, and is the

basis for the electron microscope.

Classical physics and quantum mechanics describe the dynamic behavior of a particle in

different ways. Classical physics Describes particle in terms of its location and

momentum. Quantum mechanics Uses a wavefunction ψ to describe the quantum state

and behavior of a particle.

The Schrodinger Equation is a wave equation and Ψ is a wavefunction. Ĥ is a

mathematical instruction and is called the Hamiltonian or Energy Operator. E is the

Energy of a particle described by Ψ and is an Eigenvalue. The value of E is obtained by

“solving” the Schrodinger Equation, Ĥ Ψ = E Ψ. It is the time-independent form of the

Schrodinger wave equation, Eigenvalue Equation.

Requirements of the Wavefunction: The Wavefunction must be Single-Valued. It must be

continuous, and have a continuous first derivative. Also, it must be square integrable for

each point in space.

The wavefunction for a system contains all dynamic information

The Uncertainty Principle places limits on the precision of our knowledge of certain

dynamical variables. A mathematical statement of this principle is ΔpΔq ≥ ½ħ. This means

that the product of the uncertainty in Momentum Δp and the uncertainty in position Δq at

a point in time must be equal to or greater than: ½ħ.

P a g e | 25


The Wavefunction Ψ for a system contains all dynamical information. |Ψ2|dτ is the

probability of finding the associated particle in the region dτ. A wavefunction Ψ must be

continuous, have a continuous first derivative, be single valued and be square integrable.

Dynamical variables are represented by operators formed from the position and

momentum operators.

There is a minimum amount of uncertainty in the simultaneous measurement of position

and momentum.

P a g e | 26


09: Applications of Quantum Mechanics

Chapter Summary

This tutorial covers Translational Motion: Particle in a Box, Motion in Multiple Directions, and

Tunneling. Rotational motion is quantized due to the wave nature of matter. A rotating

particle’s “wave” must “fit” evenly on the path of a rotational state in order to not have wave

interference and to be single valued.

Tutorial Features





Key Concepts

Particle in a Box

Motion in Multiple Directions

Tunneling

Rotational motion

Chapter Review

A particle confined between two infinitely high barriers and constrained to move in only

the x direction it is called a Particle in a One-Dimensional Box.

A particle moving in a one-dimensional box has Quantized Translational Motion.

Quantum Mechanics allows a particle in a box to have only certain amounts of energy, its

energy is Quantized.

Tunneling describes the entry of a particle into a region of space in which the potential

energy is greater than the total energy of the particle.

Using appropriate boundary conditions, solving the Schrodinger equation gives Normalized

Wave Functions and associated Eigenvalues.

Particle in a box wavefunctions and hydrogen atomic orbital wavefunctions are

Orthonormal sets of wavefunctions.

Separation of Variables is a technique used to solve for multidimensional Wavefunctions

and is used to obtain Wavefunctions and Eigenvalues for the three-dimensional Atomic

Orbitals and for a particle in a multi-dimensional box.

An object that experiences a restoring force F that is proportion to displacement x has

Harmonic Motion, and is a Harmonic Oscillator.

A chemical bond is approximated as a Quantum Harmonic Oscillator.

Rotational motion is quantized due to the wave nature of matter.

The angular momentum of a particle rotating in an xy-plane is limited to the values.

A rotating particle’s wave must fit evenly on the path of a rotational state in order to not

have wave interference and to be single valued.

Time-Independent Perturbation theory represents a real Hamiltonian Operator by a

sum of a Model System Hamiltonian Operator and a Perturbation Operator.

In Time-Independent Perturbation Theory, the perturbation is constant and does not

change with time.

In Time-Dedependent Perturbation Theory, the perturbation is not constant and does

change with time.

Wavefunctions for a model system is used to obtain approximate Wavefunctions for a

real system.

P a g e | 27


First-order-corrections to model wavefunctions are linear combinations of model system

wavefunctions with coefficients calculated using a first-order-Perturbation Hamiltonian

term and model wavefunctions and Eigenvalues.

Coefficients in the linear combination of model wavefunctions are calculated using a

first-order-Perturbation Hamiltonian term and model wavefunctions and Eigenvalues.

A first-order correction to the ground state energy level is calculated using a zero-order

model ground state wavefunction and the first-order-Perturbation Hamiltonian Operator.

The second-order correction to the ground state energy level is calculated using a

complete set of Model wavefunctions and the first-order-Perturbation Hamiltonian

Operator.

An important example of a Time-Dependent Perturbation is encountered when

electromagnetic radiation interacts with atoms and molecules.

P a g e | 28


10: Molecular Symmetry and Group Theory

Chapter Summary

Orbital overlap is required for bond formation and is determined using Group Theory.

Orbitals must have the same symmetry species in order to have a nonzero overlap integral.

The best Basis atomic orbitals for a LCAO-MO may be determined in part by evaluating the

symmetry species of candidate orbitals. Symmetry-adapted Linear Combinations (SALC)

Basis Atomic Orbitals are symmetry-selected linear combinations of atomic orbitals. Selection

Rules for spectroscopic transitions are determined by the Transition Dipole Moment. The

criterion for an allowed transition is that the integrand in the Transition-Dipole-Moment

Integral must have symmetry species A1 and thus be nonzero.

Tutorial Features





Key Concepts

Symmetry Operations and Elements

Symmetry Classification of Molecules

Character Tables and Symmetry Labels

Vanishing Integrals and Orbital Overlap

Vanishing Integrals and Selection Rules

Chapter Review

Molecular shapes are classified by Symmetry Operations and associated Symmetry

Elements.

A benzene molecule may be rotated in steps of 600 around the center of the molecule and

look the same after each step.

The shapes of molecules are described in terms of the set of corresponding symmetry

elements.

Unique sets of symmetry elements are called Point Groups.

A Character Table relates the symmetry operations associated with a symmetry group to

the symmetry species for the group.

Optimum LCAO-MO Bases are composed of atomic orbitals having appropriate

symmetry, which may be determined using Character Table data and Group Theory.

Orbital overlap is required for bond formation and may be determined using Group

Theory.

Orbital overlap is determined in LCAO-MO calculations by evaluating overlap integrals.

Orbitals Symmetry-adapted Linear Combinations (SALC) Bases are symmetry-

selected linear combinations of atomic orbitals tailored to specific molecules.

Selection Rules for spectroscopic transitions are determined by the Transition Dipole

Moment.

P a g e | 29


11: Intermolecular Forces and Introduction to Solids

Chapter Summary

Intermolecular forces are chemical bonds that exist within a molecule. Thus tutorial will

cover, properties of liquids, including vapor pressure, solid structure types, phase changes

and the energy of phase changes.

Tutorial Features





Key Concepts

What intermolecular forces are.

Properties of liquids.

How vapor pressure is affected by intermolecular forces and temperature.

How solids are structured.

How matter changes states.

The energy changes during a phase change.

Chapter Review

Intramolecular forces are within the molecule (the chemical bond). Intermolecular

forces are the physical attractions between two separate molecules.

All molecules have London Dispersion Forces, which are due to a temporary ganging up

of electrons as they move around the nuclei. A temporary dipole is the result. This results

in a portion of the molecule that has a partially negative charge and a portion with a

partial positive charge.

London Dispersion forces are temporary, as soon as the electrons move again and “spread

out” the force is over. It is the weakest force. All molecules have London Dispersion

Forces, as all molecules have electrons—but those with more electrons have larger London

Dispersion Forces.

Polar molecules have permanent dipoles that can be attracted to one another forming

Dipole-Dipole forces. Ions can also be attracted to the permanent dipole in a polar

molecule, forming ion-dipole forces.

Dipoles in polar molecules are permanent, and therefore dipole forces are stronger than

London Dispersion Forces. Only polar molecules exhibit dipole forces, those with stronger

dipoles having stronger dipole forces.

Hydrogen bonding is a very strong dipole. Since hydrogen atoms do not have any other

electrons to balance the positive nucleus, if the electrons are pulled away by a very

electronegative element, the hydrogen is left with a strong partial positive charge.

Hydrogen bonds are an especially strong version of a dipole-dipole interaction. They are

the strongest of the intermolecular forces, although not as strong as a chemical bond.

Only the very electronegative atoms (N, O and F) exhibit hydrogen bonding with

hydrogen.

Some general properties of liquids: volume cannot change, but shape can; molecules can

slide past one another; they are not noticeably compressible.

Vapor Pressure is the pressure created when liquid or solid particles on the surface have

enough energy to escape the intermolecular forces and become a gas.

The rate of evaporation doesn’t change over time, but as gas particles are created they

can begin to re-join the liquid. Therefore, the rate of condensation increases over time.

When the two rates are equal, equilibrium is established.

Solids cannot change volume or shape, the particles are not free to change positions, and

they are not compressible.

P a g e | 30


Amorphous Solid: Has a fair amount of disorder in the structure.

Crystalline Solid: Has a highly regular, repeating arrangement of particles.

Crystalline structures are lattice structures composed of repeated unit cells. The three

simplest types of unit cells are cubic, body-centered and face-centered cubic structures.

Crystalline solids can be categorized as atomic solids, molecular solids or ionic solids.

Atomic solids are either metallic or network solids.

Metallic bonding contains metal atoms packed closely together and bonded to the atom in

each direction equally.

Phase Change – Changes between solids, liquids and gases.

Intermolecular forces are broken or formed in phase changes—not chemical bonds. It’s a

physical change, not a chemical one.

Boiling/Condensation Point: When liquid and gas phases are at equilibrium with one

another.

A substance will boil when the vapor pressure of the liquid is high enough to overcome the

atmospheric pressure above the sample. This occurs at the boiling point.

Melting/Freezing Point: When solid and liquid are at equilibrium.

Some substances go straight from a solid to a liquid—subliming. This occurs when the

intermolecular forces in the solid are very weak and can be completely broken to form gas

particles.

The enthalpy of fusion is the energy needed to break enough IMF’s to form a liquid. When

forming a solid, the energy is released— the opposite of heat of fusion.

Enthalpy of vaporization (Hvap): energy necessary to break the rest of the IMF’s and

turns a liquid into a gas.

P a g e | 31


12: Solid State Chemistry

Chapter Summary

This tutorial will review the interaction of matter and the electromagnetic spectrum based on

the size of atoms. Crystallography will be introduced to demonstrate the Bra’vais Lattices.

How atoms pack in ionic and covalent solids, the prediction of possible packing arrangements

and common crystalline forms will also be introduced.

Tutorial Features





Key Concepts

The Periodic Trends of Atomic and Ionic Radii

The Relationship between the Electromagnetic Spectrum and Matter

How X-Rays are used to Determine the Structure of Solids

Bra’vais Lattices

How Atoms and Ions Pack

How to Predict Matter Density and Potential Structure from Atomic Size

Common Crystalline Forms and bonding in solids

Chapter Review:

Atoms are approximately 10-10 m in diameter, because of this small size, atoms cannot be

visualized even with a very powerful optical microscope.

Atomic radii increase down a column on the periodic table and decreases across a row.

Wavelengths: In order for a wave to interact with an object, the object must be larger

than the wave.

By using a wavelength of electromagnetic energy smaller that atomic distances one can

discover how the atoms are arranged.

If one has a single crystal the opportunity exists to identify every atoms position in a

molecule.

Using a random sample one cannot see single atom positions.

However, by using Bragg’s Law we can determine the distances between layers of atoms.

Solids are formed from atoms, molecules, or ions as building blocks arranged in repeating

patterns called units cells.

A Unit Cell is an array of crystal-lattice points with parallel sides that generates a Crystal

Lattice when moved along the x, y, and z axes as if it were a building block.

A crystal lattice is represented by arranging unit cells, like building blocks, in a three-

dimensional array of Lattice Points.

A Lattice Point is a point in space that may be occupied by a structural unit of a crystal

such as a molecule or atom.

14 different unit-cell patterns, Bravais Lattices: including body centered, face centered,

side centered are classified into 7 different crystal systems: including cubic, monoclinic,

and triclinic.

There are three cubic Lattices: simple cubic, face centered cubic and body centered cubic.

Face centered cubic has a structural unit at the center of each face of the unit cell.

Body centered cubic has a structural unit at the center of the unit cell.

There are two monoclinic Lattices: simple, and side-centered.

Side-centered monoclinic has structural units at the centers of two parallel sides.

There is only one triclinic Lattice.

P a g e | 32


Planes containing three or more non-collinear lattice points may be drawn through crystal

lattices in many ways and the locations/spacing of these planes are important features of

a crystalline solid.

Lattice Planes lead to X-Ray Diffraction Patterns when X-Rays pass through crystals.

X-Ray Diffraction Patterns are caused by wave interference as X-Rays are scattered by

structures in Lattice Planes.

Binary structures are made up of two elements, and the two criteria which determine

material packing are: (1) ratio of cations to anion radius, and (2) numbers of each.

Bonding in Molecular Solids Metals: When the basic bonding interaction is weak the

separation between bonding and anti-bonding states (bands) is small enough that

electrons can move from one band to the other based on thermal energy alone.

An insulator is formed when the basic bonding interaction is very strong. The band gap is

large enough, to prevent electrons moving from the valence band to the conduction band.

Most metals transfer electrons through their solid structure based on the “band gap”. As

temperature increases resistance increases. So if we cool the metal down can we get the

resistance to zero. The answer is yes, at the transition temperature. The metal becomes

a superconductor. As the temperature is decrease a critical temperature is reach and

the material does several unique things.

Zero Resistance: The flow of electrons is unrestricted. A current flowing through the

material moves without resistance and as long as the temperature remains below the

critical temperature no additional current needs to be applied. Almost perpetual motion.

The invisible force of magnetism is repelled by the material. This is known as the Meissner

Effect. This is due to the unique property, of a superconductor being cooled below its

critical temperature. In the picture a magnet is suspended over the superconductor.

Basic Theory of Superconductors: There are two basic parts to the theory of how these

materials work. (1) The first is called Cooper Pairing. An electron that is moving though

the lattice of a superconductor generates, relative to other electrons a positive charge and

attracts another electron. They push and pull each other through the lattice, (2) The

second is the Lattice Vibration: At a certain temperature the lattice vibrations of the

superconductor matches the electrons motion and assist in their movement instead of

“resist” it. Consider how this relates to our discussion of packing structures and the

possibility to regulate materials with different layers and distances between those layers.

P a g e | 33


13: Main Group Trends

Chapter Summary

This tutorial will review periodic trends and the specifics of the main group elements. This

will include: radii, 1st ionization energies, electron affinity, electronegativity, melting point

and boiling point as well as an overview of what main group elements are will be discussed.

An overview of the chemistry of each group of elements will then be covered and how to

understand the trends in a more comprehensive way.

Tutorial Features





Key Concepts

Main Group Elements

The Periodic Trends of the Main Group Elements

The Physical Properties of Each Element

The Basic Chemistry of Each Element

Chapter Review

Main group elements are sometimes called representative elements. They are also

called s and p block elements. These are the elements in group 1, 2 and 13 through 18.

The synthetic use of these industrial materials is vastly important. However, as basic

chemicals they do not determine the fundamental processes, mechanisms and catalytic

cycles so important to industry to make the vast number of downstream products. Many

organic compounds, mixed systems and transition metals do.

Atomic radii increase down a column on the periodic table and decreases across a row.

Ionization energy of an element is the energy required to remove the highest energy

electron from a neutral atom in the gas phase. The trend is that ionization energy

increases from left to right and decreases from top to bottom.

Electron affinity is the energy change when a neutral atom in the gas phase gains an extra

electron. Electron affinity increases from left to right and decreases from top to bottom

Electronegativity is an estimate of the tendency of an atom to attract a bonding pair of

electrons. Electronegativity increases from right to left and decreases from top to bottom.

The known possible number of oxidation states increases as most other trends. Down

the table the lower oxidation state or states predominate.

Hydrogen is usually placed in Group 1. It loses electron to form H+. Aqueous acid / base

chemistry and its chemistry unrelated to Alkali Metals. Group 14: Outer electron shell half

full, similar electronic properties, electronegativity, and forms covalent bonds. Group 17:

gains electron to form H-. For this reason hydrogen is usually discussed on its own and

separate from any group.

Group 1 Trends: The trends show radius increases as the mass increases. Ionization

Energy, Electron Affinity, Electronegativity, boiling point and melting point all decrease

down the family. As Francium is a short lived radioactive species no values for its boiling

point and melting point exist.

Group 2 Trends: Radii increase with increasing mass. Ionization energy decreases.

Electronegativity decreases. Boiling and melting points in general decrease although the

trend in not as regular as for group 1.

Group 2 Trends: Group 13 elements do not show nearly the magnitude of atomic trends as

seen in group 1 and 2. The effect of the transition metals and inner transition elements

diminishes what should be expected. Boiling and melting points decrease down the group.

P a g e | 34


Boron exists as a covalent network solid where the rest of the family are metallic. This

accounts for the large difference in melting points.

Group 14 Trends: The atomic radii increases down the group. The electronic properties do

not follow a nice decrease as expected due to the influence of transition and inner

transition influence.

The group 15 elements trends: Electronegativity seems to follow the normal trend by

decreasing down the group. Boiling point however increases, while melting point goes up

and then down.

In group 16, the radii increases as the mass increases, and the ionization energy

decreases down the group. The boiling and melting points increase to the metallic

element polonium (slight drop).

For group 17, radii increases down the group. Ionization Energy decreases. Electron

Affinity decreases with a small jump from fluorine to chlorine. Electronegativity decreases.

Boiling Point increases as does the melting point as you would expect the large molecules

will have larger intermolecular, London dispersion forces.

For group 18, Radii increases smoothly down the group. Ionization energy decreases.

Boiling points and melting points are also predictable based on the increasing size of the

atoms.

Not only is the chemistry of elements within a group similar so are diagonal elements.

P a g e | 35


14: Hydrogen, Groups 1 and 2 Elements

Chapter Summary

The nomenclature for the Akali metals are group lA according to the CAS system and group 1

according to the IUPAC system. The alkaline earth metals are group 2 (IUPAC) and group llA

in the CAS system. This tutorial covers the following: General properties of main group 1-2

elements, unique properties of hydrogen and its compounds, and reactivity and compounds

of group 1 and 2 elements.

Tutorial Features





Key Concepts

Main Group 1-2 Elements

Properties of Group 1-2 Elements

Unique Properties of Hydrogen

Reactivity and compounds of group 1 and 2 elements

Chapter Review

Groups 1-2 and 13-18 belong to the main group of elements. Group 1 and 2 belong to

the s-block in the periodic table (ns1, ns2 electron configuration).

Hydrogen has an electronic structure similar to group 1 elements. However, hydrogen has

different properties. Hydrogen exists as a diatomic gas and is not a metal.

Formation of H+ due to the loss of an electron: H+ (proton) exists as H3O+ in water. It

shares an electron pair with non-metals to form a covalent bond (e.g., C-H, N-H).

Gain of an electron from a highly electropositive metal to form a hydride (H-): e.g., NaH,

KH. Hydrogen has three naturally occurring isotopes: 1 H, 2H (deuterium) and 3H

(tritium).

Hydrides formed between alkali and alkaline earth metals are ionic in character. LiH, CaH2,

NaH are examples of such hydrides.

LiAlH4 is an extremely useful reducing agent in synthetic organic and inorganic

applications. It can be to reduce ketones, aldehydes, nitriles and nitro compounds.

Hydrogen can be used as a fuel instead of hydrocarbons. The burning of H2 with O2 results

in more energy per unit mass than gasoline. In fuel cells, the energy released is converted

to electricity.

Naturally occurring hydrogen is 99.985 % 1H (proton). Deuterium (2H) is also a stable

nuclei and deuterated compounds can be used to study reaction mechanisms involving H

atom transfers.

Group 1 elements are highly electropositive; they react with water-liberating hydrogen

gas. Reactivity increases from Li to Cs. All form +1 charged ions.

Physical and chemical properties correlate with size of metal atom/ion: Melting point,

lattice energy of salts and hydration energies decrease from Li to Cs.

Sodium (Na) exists in the minerals, feldspar, sodalite and rock salt. Sodium is produced

commercially by the electrolysis of molten NaCl.

Potassium is important for the function of living cells and is the seventh most abundant

element in the earth’s crust. In some countries, the principle source of K is potash (ashes

of plants).

Alkali metals form complexes with cyclic ligands containing several donor atoms. Two

major classes of such ligands are cryptands and crown ethers.

Crown ethers are cyclic compounds containing several ether groups.

P a g e | 36


Cryptands are 3-dimensional analogues of crown ethers. Cryptands offer better selectivity

and binding strength for alkali metals than crown ethers.The Boltzmann Distribution is

used to calculate populations of the energy states.

Alkalides are chemical compounds in which alkali metals are anions. Cryptands are used

to generate the alkalide salts.

Alkaline earth metals have a higher melting and boiling point compared with the alkali

metals. Mg and Ca are essential elements for living organisms. Be forms an oxide

coating, which prevents reaction with water.

Mg and Ca are one of the most abundant in the earth’s crust. Sr and Ba are not as

abundant. All isotopes of radium are radioactive.

Beryllium shows different properties than the other alkaline earth metals. Be forms mainly

covalent bonds. Be can form electron deficient bonds (lacks octet configuration) similar to

group 13 element boron and aluminum.

Beryllium chloride (or halides) can exist as linear monomeric or dimeric structures (in the

vapor phase). In crystals, the BeCl2 can polymerize to form chloride (halogen) bridged

chains.

Magnesium salts are isolated from the minerals, dolamite and magnesite. Mg is highly

reactive, hence, is not found naturally in an elemental form.

Mg-Al alloys are used in automotive and truck components. Magnesium metal is obtained

by electrolysis of Mg salts obtained from brine.

Grignard’s reagent functions by acting as a nucleophile and reacting with the electrophilic

carbon atom of the polar carbonyl (C=O) bond. Reactions involving Grignard’s reagent

must be conducted under inert conditions, since air and water can cause protonation or

oxidation of the reagent.

P a g e | 37


15: Groups 13 – 18 Elements

Chapter Summary

This tutorial will cover the properties of group 13-18 elements, the properties and structures

of the compounds of group 13-18 elements, as well as the applications of group 13-18

elements.

Tutorial Features





Key Concepts

Group 13-18 elements

Halogens: Interhalogens

Nitrogen Compounds include Oxides and Hydrides.

Chapter Review

In the old IUPAC system, the letters A and B were used to designate elements to the left

and right of the periodic table. In the CAS system, commonly used in the United States,

the letters A and B refer to main group and transition elements, respectively. In the new

IUPAC system, the numbers are increased from 1-18 from left to right side of the periodic

table.

In group 13, the metallic properties are less than that of the alkali and alkaline earth

metals. Compounds formed with non-metals are not completely ionic. Boron is a metalloid

and exhibits unique chemical properties relative to other members of the group. Members

of this group have an ns2np1 electron configuration.

Chemistry of boron is different from other elements in group 13. Boron forms compounds,

which lack the octet electron configuration. Boron forms hydrides like C and forms

oxygen-containing minerals like Si.

Aluminum also forms compounds containing bridging alkyl and bridging hydrogens.

Aluminum forms 3 and 4 coordinate structures.

Group 14: The properties of the elements in this group show considerable variation. C and

Si are non-metals whereas Sn and Pb are metals.

Allotropes of Carbon: Until recently, C was known to exist in only two allotropes

(diamond and graphite). In recent years, the discovery of fullerenes (C60 and related

structures) has generated interest in carbon chemistry and potential applications.

Diamond structure consists of C atoms tetrahedrally bonded to four other atoms,

resulting in a very hard structure. Graphite contains fused 6-membered rings in a layered

structure.

Carbides are formed when carbon combines with very electropositive metals. Metal

carbide bonds have ionic and covalent bond characteristics (the more electropositive the

metal, the higher the ionic character).

Carbon dioxide is a product of combustion reactions. CO2 absorbs thermal energy (bond

vibrational energy) and is believed to be responsible for the greenhouse effect (global

warming).

P a g e | 38


Minerals of silicon are based exclusively on SiO4 tetrahedral structural units. Silicon is

less reactive than carbon, however, is more reactive than Ge. Silicon is a metalloid and a

semiconductor.

Silanes are analogous to hydrocarbons. Silanes are more reactive than hydrocarbons due

to a weaker Si-Si and Si-O bonds.

Elements in group 15 are known to form double and triple bonds in compounds. As, Sb

and Bi are known to be toxic since they form radicals, which are not easily processed by

the liver.

Nitrogen gas is relatively inert due to the triple bond. Liquid nitrogen at 77K is used to

cool superconductors and to study low temperature reactions. Nitrogen is trivalent in most

compounds. Nitrogen bonds with certain metals to form coordination complexes.

Nitrides (N3-) do not exist in aqueous solution, since it is easily protonated. Nitrides of s,

p and d block elements are known. Li and alkali metals form ionic nitrides whereas other

nitrides show a varying degree of covalency.

Nitrous oxide is used as an anesthetic for dental purposes and is also used as a propellant

for aerosols. N2O is a better substitute for chlorofluorocarbons since it causes less damage

to the environment. However, N2O is a greenhouse gas due to absorption of thermal

energy. N2O4 dissociates into NO2 and forms an equilibrium mixture containing significant

amounts of both species.

Group 16: All elements have the ns2np4 electron configuration. Oxygen and sulfur are

non-metals; Se, Te and Po are metalloids.

Oxygen exists as a diatomic molecule. The ground state of oxygen has two unpaired

electrons in the highest antibonding (* 2p) molecular orbital. Oxygen is paramagnetic

due to the presence of the unpaired electrons.

All halogens are highly reactive due to their high electronegativity. Halogens exist as

diatomic molecules in nature. Chlorine is a gas, bromine is a liquid, and iodine is a solid

under ambient conditions. Fluorine is an extremely toxic and reactive gas.

The noble gases have ns2np6 electron configuration. The fully-filled configuration is

responsible for the inertness of the gases. Ne, Ar, Kr and Xe are obtained by liquefaction

of air and fractional distillation. He (helium) is separated from natural gas.

P a g e | 39


16: Transition Elements – The d- and f-block Metals

Chapter Summary:

The learning objectives for this chapter are the following: (1) Properties and compounds of

d-block (transition metals, 3d, 4d and 5d series of elements), (2) Properties, applications and

reactions of lanthanides (4f series), and (3) Properties and applications of actinides (5f

series).

Tutorial Features





Key Concepts

Properties and compounds of d-block (transition metals, 3d, 4d and 5d series of

elements).

Properties, applications and reactions of lanthanides (4f series).

Properties and applications of actinides (5f series).

Chapter Review

d-Block Electron Configuration: The first two rows of the transition metal series in the

periodic table have the general electron configuration, ns2(n-1)d. The third row has the

configuration, ns2, (n-2)f14, (n-1)d.

Oxidation Number: The number of valence electrons includes the number of electrons in

the highest n and n-1 d-electrons. The most common oxidation numbers for elements left

of Mn correspond to the sum of ns and (n-1)d electrons for that element (Sc: +3, Ti:+4,

V: +5, Cr: +6 and Mn: +7).

Sc and y are the only definite group 3 elements. Lanthanum and lutetium are generally

classified with the lanthanides and called the rare earth metals. Sc occurs as trace metals

in many minerals (gadolinite, euxenite). Yitrium was found in lunar rock samples in high

amounts. Lutetium is found in the phosphate mineral monazite. Sc, Y and Lu are found

in low abundance in the earth’s crust and are hard to extract.

Hf and Zr show identical chemical properties. All isotopes of Rf are radioactive. Ti shows

high strength, yet is much lighter than steel. This property has resulted in applications in

aerospace industry as components of jet engines. Ti is also used in medical devices, such

as in prosthetic devices and orthopedic implants; it is also used in cell phones.

Group 5 Elements: General electron configuration: ns2 (n-1)d3, except Nb (4d35s1). V,

Nb forms several oxides (I, II, III, IV and V oxidation states). Ta forms only Ta(V) oxide.

Group 6 Elements: Cr and Mo show d5 electron configuration. Group 6 elements form

compounds with various oxidation states (-2 to +6). The stability of the +6 oxidation state

increases down the group.

Group 7 Elements: General electron configuration: ns2 (n-1)d5. Tc and Re are found only

in trace amounts in nature. Tc has no stable isotopes. 55-Mn is the most stable isotope of

Mn.

Group 8 Elements: Tc and Re are found only in trace amounts in nature. Tc has no

stable isotopes. 55-Mn is the most stable isotope of Mn.

Iron: Biological Role: Oxygen transport and activation by heme enzymes and proteins.

Iron is the major component by mass of the earth’s inner and outer core. The common

(stable) oxidation states for iron are +2 (ferrous) and +3 (ferric).

P a g e | 40


Group 9 Elements: Ru, Rh, Pd, Os, Ir and Pt belong to the platinum group metals. They

exhibit the same chemical and physical properties and occur in the same mineral deposits.

Rhodium is found as a free metal in Pt and Ni ores. Rhodium is used in catalytic

converters of automobiles. Wilkinson’s catalyst (Rh complex) is used for the

hydrogenation of alkenes. Vaska’s complex (Ir complex) is known to bind with molecular

oxygen.

Group 10 Elements: Palladium and Pt are used in automobile catalytic converters, which

convert harmful gases (hydrocarbons, nitrogen oxides and carbon monoxide) to less toxic

gases. Pd is also used as a catalyst in several organic reactions, such as hydrogenation

and dehydrogenation reactions. Pd absorbs hydrogen gas at room temperature and is

investigated as a potential hydrogen storage device and in fuel cell applications.

The lanthanide series is comprised of 15 elements between atomic number 57

(lanthanum) and 71 (letutiem). Both lanthanum and letutiem have a single 5d electron,

however, have similar properties as other elements in the lanthanide series.

Applications of lanthanides include: (a) Superconducting magnets. Magnetic resonance

imaging, (b) Battery and electronic components, (c) Luminescent materials used in the

optoelectronic industry, and (d) lasers.

All lanthanides exhibit +3 oxidation state. The trivalent compounds are mostly ionic.

Trivalent lanthanides are known to form complexes with polydentate nitrogen ligands. The

most common divalent ion is Eu(II). All lanthanides are known to form divalent

compounds (stable f7 configuration).

The actinide series includes elements between atomic number 89 (actinium) and 103

(lawrencium). Actinides have similar properties as lanthanides. All actinides are

radioactive and release energy upon radioactive decay. Actinides show more variation in

the oxidation state than lanthanides.

Plutonium-235 is also capable of spontaneous nuclear chain reactions. Plutonium is a

product of U -235 fission and is mainly used for nuclear weapons. The neutrons released

upon fission of 235-U propagates a continuous controlled chain reaction, in reactors.

P a g e | 41


17: Coordination Chemistry I - Nomenclature, Structures and Isomers

Chapter Summary

This chapter will introduce the basic concepts in coordination chemistry. Nomenclature and

isomers will be described. The octet and 18 electron rules will be discussed, and the crystal

field and ligand field theories will be introduced. This chapter will conclude with a brief survey

of compounds with common coordination numbers.

Tutorial Features





Key Concepts

Historical background

Nomenclature of coordination compounds

Constitutional isomers

Stereoisomers

The octet rule

The 18 electron rule

Crystal field and ligand field theory

Complexes with common coordination numbers

Chapter Review

A Coordination Compound or Coordination Complex is a compound of a metal and

ligands. A ligand can be thought of as an atom, ion, or molecule that donates both

electrons of the bond to the metal. However, the bonding may be more complex than

this simple model suggests.

Some common ligands include water, ammonia, CO, amines, and phosphines. Other

neutral ligands include alkenes and aromatics. Ionic ligands include hydroxide, halide,

and cyanide ions. Organic ligands may be either neutral or anionic species.

Coordination compounds have been known since ancient times but their structures

remained unknown. An entrenched idea was the concept that a metal has a fixed

valence. CoCl2 and CoCl3 were known, but not CoCl4, for example.

Metal complexes were known. Alfred Werner proposed that cobalt could be bonded to

6 ligands.

Simple coordination compounds are named with the ligands in alphabetical order,

followed by the metal. Prefixes di-, tri-, etc. indicate the number of ligands of the

same type. These prefixes do not count towards alphabetizing.

Complexes that consist of ions are written as the cation first, then the anion. Inner

sphere ligands, those connected directly to the metal, are placed with the metal in

square brackets.

The more common convention is to indicate the metal oxidization number by a roman

numeral. Alternatively, the charge of the metal-containing fragment can be indicated

in parentheses after the name.

Isomers are two different compounds with the same molecular formula. The word

isomer comes from the Greek words meaning “same parts”.

P a g e | 42


Constitutional isomers have the atoms connected differently.

Stereoisomers have the same connectivity but the atoms are oriented differently in

space.

Ionization isomers are compounds with the same molecular formula which generate

different ions in solution.

Achiral compounds may become chiral chelated ligands due to their conformation. A

procedure similar to the one used to find ring handedness has been developed.

Consider one chelate ring of ethylene diamine on cobalt. This ring can exist in either

of two enantiomeric conformations.

The octet rule, and its extension to the 18 electron rule, is based on the stability of

filled orbitals. Most second and third row compounds use the s and p valence orbitals.

These can hold a total of 8 electrons, 2 from the s, and 2 from each of the p orbitals.

Many compounds of boron and aluminum only have 6 valence electrons. These can

achieve an octet by dimerization or acting as a Lewis acid. Examples are BH3, BF3,

and AlCl3. Compounds of P, S, Xe, and other heavy nonmetals can accommodate more

than 8 electrons by using low lying d-orbitals. Examples include PF5 and XeF4. Note

that the noble gases are not always completely inert.

Crystal field theory gives an approximate and qualitative picture of the d-orbital

energies in a coordination compound. In an octahedral complex, the dx2-y2 and dz2

orbitals are raised in energy. These are given the symmetry label “eg” by group

theory.

Ligand Field Theory reaches the same conclusions via an approximate molecular

orbital treatment. Ligand lone pair orbitals may be of the proper energy and symmetry

to interact with the metal s, p, and d-orbitals.

The 18 electron rule can be understood from Ligand Field Theory. 12 electrons fill the

6 bonding orbitals with mostly ligand character.

Compounds with a coordination number (CN) of one are rare. These usually involve a

ligand that is so sterically hindered that no other ligands can approach the metal. The

CN of 2 is also rare. Ag(NH3)2+ is one example.

Common geometries for CN5 are the trigonal bipyramid and square pyramid. The

energy difference between these two structures is small.

The octahedral geometry is the most common in transition metal coordination

compounds. Many complexes adopt distorted octahedral geometries with one

dimension lengthened or compressed.

P a g e | 43


18: Coordination Chemistry II - Bonding

Chapter Summary

There are a number of factors that affect bonding between a metal and its ligands. The

formation of coordination compounds includes: stability constants, the stability of

chelates, and the trans effect on ligand substitution.

Tutorial Features





Key Concepts

Stability constants

Factors that affect stability

Ligand types and bonding to the metal

Unpaired electrons and magnetic measurements

Valence bond theory

Crystal field and ligand field theory

Applications of the bonding models

Chapter Review

Stability constants are a special case of equilibrium constants. In general, one or

more new ligands substitute for existing water ligands.

Stability constants, also known as formation constants, are equilibrium constants for

coordination compound formation reactions.

A chelate is a complex in which a ligand is bonded via more than one atom.

Dissociation of a chelating ligand results in a smaller positive change in entropy,

compared to a non-chelating ligand.

Some ligands facilitate the substitution of other ligands trans to it. This is particularly

important in the synthesis of square planar complexes.

Ligands bonding to the metal via their lone pairs are normally σ-bonded. The ligand

lone pair orbital overlaps with a metal d-orbital.

In addition to s-bonding, some ligand can form a p bond from a filled d-orbital into a

ligand p-antibonding orbital.

Backbonding occurs when an occupied metal d-orbital overlaps with a ligand pi-

antibonding orbital. This strengthens the metal-ligand bond, but weakens the bond

between the ligand atoms (C and O in this case) as an antibonding orbital becomes

populated.

A compound with all electrons paired is diamagnetic. Diamagnetic compounds are

slightly repelled by a magnetic field. A compound with one or more unpaired electrons

is paramagnetic. Paramagnetic compounds are attracted by a magnetic field.

Magnetic susceptibility measurements use the fact that a paramagnetic substance will

weigh more in a magnetic field. Some instruments weigh the sample with and without

the magnetic field.

Magnetic susceptibility is related to the magnetic moment µ. The value of m can be

calculated from the total orbital angular momentum L and the total spin angular

momentum S.

Pauling’s valence bond theory introduced hybrid orbitals that make it easier to

visualize molecular orbitals two atoms at a time. sp3, sp2, and sp orbitals are familiar

P a g e | 44


from organic chemistry. Trigonal bipyramidal shapes can be represented by sp3d

hybrid orbitals.

Crystal field theory gives a qualitative picture of orbital energies when applied to

transition metal complexes.

Ligand Field Stabilization Energy (LFSE) is defined by the energy of a given

electron configuration, relative to the energy of the same complex with all 5

degenerate d-orbitals equally populated.

Bonding in a square planar complex uses 3 sets of ligand p orbitals: the py for sigma

bonding, the px for pi bonding in the molecular plane, and the pz for pi bonding

perpendicular to the molecular plane.

One magnetic field shift is Chemical Shift.

Crystal field and ligand field theories provide a qualitative picture of orbital splitting.

The angular overlap method provides a semi-quantitative measure of orbital energies.

Ligand positions are numbered for octahedral, tetrahedral, and other complexes. Each

ligand position is assigned an empirical parameter for its interaction with each d

orbital. The positions are shown for the octahedral and tetrahedral geometries.

For pi-donor ligands, the interacting metal orbitals are raised in energy and become

antibonding. The pi atomic overlap energy parameters are the same in magnitude but

opposite in sign.

If two (or more) degenerate orbitals are unequally occupied, the geometry will distort

to break the degeneracy. A similar distortion occurs in antiaromatic molecules where 2

degenerate orbitals are singly occupied.

Recall the antiaromatic cyclobutadiene molecule. The geometry distorts away from the

antiaromatic structure with equal bond lengths. This Jahn-Teller distortion can occur

in coordination compounds.

Applications of the bonding models: spectrochemical series, magnitude of Δ, and the

Jahn-Teller effect.

P a g e | 45


19: Coordination Chemistry III – Electronic Spectra

Chapter Summary

This Chapter will cover the absorption of light and color of compounds, the interaction of

electrons and resulting quantum numbers of atomic states, electronic spectra and selection

rules, correlation diagrams, and examples and applications.

Tutorial Features





Key Concepts

Colors of compounds

Absorption of light

Interaction of electrons and quantum numbers of atomic states

Electronic spectra of coordination compounds

Selection rules

Correlation diagrams

Chapter Review

When light passes through matter, some colors are absorbed. The matter appears as the

complementary color to that of the absorbed light.

Complementary colors can be determined from a color wheel. A substance that

absorbs red light will appear green. If it absorbs yellow light it will appear blue.

Many compounds absorb at more than one visible wavelength. This makes it difficult to

predict exactly what color the compound will be.

Beer’s Law, also known at the Beer-Lambert law, relates the absorption of light to the

concentration and path length. log(I0/I) = A = εlc, I = intensity of light entering the

sample and I0 = intensity leaving the sample. A = absorbance; l = path length; c =

concentration; ε = molar absorptivity.

Individual electron quantum numbers are obtained from solving the Schrodinger

equation. The orbital type is obtained from quantum number l: l =0, s orbital; l =1, p

orbital; l = 2, d orbital, etc.

Consider the electron configuration s1p1. For both electrons, n = 2 for a 2s and 2p orbital.

For the s electron, l = 0. For the p electron, l = 1. For the s electron, ml = 0 and ms =

±1/2.

The quantum numbers n, l, ml, and ms define the states of individual electrons in an

atom. The atomic quantum numbers ML and MS define individual microstates of an

electron configuration.

Term symbols are used to indicate atomic states. They consist of a letter and a

superscript to the upper left. Sometimes a subscript to the lower right is also used. The

letter indicates the total angular momentum L. Note that the symbol “S” is used in 2

different ways: To indicate the total spin angular momentum, and to indicate the state

corresponding to L = 0.

Hund’s rules predict which atomic state is the ground state: (1) The ground state has the

highest spin multiplicity. In the p2 example, that is the 3P state, (2) If the multiplicity is

tied, the ground state is the one with the larger L value. A 1D state will be lower in energy

than a 1S state, (3) For subshells (orbital sets) that are less than half filled, the ground

state has the lowest J value. For more than half filled subshells, the state with the highest

J value is the ground state.

P a g e | 46


Selection rules determine which electronic transitions are most probable. These rules are

not entirely rigid, and many colors of compounds arise largely from “forbidden”

transitions.

There are 2 primary selection rules. The Laporte rule forbids d-d transitions and other

transitions between states with the same parity. These transitions are often weak but still

observed. The spin selection rule requires transitions to be between states with the same

multiplicity. If not, the transitions are improbable and the spectral bands are weak. Spin

selection rule: Transitions are only allowed between states with the same multiplicity. 4A2 to 4T1 is allowed. 4F to 3H is forbidden.

Up to now we have considered free ion states in an atom or ion with spherical symmetry.

When placed into an octahedral ligand field, the degeneracy of the free ion terms breaks.

This splits the terms into several new terms that reflect the octahedral symmetry of the

complex. Term splitting is similar to the orbital splitting into the eg and t2g sets. We first

consider a system with a single d electron.

Another form of correlation diagram was developed by Y. Tanabe and S. Sugano. The y

axis is the energy of each state in units of E/B. The x axis is the ligand field splitting

parameter, in units of Do/B. B is an energy parameter that represents repulsion between

terms of the same multiplicity. It is called the Racah parameter. Spin allowed transitions

are easily visualized in the Tanabe-Sugano diagrams.

The Tanabe-Sugano diagram predicts one spectral band for each spin allowed

transition. It does not guarantee that each band will be visible. Bands may be outside the

spectral region, or obscured by charge transfer bands.

Charge transfer spectra occur when electrons are excited from a ligand (M-L bonding)

orbital to a vacant metal d orbital. This is called Charge Transfer To Metal (CTTM). These

transitions are allowed and usually result in very strong absorptions. They result in a

formal reduction of the metal.

Electrons in filled metal d orbitals can also be excited to ligand pi* orbitals. This is called

Charge Transfer To Ligand (CTTL). These occur when the ligand has an empty pi* orbital,

such as CO and aromatic ligands. It results in a formal oxidization of the metal.

P a g e | 47


20: Coordination Chemistry IV – Reactions and Mechanisms

Chapter Summary

This chapter will cover reaction coordinate diagrams and kinetics, ligand substitution

reactions and their stereochemistry, oxidative addition/reductive elimination reactions, and

insertion/elimination reactions.

Tutorial Features





Key Concepts

Reaction coordinate diagrams and kinetics

Ligand substitution reactions

Substitution reaction stereochemistry

Redox reactions

Oxidative addition and reductive elimination reactions

Insertion and elimination reactions

Chapter Review

A single step reaction goes from reactants to products via a single transition state. The

transition state is the maximum on the potential energy curve.

An example of a two-step reaction is the SN2.

The activation energy is the energy difference between the reactant and transition state.

For an n step reaction, there are n transition states and n-1 intermediates. An

intermediate may or may not be observable by spectroscopy.

Graphing different powers of reaction rate on the y-axis versus concentration of a reactant

on the x-axis, the order with respect to that reactant may be determined.

Among the most common reactions of coordination compounds are those in which one

ligand is displaced by another. If the compound reacts fast (kinetics), it is labile. If it

reacts slowly, the compound is inert. The inertness refers to a slow reaction, not a

thermodynamically inert species.

Reaction rates depend on the temperature of reaction in addition to the initial

concentrations of reactants.

Ligand Field Stabilization Energies of the reactant and intermediate can be used to obtain

a rough estimate of the activation energy. Ligand Field Activation Energy = Difference in

LFSE’s.

The rate of a dissociative reaction can be influenced by electrostatic effects that make the

ligands less apt to dissociate, or by steric effects, which enhance the rate of ligand

departure. The incoming ligand enters after the rate determining step, so it has little or

no effect on the reaction rate.

In the associative mechanism, an additional ligand Y adds to the coordination compound,

followed by the departure of ligand X.

The associative mechanism involves a pentacoordinate complex as an intermediate.

Several of these trigonal bipyramidal complexes have been isolated. The degree of

incoming ligand bond formation or departing ligand bond cleavage can vary. The terms Ia

and Id are used to describe these variants, similar to octahedral complexes.

The interchange mechanism lies between the dissociative and associative mechanisms.

P a g e | 48


Some ligand substitution reactions are base catalyzed. A coordinated ammonia, amine, or

water ligand loses a proton. The ligand trans to the amido or hydroxo group is usually

displaced.

The trans effect is caused by two factors. The first is weakening of the M-l sigma bond

trans to the ligand T. A strong trans effect ligand will hog the metal d orbital, making it

less available for bonding to X.

The other factor causing the trans effect is: Stabilization of the pentacoordinate

intermediate. This is a pi bonding effect. Strong pi-acceptor ligands tend to be strong

trans effect ligands.

Insertion occurs when two ligands are bonded to the same metal. One ligand inserts into

the other. The reverse process is elimination. The reaction is usually favored in one

direction.

β-Elimination is the most common elimination reaction. It is a common decomposition

pathway for metal alkyls with a β-hydrogen.

Insertion and elimination are the reverse of each other, although thermodynamics usually

favors one process over the other. In a 1, 1 insertion, the metal and ligand X end up on

the same carbon of the L type ligand. This is typical for insertions involving CO. In a 1, 2

insertion, the metal and ligand X end up on adjacent carbons. This is typical for insertions

involving alkenes.

For compounds without a β-hydrogen, -elimination is sometimes possible. The product of

-elimination is a metal carbene, with a C=M bond.

P a g e | 49


21: Organometallics I – Introduction

Chapter Summary

Organometallics are a subset of coordination compounds in which one or more ligands has a

carbon directly bonded to the metal. There are several different ligand types and several

kinds of bonding. The chemistry is also affected by other important ligands, particularly

hydride and phophine ligands. This tutorial will cover the chemistry of metal alkyls,

carbonyls, and carbenes.

Tutorial Features





Key Concepts

Historical Background

Organometallics as coordination compounds

Metal alkyls

Metal hydrides

Metal carbonyls and related compounds

Metal phosphine complexes

Transition metal carbenoids

Chapter Review

In 1900 Victor Grignard discovered a reaction between alkyl halides and magnesium

metal. Magnesium inserts into the alkyl-halogen bond by an oxidative addition

reaction. This was the first example of a metal alkyl.

The first metal carbonyl, Ni(CO)4, was discovered in 1884 by Ludwig Mond. He was

investigating damage to nickel valves by CO. He intentionally heated Ni metal in a

stream of CO, and discovered the reversible formation of nickel carbonyl.

As with other coordination compounds, electrons can be counted with the covalent

model or the ionic model. In the covalent model, the metal electron count is the same

as in the free metal atom. Each ligand is considered as a neutral fragment.

In the ionic model, the metal electron count is the same as in the metal ion

corresponding to its oxidation state.

Most organometallic compounds have 18 electrons. Common exceptions are: Radical

species with an odd number of electrons. These are often oxidized or reduced to an

18 electron species. 17 Electron compounds may also dimerize. 16 electrons are

common with Ni, Pd, Pt, and some other metals.

Organometallics are coordination compounds with one or more organic ligands. The

organic ligand may be an alkyl group, CO, or a pi-bonded ligand. Reactions can occur

at the metal center, or at the organic fragment.

Main group organometallics: Alkyllithiums, Grignard reagents, and related

compounds. Metal alkyls and metal hydrides: A metal atom is directly bonded to an

alkyl group or hydride.

Alkyl magnesium halides, the Grignard reagent, actually exists as an equilibrium

mixture of R2Mg, MgCl2, and RMgCl, or the analogous bromide or iodide species. This

is called the Schlenk equilibrium.

Trialkylaluminums have the dual properties of alkyl group donor and Lewis acid, and

they are commonly used as Lewis acids in organic and polymer synthesis. Except for

P a g e | 50


alkyllithiums, alkali metal alkyls are difficult to prepare, extremely reactive and

pyrophoric, and are seldom used.

Transition metal alkyls were usually difficult to synthesize. It is now understood that

the difficulty was due to the easy decomposition of those compounds. When the

decomposition pathways are blocked, transition metal alkyls can be synthesized and

isolated.

Cyclic metal dialkyls are called metallocycles. They often exist in equilibrium with

metal-coordinated alkenes. Metallocyclopropanes are resonance structures of metal

alkene complexes.

Some metal alkyls are prepared by reaction of the metal with an electrophile. Common

electrophiles are alkyl halides, acyl halides, and iodonium salts. This method is

common for anionic metal complexes. The mechanism may be an SN2 or a radical

mechanism.

Main group metal hydrides are often simple ionic compounds. Examples: NaH, CaH2

Borane tends to dimerize in the gas phase using 3 center, 2 electron bonds. It also

forms complexes with Lewis bases.

Metal hydrides can be prepared from protonation of a basic metal complex. The

reverse reaction is a common reaction of metal hydrides if the conjugate base is

stabilized by ligands. CO ligands are particularly good at stabilizing a negative charge. The carbon lone pair of CO can form a σ bond with an empty metal d-orbital. Metal to

ligand back bonding is also possible. IR spectroscopy provides clues about the amount

of back bonding.

Back bonding requires occupied metal d-orbitals. The metal induced polarization

activates CO to attack by nucleophiles at carbon, or electrophiles at oxygen.

Metal carbonyls can be prepared directly from the metal and CO. They can also be

prepared as a metal complex. This may involve ligand exchange and/or addition of a

CO ligand.

Phosphines bond to the metal via the phosphorous lone pair. Back bonding can also occur between filled metal d-orbitals and the P-R σ* orbital.

Transition metal carbenes have a formal metal-carbon double bond. They can be

thought of as arising from a free carbene interacting with a metal.

Triple bonded metal-carbon complexes are called carbynes. Common reactions are

coupling with another carbyne, with alkyne formation.

P a g e | 51


22: Organometallics II – Synthesis and Catalysis

Chapter Summary

Major uses of organometallic compounds are in synthesis and catalysis. Some useful

products are made by reaction of a stoichiometric amount of an organometallic compound.

Other organometallics are used in catalytic cycles. This chapter will cover applications of

transition metal organometallic compounds in synthesis and catalysts.

Tutorial Features





Key Concepts

Organometallic reactions

Catalytic cycles as a sequence of simple reaction steps

Important synthetic reactions that involve organometallic reagents or organometallic

catalytic cycles.

Chapter Review

Organometallic: A compound in which a carbon atom is directly bonded to a metal.

Catalyst: speeds up a reaction without being consumed. The catalyst may undergo

changes during the reaction, but it is always regenerated at the end.

Ligand dissociation and association opens or closes a coordination site. Dissociation and

association of carbonyl and phosphine ligands are of particular importance.

A metal complex reacts with a molecule A-B, adding the fragments A and B. The

coordination number and the oxidization state of the metal both increase by 2.

Metal complexes are frequently nucleophilic enough to react in an SN2-like reaction.

Catalytic cycles consist largely of a series of simple reaction steps. These steps may

include ligand dissociation and association, oxidative addition and reductive elimination,

and insertion and elimination.

Examples of catalytic cycles include: Deuteration of Benzene, Palladium catalyzed coupling

reactions: Suzuki, Buchwald-Hartwig, and Heck Reactions.

Several catalytic coupling reactions have been developed using palladium. These employ a

cycle in which palladium is sequentially oxidized and reduced between Pd(0) and Pd(II).

Suzuki Coupling Reaction: This is a coupling reaction between an organoboronic acid

and alkyl halide. Organoboranes, boronate esters, or potassium trifluoroborates can be

substituted for the boronic acid.

The Buchwald-Hartwig coupling reaction is similar to the Suzuki coupling, except that

one fragment is an amine. A version of this reaction uses a phenol, and the product is an

aromatic ether.

The Heck reaction is a palladium catalyzed coupling of a vinyl ester with an aryl halide.

This reaction is selective for the trans-olefin group.

Heck Reaction Mechanism: The key steps are: Oxidative addition of the aryl halide, Pd

ligand displacement by alkene, Insertion of aromatic group into the coordinated alkene,

Elimination, resulting in another alkene-Pd complex, Alkene ligand dissociation, Reductive

elimination of HBr.

P a g e | 52


The Monsanto process uses a rhodium catalyst to produce acetic acid from CO and

methanol. The process involves oxidative addition/reductive elimination and

insertion/elimination steps.

Hydroformylation uses a metal catalyst to convert an alkene, CO, and hydrogen to an

aldehyde.

Alkyllithiums react with Cu(I) salts to form lithium dialkylcopper reagents. These are also

known as Gilman reagents. The structure is complex, and it is dependent on the solvent

and the lithium cuprate structure.

Lithium dimethylcuprate is believed to exist as a monomer- or solvent separated ion pair

(SSIP)-dimer equilibrium in ethereal solvents.

THF solvation favors the monomer or SSIP. Less polar solvents like diethyl ether favor

the dimer. Higher aggregates have also been reported in diethyl ether.

Lithium cuprates interact with lithium halides and LiCN. Although the structure of these

species is not completely understood, it appears that mixed aggregates of lithium cuprates

and other lithium salts are formed.

The mechanism of Gilman coupling is not yet fully understood, and it is a research topic of

current interest.

The usual oxidization states for copper are Cu(I) and Cu(II). Cu(III) is an unusual

oxidization state for copper, but recent studies indicate that the Cu(I) reagent undergoes

an oxidative addition of alkyl halide.

Olefin Metathesis: A metathesis reaction involves a formal exchange of a: CR2 group of an

olefin. It usually involves a metallacyclobutane intermediate.

Olefin metathesis reactions are often performed as part of a catalytic cycle if the carbene

can be regenerated. These are frequently used in polymer synthesis.

The catalyst may be heterogeneous, or a separate phase from the reagents. Metal oxides

are commonly used as heterogeneous catalysts. The reaction mechanisms are often

poorly understood.

Heterogeneous metal catalysts are used to convert syngas into useful products, such as

alkanes and methanol. The water gas shift reaction produces additional H2 from CO and

water.

P a g e | 53


23: Organometallics III – Main Group Comparison

Chapter Summary

Some similarities in properties of main group and transition metals can be explained by

comparison of their molecular orbitals. Boron and transition metals can form cluster

compounds with some similarity in their structures and properties. This chapter will compare

some properties of main group and organometallic compounds.

Tutorial Features





Key Concepts

Similarities in properties between main group and organometallic compounds

Isolobal analogy

Metal-metal bonds

Main group and transition metal cluster compounds

Chapter Review

Cluster compound: A compound containing several of the same kind of atom arranged

as a polyhedron.

Main group elements require an octet (8) of electrons to fill the s- and 3-p orbitals of the

valence shell. Transition metals also have 5 d-orbitals that hold an additional 10 electrons.

Therefore, a transition metal “octet” is 18 electrons. The number of electrons needed to

complete the octet can be used to understand properties of main group and transition

metal compounds.

Halogens form diatomic molecules, add to alkenes, form insoluble Ag salts, and

disproportionate with Lewis bases. HCl, HBr, and HI are strong acids.

The 16 electron Fe(CO)4 unit is analogous to a divalent atom like oxygen or sulfur.

The isolobal analogy was introduced by Roald Hoffmann in 1982. Simply put, it states that

molecular fragments with similar orbital lobes can undergo similar bonding.

Molecular fragments are isolobal if the frontier orbitals have similar: Number, Shape,

Symmetry properties, approximate energy, Number of electrons.

To a good approximation, the bonding between the sp3 carbon and Cl is independent of

what else is attached to that carbon. All 3 sp3 carbon-containing fragments are isolobal

with each other.

The most direct isolobal analogy is for fragments with the same number of orbitals and

the same number of holes. The number of orbitals and holes can be defined by simple

formulae for metal complexes.

Electron Counting: The metal is considered as an ion. Lone pair donors and p-bonded

ligands donate 2 electrons as in the covalent model. Ligands such as Cl, H, methyl, and

phenyl are considered as anions and contribute 2 electrons.

X-type ligands are 1 electron ligands in the covalent model and 2 electron donors in the

ionic model. Examples include halogens, alkyl and aryl groups.

More complex ligands can be thought of as combinations of X and L ligand types.

The electron count of complex ligands is the sum of the electrons in each X and L

fragment. Each X ligand donates one electron in the covalent model and 2 electrons in the

ionic model. Each L ligand donates two electrons in both models.

P a g e | 54


Square planar complexes have a high energy d-orbital that is normally inaccessible. The

isolobal analogy is extended to square planar complexes using modified formulas for the

number of holes and available orbitals.

Metal-metal bonds were not definitively demonstrated until 1935. The W-W bond

distance in the W2Cl93- anion was shown by x-ray crystallography to be shorter than the

tungsten interatomic distance.

Metal-metal multiple bonds are also possible. Like main group elements, metals can form

single, double, and triple bonds.

Sigma, pi, and delta bonds are formed from the metal d-orbitals. Delta bonds are the

weakest, sometimes with bond energies of only about 6 kcal/mol.

Cluster compounds contain a framework of atoms. The best known main group cluster

compounds are the boranes.

Boron hydrides form cluster compounds similar to cluster compounds of many transition

metals.

Diborane is the dimer of borane. The two boron atoms are bridged by three center, two

electron bonds. These bonds are common in boron compounds.

Boron forms a large number of compounds with hydrogen, known collectively as boranes.

These boranes can be neutral compounds or anions.

The simplest class of higher order boranes are made of BnHn polyhedra, and exist as

BnHn2- anions. These are called closoboranes, named for the cage-like structure.

Nidoboranes: have a nest-like structure formed by removal of one corner of a closoborane

polyhedron. Arachnoboranes: have a web-like structure formed by removing two B-H

units from a closoborane polyhedron.

A carborane is a borane with one or more boron atoms substituted by carbon. An

equivalent borane formula is used to classify carboranes as closo, nido, or arachno.

Metallaboranes and metalla-carboranes have a metal substituting for a boron or carbon

atom. B is isolobal with a 13 electron fragment. B-H and C are isolobal with a 14 electron

fragment.

Transition metal cluster compounds are well known. Some metal clusters have structures

analogous to closo-, nido-, and arachno-boranes and carboranes.

Not all cluster compounds fit the borane analogy, and other structures are known. Some

metal clusters have a carbon atom with an unusual valence or coordination geometry.

CO can bridge 2 or 3 metal atoms. Metal cluster compounds with bridging CO ligands are

particularly common.

Metal-metal bond formation makes electron counting cumbersome for cluster compounds.

P a g e | 55


24: Bioinorganic and Environmental Chemistry

Chapter Summary

Metals may be part of a biomolecule, like an enzyme, for example. Coordination compounds

may also interact with biomolecules. These compounds may be toxic, or they may be used

to treat cancer or other diseases. This chapter will cover the role of metals in biological

molecules, molecules for photosynthesis and to carry and store oxygen, and coordination

compounds in medicine. The toxicity of metal compounds will also be discussed.

Tutorial Features





Key Concepts

The role of metals in biological molecules

Oxygen carrying and storing molecules

Compounds for photosynthesis

Coordination compounds in medicine

Toxicity of metal compounds

Chapter Review

Porphyrin: A metal complex with a derivative of the organic ring “porphine”.

Chemotherapy: the treatment of a malignancy or cancer with drug therapy.

Enzymes are large biomolecules that catalyze biological processes. The active site of an

enzyme is where the catalysis takes place.

The active site may allow other molecules to selectively fit into its size and shape. The pH

at the active site may be higher or lower than in the surrounding environment.

Metalloenzymes contain a metal atom at the active site. The metal may coordinate to the

substrate and hold it in place.

Chlorophylls are magnesium porphyrins that are involved in photosynthesis. The various

chlorophylls differ in their side chains.

The following first row transition metals have biological importance: Vanadium, Chromium

and Molybdenum, Manganese, Iron and Cobalt, Nickel, Copper and Zinc.

Vanadium compounds have a limited role in biological systems, but large concentrations

are present in some sea organisms.

Manganese is found in many enzymes and cofactors. It is needed for proper iron and

copper metabolism.

One important enzyme is superoxide dismutase (SOD). It destroys harmful superoxide

ion produced in the body during respiration. Other SOD enzymes use iron, nickel, copper,

and zinc.

Iron is among the most important transition metals in biological systems. The best known

iron biomolecule is hemoglobin, which carries oxygen in the blood and delivers it to the

cells.

The hemoglobin molecule consists of 4 globin subunits. Each can bind one O2 molecule,

and each bound O2 increases the binding constant for additional O2 molecules.

P a g e | 56


Each hemoglobin subunit contains an iron porphyrin unit. The geometry of the active site

changes as oxygen binds. Without oxygen, the iron is pulled out of the porphyrin ring

plane.

CO binds much more strongly to hemoglobin than oxygen and the reverse reaction is

slow. The product of CO and hemoglobin is called carboxyhemoglobin. It resembles a

typical metal carbonyl compound.

Myoglobin is another important iron containing biomolecule. Its function is oxygen

storage in the muscles. Myoglobin has one heme group per molecule and each molecule

can bind one oxygen molecule.

Copper complexes are widely found in biological systems. The blood of the horseshoe

crab uses a copper complex to carry oxygen.

Zinc is a necessary trace element in both plants and animals. It is one of the most

abundant elements in biological systems. Zinc fingers are fragments of proteins that can

coordinate to zinc. They can bind to DNA, RNA, proteins, or small molecules. DNA

binding is shown at the lower left.

The best known coordination compound used in medicine is cis diamminedichloroplatinum,

or cisplatin. Cisplatin kills cancer cells by interfering with DNA replication.

Since cancer cells divide rapidly, they are more sensitive than healthy cells to agents that

interfere with cell division. The trans isomer is not active against cancer cells.

The ammonia ligands in cisplatin are inert, or slow to be displaced by another ligand. The

chloride ligands are labile, and can be displaced by water, amine groups, sulfur, and other

ligands.

This accounts for both the mode of action and toxicity of cisplatin. Cisplatin therapy has

severe side effects. Mono- and diaquo species are partially responsible for the toxicity.

The toxicity of cisplatin and tumor cell resistance led to a search for new platinum drugs.

Carboplatin is like cisplatin with two inert ligands and two labile ligands.

Derivatives of ferrocene are being researched as antimalarial compounds. Chloroquine

(top left) is a common antimalarial but the parasite is developing resistance.

The toxicity of metal compounds is a major environmental concern.

Both metallic mercury and mercury compounds are toxic. Mercury compounds are

sometimes use in organic synthesis, such as oxymercuration in the synthesis of alcohols.

Mercury binds readily to sulfur, and deactivates many sulfur-containing enzymes. Mercury

accumulates in food chains, particularly those involving fish and shellfish.

Lead poisoning often occurs when lead compounds are ingested, such as from paint chips.

Lead ions have the proper size and charge to mimic calcium and inhibit enzyme reactions

that depend on calcium. It also binds to sulfur containing enzymes.

Most heavy metals are toxic. Cadmium exposure is mostly from industrial sources and

cigarette smoking. Arsenic inhibits the formation of ATP by several mechanisms. In one

mechanism it competes for phosphorous in biological pathways.

A chelate is a metal complex in which the metal is complexed to two or more atoms in the

ligand.

EDTA, or ethylenediamine-tetraacetate, is a common chelating agent for metals. Chelation

therapy is the use of drugs that will form a chelate with heavy metal ions and remove

them from the body.

Catalytic converters use metal catalysts to convert pollutants from car exhaust to less

harmful compounds. Palladium, platinum, and rhodium metals are used in catalytic

converters. Most late model cars use a 3-way converter. Together with air, this causes

unburned hydrocarbons to completely burn to CO2 and H2O.

Oxides of nonmetals form acidic solutions. Nitrogen combines with oxygen at high

temperatures during combustion of fossil fuels. The resulting nitrogen oxides, collectively

represented as NOx, contribute to acid rain formation. Coal and petroleum contain sulfur,

which forms sulfur oxides during combustion. These also contribute to acid rain.

inorganic chemistry rapid learning series

Documents