Atomic Models Atomic Theory I Love Chemistry

1. Atomic Models Atomic Theory I Love Chemistry

2. This power point will take us on a journey through time, thought, and imagination. Many ideas about the atom will be covered, how it was figured out, the good turns, the wrong turns, and where we are now, after more than 2000 years of trying. Welcome to chemistry proper. Fasten your seatbelts, it’s going to start slow but it will pick up speed shortly.

3. Subatomic Particles All atoms are made up of the same three particles. This chart outlines what you must remember about these three particles…

4. The history of the Atom, & the models of the atom over time.

5. Democritus came up with the original idea of the atom being the ultimately small particle of matter. He knew of no chemistry, but he was a good thinker. His ideas were more on the philosophical side rather than experimental. He felt that matter could be cut, cut, and cut again over and over until you reach the smallest bit, which he called ATOMOS, which is Greek for atom. The name stuck. Surely there is no “atom of plant”or atom of fish, atom of rock, etc. But there are atoms. I could only hope to think of something that anyone would remember for more than 2000 years! Chemistry is born.

6. John Dalton is an English scientist farmer. One day, bored watching the corn grow, or the cows milk, he goes to the barn and pretty much reinvents chemistry. He is sure that the atom idea is correct. He makes a bunch of measurements and figures out what becomes Atomic Theory around 1803. His ideas are nearly perfect, and we’ll examine them in the next slide. He did make a couple of small errors, but really, so what. This is a giant of thinkers, the person who described matter in such a way that we study his ideas over 200 years later.

7. Dalton's Atomic Theory 1 All matter is made up of extremely small particles called atoms. 2 All elements are made up of only one kind of atom, each identical to the others in properties and mass. 3 Two or more atoms can combine in small, whole number ratios, to form compounds 4 In a chemical reaction, atoms are re-arranged (combined or separated) – but not destroyed • Perfectly correct, still. • Atoms are now known to have isotopes, but all atoms of an element are CHEMICALLY identical. Isotopes have different numbers of neutrons. • Perfectly true still. • Also perfect, although in nuclear chemistry this sometimes is a rule to break.

8. JJ Thompson’s Plum Pudding Model of the Atom Chemistry marches on, and by 1897 he was able to prove the existence of electrons. Atoms were getting more sophisticated. He knew that they were electrically neutral, and he knew that electrons existed in the atom. Where? How? He wondered, as he ate his wife’s Plum Pudding (which is made up of ripped up bread soaked in cream and sugar, with bits of plum, then re-baked to a warm dessert to be covered in a sweet sauce. Atoms, he thought, plum pudding! That’s it! The electrons were embedded into a positive “mass” of atom stuff. That kept them neutral, allowed for electrons, and won him a Nobel Prize in 1906 in Physics. Wow, desserts really impacted us here. Stop for a second, look at this man’s eyes. He discovered electrons. Wow.

9. My hero, Ernest J. Rutherford, the man who discovered the nucleus (and more). Imagine, for about forty bucks this man comes up with one of the top 10 experiments of all time, and proves that Thompson was wrong about the model of the atom. Student learns from teacher, then surpasses him. Nobel Prize in chemistry. Walrus mustache. Who’s living better than this guy? The Gold Foil Experiment is on the next slide, and you must know it, know how to redraw it, and how to explain what he did, what he discovered, and what were the flaws in his thoughts. You’ve heard of that game, who’d you like to have dinner with from history? This is the man I’d have dinner with in that game, Who would you want to eat dinner with? RUTHERFORD

10. Gold Foil Experiment Rutherford thought to blast a thin, thin sheet of hammered gold foil with alpha particles. Alpha particles are positive radioactive particles given off by a radioactive element called polonium. An alpha particle has 2 neutrons and 2 protons, with a net charge of +2. He put them in the lead box and aimed them through a little hole at his gold foil. He soon discovered that most of these particles went directly through the solid gold foil, proving that most the gold atoms (all atoms) were empty space. Then he saw some of these positive charged particles deflected, at angles. He figured that the gold atom must have a large nucleus, which was positive charged as well (or the positive alpha particles would stick rather than bounce off). If the gold atoms had a positive nucleus, then the electrons were flying around outside (far far away from the nucleus) to keep the atoms neutral and mostly empty space. Since a few of his alpha particles bounced back, he was sure that the nucleus was in fact huge and dense besides. He discovers the nucleus (wow), puts electrons outside the nucleus (wow again), makes the nucleus positive (hmmm), and says correctly that atoms are mostly empty space (hard to swallow, but perfect!).

11. Rutherford's small flaws His ideas were great, but not quite perfect. The normal person would have a hard time understanding that all matter was by far, mostly empty space. I mean, do you even believe that yet? And, where were these pesky electrons? Flying around so far away? Why did they stay, not fly off, not collapse into the nucleus? These were real difficulties for EJ, and he had few answers. But luckily for him, one of his mathematical students, Niels Bohr, pictured here, was able to think his way through these puzzles and save his teacher. Niels Bohr

12. Bohr Model of the Atom The math was hard, and he could only “prove it” for hydrogen, but Bohr saved the Rutherford model of the atom. Bohr showed that electrons had to fly around the nucleus, not willy-nilly, but rather in specific orbits, which were also energy levels. As long as the electrons stayed in these fixed orbits, like planets around the Sun, they never lost any energy, and always kept flying right where they were supposed to. He could not do the math for larger atoms (no one could), but clearly the electrons don’t collapse into the nucleus, so he must have been right. He showed that the electrons fit into these orbits as follows:

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14. electron orbits First orbit: smallest, closest to the nucleus, holds up to two electrons. Electrons here have the lowest energy levels. Second orbit: bigger, holds up to 8 electrons. Electrons here have more energy associated with them than those electrons in the first orbit. Third orbit: highest energy so far, holds up to 8 electrons as well, although sometimes it can fit up to 18 electrons instead. The math gets so dense and we’ll just accept that this is correct. The electrons always hang out in the lowest possible energy levels. This is called the GROUND STATE. The electron configurations of your periodic tables have electrons in the ground state.

15. Ground vs. Excited State Bohr wasn’t quite done, he was a real thinker! He showed clearly that electrons in the ground state could absorb exact amounts of energy to be able to jump from their ground state orbital to the next highest energy levels. These levels could be imagined like steps that are uneven. Each step takes a certain amount of energy to climb onto. Each orbit would have a certain amount of energy to get to. Think of the lines as the orbits of the electrons.Orbits r1, r2, and r3 for this one atom. Electrons in orbit r1 could only get to r2 if theyabsorbed an exact amount of energy, not morenot less. To go from r2 to r3 would requirea different amount of energy. If an electron gained the right amount of energyit could jump from the ground state to an excited state, in a higher energy level orbit.

16. Let’s look at sodium for our example. The electron configuration from our periodic table shows us it is 2-8-1. Eleven electrons, three orbits (but we call them orbitals now), all electrons filling up the lowest electron orbits first, so that the atom has the lowest amount of energy associated with it. The electrons are in the ground state. Zap this atom with some energy (heat, light, electricity, etc.) and the electrons get “excited”. So excited that they might be able to “jump” to higher energy levels for a bit of time. Instead of 2-8-1, it might adjust to a 2-7-2 arrangement. Or a 2-6-3 if enough energy is absorbed. Now the atom is in the excited state. No new electrons, no losing any either, just moving some to higher energy orbits. (wow). When this unstable arrangement can’t be tolerated a moment more the electrons that got excited go back to the ground state, back to the lowest energy levels or orbits possible. To do so they must emit the exact amount of energy that they had gained. This energy is given off as colors of light called a spectrum, or the spectra.

17. Energy In: heat, electricity, or other forms. Same amount of Energy Out: Colored Light, colored flames, or spectra.

18. A copper flame burns green, other atoms or compounds will make their flames unique colors as well. The colors are given off by the excited electrons when they release that excess energy they gained by the heat, and that color is the mixture of unique wavelengths that our eyes see as one color. The excited electrons go back to the ground state by releasing the exact amount of energy they gained to become excited in the first place.

19. The atoms of lithium get excited when they heat up as well. Lithium is different than copper, it has different electrons in different orbitals. They get excited by gaining a different, unique amount of energy. Once excited they are also unstable, and revert back to the ground state. To go back to the ground state, they must release exactly as much energy it took for them to become excited in the first place. That is why each atom, or compound, gives off a unique mixture of colors when heated (or electrified, or energized in any way. The lithium flame is magenta, which is a cool name for a color.

20. Mixtures of Colors of Flames Our eyes see one color, sort of like when we see magic markers. They appear to be one color, but they really are mixtures of colors that fool our eyes in to seeing just one color. Special lenses, called REFRACTIVE LENSES can break apart the color flames, or color given off by Neon Lights, or even white lights. They separate out the specific wavelengths of light, so we can see a unique pattern of intensity, and wavelength, which can be readily measured in AP or college chemistry classes with a spectroscope. An example spectrograph is on the next slide. It shows the unique pattern of wavelengths given off by hydrogen and also by helium when the H and He electrons are returning to the ground state, from their excited states. Inexpensive cardboard refractive lenses will make us all look silly, but allow us to see the spectra produced by a variety of gases in our lab. The flames will be too jumpy to see spectra clearly in, but we will peek at the fires as well.

21. Spectra lines for H and for He. The gases give off a specific color of light to our eyes, but this color is really a mixture of specific wavelengths of specific intensities, which our eyes see as one color. These unique patterns are the same for H or for He anywhere in the universe. They are like the atoms optical fingerprints. The wavelengths can be measured on a numeric scale in nanometers. It’s very quantitative like this, although the mixture of colors is rather qualitative.

22. The Modern Model of the Atom During the early 20th century the quantum mechanics physics and math developed and the atom got both much more complex and easier to understand. The atom is just not that simple, even though we’d probably like it better that way. The math showed us that electrons do not fly in neat little orbits around the nucleus, rather they inhabit zones called orbitals. These orbitals have funny shapes, funny sorts of names, and the electrons are only in them most of the time. The electrons do not follow the “rules” we’d like to make them. The electrons are understood to be more statistical, say, most of the time they do this, or are here, but sometimes, they are over there, or somewhere else. It’s not neat, it’s probability math to the highest level. But, that’s what’s really happening in those atoms. This model is also known as the wave mechanical model, because the electrons don’t always act like little particles, they often act like waves of energy. The math is way beyond our course, chemistry and physics are big topics, deep breaths now, just remember the bold words above.

23. Isotopes Isotopes are chemically identical atoms that have a different number of neutrons than their fellow atoms. They could have a couple more or less, They all react the same way in all chemical reactions, they all have the same properties, except for the ones concerning mass. They each have a slightly different mass. That’s what makes them isotopes. Here are the three isotopes of hydrogen. Atop is the “normal” one, one proton, no neutrons, and one electron. Deuterium is identical chemically, but has one neutron. Tritium is also identical chemically, but has 2 neutrons.

24. The Periodic Table of the Elements Groups go up and down. Periods go left to right. Group 1 has hydrogen detached, and the alkali metals from lithium to francium. Hydrogen is a non-metal that bonds as if it belongs in group 1. Group 2 are the alkaline earth metals. Group 18 are the noble gases. Group 17 are the halogens. The central part of the table from group 3 through group 12 are the transitional metals. The dark colored staircase separates the metals on left, from the non-metals at right. Atoms touching the staircase are called the Metalloids, metals with some non-metallic properties, and some non-metals with some metallic properties. Two atoms that touch this staircase, Al and Po (the dog food exception) are just metals, not metalloids, but they touch the line because there is no avoiding it. The rest of the metals in the main part of the table are the “other” metals, although they too sometimes are called transitional. The bottom two rows are the “Alaska and Hawaii” of the table. Just like those two states on any American map, they get stuck where they fit, not where they really belong. Strangely all of these two rows fit below Yttrium in group 3. If we put them there, the table would become too wide to fit nicely, so we stick them on the bottom with a few asterisks to keep you aware of their real placement!

25. Alkali metals group 1 are reactive and soft. Alkaline earth metals are also reactive and soft, but bond differently than those group 1 metals. Transitional metals break many rules we’ll learn about bonding and electron location. Other metals often act like transitional metals. The La and the Ac series, at the bottom, are rare, and/or radioactive. We won’t be touching them in our class. The non-metals are very important. They are mostly gases, but some solids and liquid bromine live in that part of the table. Halogens are in group 17 and are all non-metal gases. Po is not a halogen. The Noble gases are in the last group, they do not make compounds with other atoms. Atoms in a group have many similarities, atoms in periods have few similarities. Period number is the number of electron orbitals, while group numbers just help keep us organized. Hydrogen is a strange element, it thinks it’s a metal, but it’s not.