Published on MATSE 81: Materials In Today's World (https://www.e-education.psu.edu/matse81)

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Lesson 3: Atomic Structure and Bonding

Overview

Electronic configuration for elements and the interatomic bonding between atoms and molecules determine some of the important properties of solid materials, including a correlation between bonding type and material classification—namely, ionic bonding (ceramics), covalent bonding (polymers), metallic bonding (metals), and van der Waals bonding (molecular solids). In this lesson, we will review briefly atomic structure, electron configurations in atoms, the periodic table, and atomic and interatomic bonding. These fundamental and important concepts will be applied to the understanding of solid materials in this and subsequent lessons of this course.

We will see in later lessons that important properties of solid materials depend on the way in which the atoms are arranged. In this lesson, we will consider some fundamental and important concepts about how the atoms are held together that compose a solid. These concepts: atomic structure, electron configuration, the periodic table, and the various types of primary and secondary interatomic bonds, are discussed with the assumption that the student has already encountered this material in a high school chemistry course. 

Learning Objectives

By the end of this lesson, you should be able to:

  • Describe and compare the Bohr and wave mechanical atomic models.
  • Describe the important quantum-mechanical principle that relates to electron energies.
  • Recognize the effect of the Pauli exclusion principle on atomic structure.
  • Produce the electronic configuration for the ground state of a given element and any ions.
  • Identify the locations of metallic, non-metallic, and intermediate elements on the periodic table.
  • Describe the general rule for electronegativity on the periodic table.
  • Contrast the behavior of valence electrons for electropositive with the valence electrons of electronegative elements.
  • Briefly describe ionic, covalent, metallic, hydrogen, and van der Waals bonds.
  • Find examples of materials with the following bond types; ionic, covalent, metallic, hydrogen, and van der Waals bonds.
  • Briefly explain the fact that water expands upon freezing to ice from its liquid phase.

Lesson Roadmap

Lesson 3 will take us 1 week to complete. Please refer to Canvas for specific due dates.

Lesson Roadmap
To Read

Read pp 37-65 (Ch. 3) in Introduction to Materials ebook

Reading on course website for Lesson 3

To Watch Chapters from Hunting the Elements, TED-Ed talks on Atoms and Periodic Table
To Do Lesson 3 Quiz

Questions?

If you have general questions about the course content or structure, please post them to the General Questions and Discussion forum in Canvas. If your question is of a more personal nature, feel free to send a message to the instructor through Canvas email. The instructor will check daily to respond.

Reading Assignment

Things to consider...

While you read the material for this lesson in your e-book and on the course website, use the following questions to guide your learning. Also, remember to keep the learning objectives listed on the previous page in mind as you learn from this text. 

  • What is the main difference between the Bohr and the wave particle atomic models?
  • How does the electronic configuration of an atom determine its materials classification?
  • What are some of the materials properties that the location of the element on the Periodic Table predicts?
  • How does bonding affect the materials properties of atoms and molecules?
  • How does bonding type determine a material's probable materials characterization?

Reading Assignment

Read pp 37-65 (Ch. 3) in Introduction to Materials ebook

The Atom

The word atom is derived from the ancient Greek adjective atomos, meaning "uncuttable" or "indivisible." The earliest concepts of the nature of the atom were debated in ancient India and ancient Greece. We now know that the atom has a nucleus composed of protons and neutrons surrounded by clouds of electrons. The protons are positively charged, electrons are negatively charged, and neutrons possess no charge. Neutrons and protons are held in the nucleus by the nuclear force, and neutrons are not simply a proton plus an electron. In fact, neutrons are required to make the nucleus stable once you have more than one proton in the nucleus.

Atoms are the fundamental building blocks of matter; they cannot be divided using chemicals. Chemical reactions to move electrons can affect how atoms bind to each other but cannot be used to divide atoms. Most of the mass of the atom is located in the nucleus, with the mass of the proton roughly equal to the larger neutron, but 1840 times the mass of the electron. In contrast, most of the volume of the atom is filled with electrons. Now please watch this brief (5:22) video on the (brief) history of atomic theory.

To Watch

TED talk, the 2,400-year Search for the Atom
Click for transcript of The 2,400-year Search for the Atom

What do an ancient Greek philosopher and a 19th century Quaker have in common with Nobel Prize-winning scientists? Although they are separated over 2,400 years of history, each of them contributed to answering the eternal question: what is stuff made of?

It was around 440 BCE that Democritus first proposed that everything in the world was made up of tiny particles surrounded by empty space. And he even speculated that they vary in size and shape depending on the substance they compose. He called these particles "atomos," Greek for indivisible. His ideas were opposed by the more popular philosophers of his day. Aristotle, for instance, disagreed completely, stating instead that matter was made of four elements: earth, wind, water, and fire, and most later scientists followed suit.

Atoms would remain all but forgotten until 1808 when a Quaker teacher named John Dalton sought to challenge the Aristotelian theory. Whereas Democritus's atomism had been purely theoretical, Dalton showed that common substances always broke down into the same elements in the same proportions. He concluded that the various compounds were combinations of atoms of different elements, each of a particular size and mass that could neither be created nor destroyed. Though he received many honors for his work, as a Quaker, Dalton lived modestly until the end of his days.

Atomic theory was now accepted by the scientific community, but the next major advancement would not come until nearly a century later with the physicist J.J. Thompson's 1897 discovery of the electron. In what we might call the chocolate chip cookie model of the atom, he showed atoms as uniformly packed spheres of positive matter filled with negatively charged electrons. Thompson won a Nobel Prize in 1906 for his electron discovery, but his model of the atom didn't stick around long.

This was because he happened to have some pretty smart students, including a certain Ernest Rutherford, who would become known as the father of the nuclear age. While studying the effects of X-rays on gases, Rutherford decided to investigate atoms more closely by shooting small, positively charged alpha particles at a sheet of gold foil. Under Thompson's model, the atom's thinly dispersed positive charge would not be enough to deflect the particles in any one place. The effect would have been like a bunch of tennis balls punching through a thin paper screen. But while most of the particles did pass through, some bounced right back, suggesting that the foil was more like a thick net with a very large mesh. Rutherford concluded that atoms consisted largely of empty space with just a few electrons, while most of the mass was concentrated in the center, which he termed the nucleus. The alpha particles passed through the gaps but bounced back from the dense, positively charged nucleus. But the atomic theory wasn't complete just yet.

In 1913, another of Thompson's students by the name of Niels Bohr expanded on Rutherford's nuclear model. Drawing on earlier work by Max Planck and Albert Einstein he stipulated that electrons orbit the nucleus at fixed energies and distances, able to jump from one level to another, but not to exist in the space between. Bohr's planetary model took center stage, but soon, it too encountered some complications. Experiments had shown that rather than simply being discrete particles, electrons simultaneously behaved like waves, not being confined to a particular point in space. And in formulating his famous uncertainty principle, Werner Heisenberg showed it was impossible to determine both the exact position and speed of electrons as they moved around an atom. The idea that electrons cannot be pinpointed but exist within a range of possible locations gave rise to the current quantum model of the atom, a fascinating theory with a whole new set of complexities whose implications have yet to be fully grasped.

Even though our understanding of atoms keeps changing, the basic fact of atoms remains, so let's celebrate the triumph of atomic theory with some fireworks. As electrons circling an atom shift between energy levels, they absorb or release energy in the form of specific wavelengths of light, resulting in all the marvelous colors we see. And we can imagine Democritus watching from somewhere, satisfied that over two millennia later, he turned out to have been right all along.

Credit: Theresa Doud, TED-Ed

To Read

Now that you have watched the video, please go to your e-textbook and read the first four sections (pages 36 to 46 in Chapter 3 of Materials for Today's World, Custom Edition for Penn State University) of this lesson's reading. When finished with the reading proceed to the next web page.

The Periodic Table

The periodic table classifies the elements according to their electron configuration. The scientist given credit for the modern periodic table is Russian chemist Dmitri Mendeleev. Please watch the following video (4:24) which explains the true genius of what Mendeleev accomplished.

To Watch

TED Talk: The Genius of Mandeleev's Periodic Table
Click for transcript of The Genius of Mendeleev's Periodic Table.

The periodic table is instantly recognizable. It's not just in every chemistry lab worldwide, it's found on t-shirts, coffee mugs, and shower curtains. But the periodic table isn't just another trendy icon. It's a massive slab of human genius, up there with the Taj Mahal, the Mona Lisa, and the ice cream sandwich -- and the table's creator, Dmitri Mendeleev, is a bonafide science hall-of-famer. But why? What's so great about him and his table? Is it because he made a comprehensive list of the known elements? Nah, you don't earn a spot in science Valhalla just for making a list. Besides, Mendeleev was far from the first person to do that. Is it because Mendeleev arranged elements with similar properties together? Not really, that had already been done too. So what was Mendeleev's genius?

Let's look at one of the first versions of the periodic table from around 1870. Here we see elements designated by their two-letter symbols arranged in a table. Check out the entry of the third column, fifth row. There's a dash there. From that unassuming placeholder springs the raw brilliance of Mendeleev. That dash is science. By putting that dash there, Dmitri was making a bold statement. He said -- and I'm paraphrasing here -- Y'all haven't discovered this element yet. In the meantime, I'm going to give it a name. It's one step away from aluminum, so we'll call it eka-aluminum, "eka" being Sanskrit for one. Nobody's found eka-aluminum yet, so we don't know anything about it, right? Wrong! Based on where it's located, I can tell you all about it. First of all, an atom of eka-aluminum has an atomic weight of 68, about 68 times heavier than a hydrogen atom. When eka-aluminum is isolated, you'll see it's a solid metal at room temperature. It's shiny, it conducts heat really well, it can be flattened into a sheet, stretched into a wire, but its melting point is low. Like, freakishly low. Oh, and a cubic centimeter of it will weigh six grams.

Mendeleev could predict all of these things simply from where the blank spot was, and his understanding of how the elements surrounding it behave. A few years after this prediction, a French guy named Paul Emile Lecoq de Boisbaudran discovered a new element in ore samples and named it gallium after Gaul, the historical name for France. Gallium is one step away from aluminum on the periodic table. It's eka-aluminum. So were Mendeleev's predictions right? Gallium's atomic weight is 69.72. A cubic centimeter of it weighs 5.9 grams. It's a solid metal at room temperature, but it melts at a paltry 30 degrees Celcius, 85 degrees Fahrenheit. It melts in your mouth and in your hand.

Not only did Mendeleev completely nail gallium, he predicted other elements that were unknown at the time: scandium, germanium, rhenium. The element he called eka-manganese is now called technetium. Technetium is so rare it couldn't be isolated until it was synthesized in a cyclotron in 1937, almost 70 years after Dmitri predicted its existence, 30 years after he died. Dmitri died without a Nobel Prize in 1907, but he wound up receiving a much more exclusive honor. In 1955, scientists at UC Berkeley successfully created 17 atoms of a previously undiscovered element. This element filled an empty spot in the periodic table at number 101, and was officially named Mendelevium in 1963. There have been well over 800 Nobel Prize winners, but only 15 scientists have an element named after them. So the next time you stare at a periodic table, whether it's on the wall of a university classroom or on a five-dollar coffee mug, Dmitri Mendeleev, the architect of the periodic table, will be staring back.

Credit: Lou Serico, TED-Eed

As mentioned in the video the true power of Mendeleev’s periodic table was the predictive ability of his table. This concept is at the heart of science. Scientists cannot just model behavior, but are required to make predictions, which later can be verified or refuted, thus, providing a test for the validity of their model or theories. It is interesting to note that Mendeleev’s work in the 1870s preceded the discovery of the atom which occurred with J.J. Thompson’s discovery of the electron in 1897 and the later work on the nucleus after 1900.

To Read

Now proceed to your e-textbook and finish reading this lesson’s reading assignment (pages 47 to 64 in Chapter 3 of Materials for Today's World, Custom Edition for Penn State University). Please proceed to the next webpage when you have completed this reading assignment.

Bonding and Bonding Type - Material Correlations

As you've recently read, there are four principal bonding types: ionic, covalent, metallic, and van der Waals. Ionic bonding involves the exchange of electrons between atoms to complete shells, either by adding or giving up electrons. The resulting atoms are oppositely charged and attract each other, resulting in an ionic bond. Covalently bonded materials have bonds in which electrons are shared between atoms. In metallic bonding, a "sea of electrons" is uniformly distributed throughout the solid and acts as a glue to hold the atoms together. Van der Waals bonds are relatively weak compared to the other three principal bond types and result when attractive forces from permanent or induced dipoles form.

diagram showing the 4 bonding types as discussed in the text above
Bonding tetrahedron.
Credit: Callister

In addition, the reading noted a correlation between materials classification and bonding time. Ionic bonding is associated with ceramics, covalent bonding is associated with polymers, metallic bonding is associated with metals, and van der Waals bonding is associated with molecular solids. As we study materials in further detail in this course we will utilize these associations to explain observed materials properties in the different materials classifications. Before we proceed to this lesson’s video assignment, there are a couple of more topics that I would like to address. Your textbook highlighted water as a material of importance and its volume expansion upon freezing. We will explore this topic further in the next section.

Water (Its Volume Expansion Upon Freezing)

Water is an extremely important molecule for life as we know it. An uncommon property that water possesses is the fact that frozen water (ice) is less dense than liquid water. This effect occurs due to the structure that occurs when water is cooled to form ice. The following video (3:55) takes a lighthearted approach to explain why ice floats.

To Watch

TED talk: Why Does Ice Float in Water?
Click for transcript of Why Does Ice Float in Water?

Water is the liquid of life. We drink it, we bathe in it, we farm, cook, and clean with it. It's the most abundant molecule in our bodies. In fact, every life form we know of would die without it. But most importantly, without water, we wouldn't have iced tea. Mmmm, iced tea.

Why do these ice cubes float? If these were cubes of solid argon in a cup of liquid argon, they would sink. And the same goes for most other substances. But solid water, a.k.a. ice, is somehow less dense than liquid water. How's that possible?

You already know that every water molecule is made up of two hydrogen atoms bonded to one oxygen atom. Let's look at a few of the molecules in a drop of water, and let's say the temperature is 25 degrees Celcius. The molecules are bending, stretching, spinning, and moving through space. Now, let's lower the temperature, which will reduce the amount of kinetic energy each of these molecules has so they'll bend, stretch, spin, and move less. And that means that on average, they'll take up less space.

Now, you'd think that as the liquid water starts to freeze, the molecules would just pack together more and more closely, but that's not what happens. Water has a special kind of interaction between molecules that most other substances don't have, and it's called a hydrogen bond. Now, remember that in a covalent bond two electrons are shared, usually unequally, between atoms. In a hydrogen bond, a hydrogen atom is shared, also unequally, between atoms. One hydrogen bond looks like this. Two look like this. Here's three and four and five, six, seven, eight, nine, ten, eleven, twelve, I could go on. In a single drop of water, hydrogen bonds form extended networks between hundreds, thousands, millions, billions, trillions of molecules, and these bonds are constantly breaking and reforming.

Now, back to our water as it cools down. Above 4 degrees Celcius, the kinetic energy of the water molecules keeps their interactions with each other short. Hydrogen bonds form and break like high school relationships, that is to say, quickly. But below 4 degrees, the kinetic energy of the water molecules starts to fall below the energy of the hydrogen bonds. So, hydrogen bonds form much more frequently than they break and beautiful structures start to emerge from the chaos.

This is what solid water, ice, looks like on the molecular level. Notice that the ordered, hexagonal structure is less dense than the disordered structure of liquid water. And you know that if an object is less dense than the fluid it's in, it will float. So, ice floats on water, so what? Well, let's consider a world without floating ice. The coldest part of the ocean would be the pitch-black ocean floor, once frozen, always frozen. Forget lobster rolls since crustaceans would lose their habitats, or sushi since kelp forests wouldn't grow. What would Canadian kids do in winter without pond hockey or ice fishing? And forget James Cameron's Oscar because the Titanic totally would have made it. Say goodbye to the white polar ice caps reflecting sunlight that would otherwise bake the planet. In fact, forget the oceans as we know them, which at over 70% of the Earth's surface area, regulate the atmosphere of the whole planet. But worst of all, there would be no iced tea. Mmmmm, iced tea.

Credit: George Zaidan and Charles Morton, TED-Ed

Now that you have watched this video, please proceed to the next section which highlights van der Waals forces and the gecko’s ability to walk on ceilings.

How Do Geckos Defy Gravity?

Close-up of gecko from underneath
Gecko feet
Credit: Magen McCrarey via Flickr [1]

Please watch the following video (4:29) which explains how geckos use van der Waals forces to walk on ceilings. While watching this video, see if you can answer the following question: how is the gecko’s ability to walk on ceilings an example of nanomaterials?

To Watch

TED talk: How do Geckos Defy Gravity?
Click for the transcript of How do Geckos Defy Gravity?

It's midnight and all is still, except for the soft skittering of a gecko hunting a spider. Geckos seem to defy gravity, scaling vertical surfaces and walking upside down without claws, adhesive glues or super-powered spiderwebs. Instead, they take advantage of a simple principle: that positive and negative charges attract. That attraction binds together compounds, like table salt, which is made of positively charged sodium ions stuck to negatively charged chloride ions. But a gecko's feet aren't charged and neither are the surfaces they're walking on. So, what makes them stick?

The answer lies in a clever combination of intermolecular forces and structural engineering. All the elements in the periodic table have a different affinity for electrons. Elements like oxygen and fluorine really, really want electrons, while elements like hydrogen and lithium don't attract them as strongly. An atom's relative greed for electrons is called its electronegativity. Electrons are moving around all the time and can easily relocate to wherever they're wanted most. So when there are atoms with different electronegativities in the same molecule, the molecule's cloud of electrons gets pulled towards the more electronegative atom. That creates a thin spot in the electron cloud where positive charge from the atomic nuclei shines through, as well as a negatively charged lump of electrons somewhere else. So the molecule itself isn't charged, but it does have positively and negatively charged patches.

These patchy charges can attract neighboring molecules to each other. They'll line up so that the positive spots on one are next to the negative spots on the other. There doesn't even have to be a strongly electronegative atom to create these attractive forces. Electrons are always on the move, and sometimes they pile up temporarily in one spot. That flicker of charge is enough to attract molecules to each other. Such interactions between uncharged molecules are called van der Waals forces. They're not as strong as the interactions between charged particles, but if you have enough of them, they can really add up.

That's the gecko's secret. Gecko toes are padded with flexible ridges. Those ridges are covered in tiny hair-like structures, much thinner than a human hair, called setae. And each of the setae is covered in even tinier bristles called spatulae. Their tiny spatula-like shape is perfect for what the gecko needs them to do: stick and release on command. When the gecko unfurls its flexible toes onto the ceiling, the spatulae hit at the perfect angle for the van der Waals force to engage. The spatulae flatten, creating lots of surface area for their positively and negatively charged patches to find complimentary patches on the ceiling. Each spatula only contributes a minuscule amount of that van der Waals stickiness. But a gecko has about two billion of them, creating enough combined force to support its weight. In fact, the whole gecko could dangle from a single one of its toes. That super stickiness can be broken, though, by changing the angle just a little bit. So, the gecko can peel its foot back off, scurrying towards a meal or away from a predator.

This strategy, using a forest of specially shaped bristles to maximize the van der Waals forces between ordinary molecules has inspired man-made materials designed to imitate the gecko's amazing adhesive ability. Artificial versions aren't as strong as gecko toes quite yet, but they're good enough to allow a full-grown man to climb 25 feet up a glass wall. In fact, our gecko's prey is also using van der Waals forces to stick to the ceiling. So, the gecko peels up its toes and the chase is back on.

Credit: Eleanor Nelsen, TED-Ed

So now that you have watched the video, can you see how this is an example of nanomaterials? A nanomaterial can utilize size and structure to perform unique abilities. The gecko utilizes van der Waals forces which operate on the scale of nanometers. In addition, the gecko utilizes the unique geometry of its feet to adhere to and release from surfaces. This is an example of using structure (or geometry) to perform a unique ability. At this time, please proceed to the lesson’s video assignment.

Video Assignment: Hunting the Elements

Video Assignment

Now please go to Lesson 3 in Canvas and watch chapters 3, 4, 5, 6, and 7 from the NOVA "Hunting the Elements" documentary. You will be quizzed on the content of these videos.

Summary and Final Tasks

Summary

Presented in this lesson were several fundamental and important concepts—namely, atomic structure, electron configurations in atoms and the periodic table, and the various types of inter-atomic bonds that hold together the atoms that compose a solid. The various types of atomic bonding, which are determined by the electron structures of the individual atoms, along with geometric atomic arrangements can determine some of the important properties of solid materials. Later in the course, we will move to the next level of the structure of materials, specifically, to some of the geometric atomic arrangements that may be assumed by atoms in the solid state.

Reminder - Complete all of the Lesson 3 tasks!

You have reached the end of Lesson 3! Double-check the to-do list on the Lesson 3 Overview page to make sure you have completed all of the activities listed there before you begin Lesson 4.


Source URL:https://www.e-education.psu.edu/matse81/node/2113

Links
[1] https://www.flickr.com/photos/95784928@N07/12127892834/in/photolist-jtGC1A-nf6vTM-ej7Evj-nUdA8C-8sS6Pg-gqxvHz-8sS4Zr-er3M83-8gPMpp-rvcMf8-7v46ad-93Wvs-nHDr9f-RfWtq7-2fzLrG-8gPMi6-nPnpra-fzvjEK-boFvGo-er3WAN-er4h7Q-5r8h85-inse7U-AiUbcL-oRsCxd-i3GLgJ-DNDQ4t-AAvoPi-bp78z7-eq7Kz6-44TdcZ-ptAz3H-ce89wY-mkwKkr-533ZDa-nm2KQ6-er4eH5-oQkQu-eq7Duv-8oJUL7-dA9hK-eq7Bdi-eq7S74-eq7yWH-er4cmU-6j9YHr-btyqCD-er42VC-3iwikt-6uHT9N