The Nature of The Chemical Bond
AND THE STRUCTURE OF MOLECULES AND CRYSTALS:
An Introduction to Modern Structural Chemistry by Linus Pauling
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Linus Carl Pauling (1901-1994) was one of the most influential scientists of the 20th century, whose work bridged multiple disciplines such as chemistry, physics, and biology. Born in Portland, Oregon, Pauling was a pioneer in the field of quantum chemistry and molecular biology, forging new paths of scientific inquiry and fundamentally altering our understanding of the natural world.
Educated at Oregon State University and the California Institute of Technology, Pauling's groundbreaking research in the field of molecular structure won him the Nobel Prize in Chemistry in 1954. His work on the nature of the chemical bond, specifically his introduction of the concept of electronegativity and the elucidation of resonance structures, transformed the way chemists visualize molecules and their interactions.
Not only was Pauling a titan in the scientific field, but he was also a passionate advocate for peace. His efforts to curb nuclear weapon proliferation won him a second Nobel Prize, this time for Peace, in 1962, making him the only person to have been awarded two unshared Nobel Prizes.
"The Nature of the Chemical Bond", first published in 1939, is considered a classic in scientific literature. In this text, Pauling dives deep into the principles underlying chemical bonding and structure, bringing forth his revolutionary ideas that melded principles from quantum mechanics with structural chemistry. His revolutionary hybridization concept and the introduction of orbital theories stand as cornerstones of modern chemistry.
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Linus Pauling's educational journey was marked by early achievements and profound curiosity. His interest in chemistry was sparked during his adolescence, around the age of 15, when a friend's chemistry set became a tool for independent study. After graduating from Washington High School in Portland, Pauling attended Oregon Agricultural College (now Oregon State University), where he studied chemical engineering to meet his mother's hopes of him becoming a vocational chemist. However, his interest in theoretical chemistry grew, and he completed his bachelor's degree in this subject in 1922.
Intrigued by the recent advances in quantum mechanics, Pauling moved to the California Institute of Technology (Caltech) for his graduate studies, where he focused on X-ray crystallography — a technique that uses X-rays to study the structure of crystals. Under the guidance of Roscoe G. Dickinson, Pauling used this method to study the structure of minerals. His Ph.D. dissertation in 1925 was titled "The Determination with X-Rays of the Structures of Crystals," marking his early steps into the world of structural chemistry.
Pauling's deepening interest in quantum mechanics led him to Europe, where the field was rapidly evolving. With the support of a Guggenheim Fellowship, he spent 1926 and 1927 studying with physicists including Arnold Sommerfeld in Munich, Niels Bohr in Copenhagen, and Erwin Schrödinger in Zurich. This exposure to quantum mechanics was crucial as Pauling started to apply these principles to the understanding of chemical bonds, leading to his development of the field of quantum chemistry.
Returning to Caltech as a faculty member, Pauling focused on understanding the nature of the chemical bond. One of his earliest and significant successes was explaining the tetrahedral arrangement of electron pairs around the central atom in a molecule, using quantum mechanics. This understanding provided a theoretical basis for Gilbert N. Lewis's observation that atoms bond in a way that achieves an outer shell of eight electrons (the "octet rule").
In the mid-1930s, Pauling introduced the concept of electronegativity, defining it as the power of an atom in a molecule to attract electrons to itself. He developed the electronegativity scale, which is still used today to predict the type of bond that will form between two atoms.
His seminal work, "The Nature of the Chemical Bond," was published in 1939. It summarized his theories of chemical structure, including the concepts of hybridization and resonance, effectively fusing quantum physics and chemistry. This transformative work still stands as a fundamental text for students of chemistry.
Pauling's early education and research thus laid the groundwork for his later revolutionary contributions, marking him as one of the most influential scientists of the 20th century. His insatiable curiosity and his ability to combine theoretical physics with experimental chemistry played a crucial role in developing our modern understanding of the chemical bond.
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"The Determination with X-Rays of the Structures of Crystals" was Linus Pauling's doctoral dissertation, completed in 1925 under the guidance of his mentor, Roscoe G. Dickinson, at the California Institute of Technology (Caltech).
Pauling's research focused on X-ray crystallography, a powerful technique that allows scientists to study the arrangement of atoms within a crystal structure. When a crystal is hit with an X-ray beam, the X-rays are diffracted, and this diffraction pattern can be analyzed to determine the crystal's atomic structure. This technique was crucial in the early 20th century for revealing the structures of many types of crystals and molecules, including complex biological substances like proteins and DNA.
In his dissertation, Pauling explored the structures of various crystals, demonstrating the power and precision of X-ray crystallography in revealing atomic arrangements. His early investigations included the structure of molybdenite (MoS2), demonstrating the layered structure of molybdenum and sulfur atoms.
This foundational work provided Pauling with a deep understanding of structural chemistry and the use of X-ray diffraction as a tool, which later became instrumental in his theories about the nature of the chemical bond. Pauling's early work in X-ray crystallography was a precursor to his future pioneering research, highlighting the ways that physical techniques can provide invaluable insights into chemical structures.
It's important to note that the methods and techniques Pauling used in his dissertation were still in their infancy. His work, along with that of many other scientists during this time, helped refine and develop X-ray crystallography into the powerful and precise tool that it is today, used in a wide array of scientific disciplines, from materials science to pharmaceutical research.
12-4 ICE AND WATER; CLATHRATE COMPOINDS
This section seemed particularly interesting so let’s go through it in detail.
Investigations by Dennison (1921), Bragg (1922), Barnes (1929), Megaw (1934) and Hillesund (1942) into the crystal structure of ice by x-ray diffraction indicate each oxygen atom is surrounded by four other oxygen atoms at a distance of 2.76 Å.
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In a water molecule (H2O), the two hydrogen atoms are covalently bonded to the central oxygen atom. The molecule is also capable of forming two additional hydrogen bonds with two other water molecules, via the lone pairs of electrons on the oxygen atom. This forms a tetrahedral configuration around the oxygen atom: two close hydrogen atoms (part of the same water molecule), and two further away (each part of a different water molecule). The four surrounding oxygen atoms form a distorted tetrahedron around each oxygen atom.
The typical distance between the oxygen atom and the two closer hydrogen atoms is about 0.96 Å (angstroms). Meanwhile, the distance between the oxygen atom and the two more distant hydrogen atoms (i.e., those involved in hydrogen bonding) is about 1.76 Å. When you add these two distances together, you get the value mentioned in your statement: approximately 2.76 Å.
This is the approximate distance between an oxygen atom and the oxygen atoms of adjacent water molecules in the crystal lattice of ice Ih. This structure is responsible for the unique properties of ice, such as its lower density compared to liquid water.
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When water freezes into ice, it undergoes a fascinating transformation at the molecular level. Individual water molecules, each made up of one oxygen atom and two hydrogen atoms (H2O), align themselves in a way that maximizes the attractive forces between them, known as hydrogen bonds.
Hydrogen bonds occur due to the polarity of water molecules. In each molecule, the oxygen atom is more electronegative than the hydrogen atoms, meaning it pulls shared electrons closer to itself. This creates a charge separation within the water molecule: a slightly negative charge on the oxygen side (δ-) and a slightly positive charge on the hydrogen side (δ+). This polarity allows water molecules to form hydrogen bonds, where the δ+ hydrogen atom of one water molecule is attracted to the δ- oxygen atom of another.
Upon freezing, water molecules arrange themselves into a hexagonal lattice structure to optimize these hydrogen bonds. In this arrangement, each water molecule forms hydrogen bonds with four others, creating a three-dimensional network with a tetrahedral configuration around each water molecule. This open structure has a lot of empty space, which gives ice its lower density compared to liquid water.
Contrary to what one might initially think, it's not a process of the molecules repelling each other that leads to this open structure. Instead, it's about the system trying to lower its energy by maximizing the attractive hydrogen bonding interactions. This unique behavior contributes to many of the special properties of water and ice, and it is part of what makes water such a vital substance for life on Earth.
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Ice Ih
Ice Ih is the form of ice that is most familiar to us. It is the ice we see in our daily life, forming ice cubes, frost, and snowflakes. The 'h' in ice Ih stands for hexagonal, reflecting its crystal structure. In ice Ih, each water molecule is hydrogen bonded to four others, creating a three-dimensional network with a tetrahedral configuration around each water molecule. This results in an open, hexagonal lattice structure with a lot of empty space, giving ice Ih its lower density compared to liquid water.
Ice Ic
Ice Ic (the 'c' stands for cubic) is a less common form of ice. It is a metastable cubic crystalline variant of ice. It is produced at temperatures between 130 and 220 Kelvin, usually upon cooling from the gas phase. In the crystal structure of ice Ic, oxygen atoms are arranged in a diamond-like structure. Like ice Ih, each water molecule in ice Ic is hydrogen bonded to four others, but the arrangement leads to a cubic structure rather than a hexagonal one.
It's important to note that these are just two forms of ice out of many. Under different conditions of temperature and pressure, water ice can exist in several different solid states. These different forms are known as polymorphs and have different physical properties, even though they are all made up of the same water molecules. This polymorphism of ice is a topic of ongoing scientific study, and it contributes to the fascinating complexity of this common substance.
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Linus Pauling calculated the residual entropy of ice by considering the different ways water molecules can be arranged in the crystal lattice of ice. This calculation is based on the rules of hydrogen bonding in ice: each oxygen atom is hydrogen-bonded to four others, two through its hydrogen atoms (donor bonds) and two through its lone pairs of electrons (acceptor bonds).
In an ideal, defect-free crystal of ice, each water molecule is part of two hydrogen bonds as a donor and two hydrogen bonds as an acceptor. However, there are different ways these bonds can be arranged. In essence, the two hydrogen atoms of each water molecule can "point" towards any two of the neighboring oxygen atoms, leading to different arrangements or "states" for the crystal as a whole.
Pauling considered a specific water molecule and its four nearest neighbors in the ice lattice. He calculated the number of ways these five water molecules could arrange themselves while still obeying the hydrogen-bonding rules. This leads to six possible configurations, two of which have an extra degree of freedom because they can be transformed into another valid configuration by rotating a single water molecule about its oxygen atom.
After careful statistical consideration of these possibilities, Pauling concluded that each water molecule, on average, can exist in about 1.5 different "states" or orientations that are consistent with the hydrogen-bonding rules. This leads to a large number of total states for the crystal, which is related to the entropy of the system according to Boltzmann's entropy formula, S = k log W, where S is the entropy, k is Boltzmann's constant, and W is the number of possible states.
Thus, despite the regularity of the crystal lattice, the water molecules in ice have some degree of disorder due to their different possible orientations, leading to a small, non-zero residual entropy even at absolute zero temperature.
"The Nature of the Chemical Bond" is a classic book written by Linus Pauling, an American chemist and one of the most influential scientists of the 20th century. The book was first published in 1939 and is based on a series of lectures Pauling delivered at Cornell University in 1937. It is considered a groundbreaking work in the field of chemistry and has played a pivotal role in shaping the modern understanding of chemical bonding.
In "The Nature of the Chemical Bond," Pauling introduced several important concepts and principles, including the concept of electronegativity, which is the measure of an atom's ability to attract electrons in a chemical bond. He also developed the concept of resonance, which explains the stability and properties of molecules by considering multiple possible structures (resonance structures) that contribute to the actual structure of the molecule.
The book covers various types of chemical bonds, such as ionic, covalent, and metallic bonds, as well as hydrogen bonding and van der Waals forces. It also delves into the theories of valence, molecular orbitals, and crystal structures.
"The Nature of the Chemical Bond" remains an essential reference in the field of chemistry and has been a key resource for generations of chemists and scientists studying chemical bonding, molecular structure, and chemical reactions. Linus Pauling's work in this area earned him the 1954 Nobel Prize in Chemistry.
