The History of Atomic Theory

1. The History of Atomic Theory

2. Atomic Models • At right is the image most people have of what an atom looks like. It shows electrons in various energy levels orbiting a nucleus. • Models use familiar ideas to explain unfamiliar facts observed in nature. Models can be changed as new information is collected. • So, is this the most accurate model of the atom? Or is it far from reality? • This presentation discusses six major models of what we call “atoms”. Each new model was developed based on new discoveries at the time. The Six Atomic Models are: • The Greek Model • Dalton’s Model • Thomson’s Model • Rutherford’s Model • Bohr’s Model • Wave-Mechanical Model

3. 1. The Greek Model • Democritus was a Greek philosopher who began the search for a description of matter more than 2400 years ago (around 400 BC). • He asked: could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided?

4. It’s only logical! • Democritus knew this question could not be easily answered directly. Let’s say you were breaking up a piece of rock -- how would you know when the smallest piece of rock was actually reached? • Democritus logically concluded that a piece of matter could not be divided into smaller pieces forever. Eventually the smallest possible piece would be obtained. • Since this piece is the smallest possible it would be indivisible, meaning it could not be further divided.

5. He named the smallest possible piece of matter “atomos,” which means “not to be cut.” • To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes. They were infinite in number, always moving, and capable of joining together. • However, Democritus’ theory was ignored and forgotten for more than 2000 years! Why? • The ancient Greeks did not experiment, but tried to win arguments through logic and debate. This would cause a problem for Democritus.

6. The famous philosopher Aristotle had a different theory of matter. He proposed that matter was composed of four elements: earth,fire, air and water. These four elements were supposedly blended in different proportions to make up all the various types substances in the world. • Aristotle had such a great reputation as a debater and logical thinker that many of his ideas were simply accepted as true without any sort of experimental evidence. • It wasn’t until the early 19th century that Aristotle’s idea of the nature of matter was seriously challenged.

7. 2: Dalton’s Model • John Dalton proposed his “modern” atomic theory in 1803, about 2000 years after Democritus and Aristotle. • Prior to Dalton matter was considered to be continuous in nature—meaning it could be infinitely divided--and not composed of indivisible bits as Democritus had proposed. John Dalton 1766-1824 See Sec 4.3 of text.

8. Dalton’s Atomic Theory • All elements are composed of atoms. • Atoms of the same element are identical. • Atoms of a given element are different from those of another element. • Compounds are formed by the joining of atoms of two or more elements. A given compound always has the same relative number and types of atoms. • Atoms are indivisible and may not be created or destroyed in chemical reactions. Chemical reactions simply rearrange how atoms are grouped together.

9. Dalton thought of atoms as little more than very tiny spheres – much like billiard balls. He knew nothing about any subatomic parts. • By experimentation he showed that compounds are formed by the joining of atoms of two or more elements. A given compound always has the same relative number and types of atoms. • Atoms are indivisible and may not be created or destroyed in chemical reactions. Chemical reactions simply rearrange how atoms are grouped together. • By experimentation he showed that compounds could be formed in ratios However, do the points of Dalton’s theory sound familiar? • They should! Dalton’s atomic theory became the foundation of modern chemistry.

10. 3: Thomson’s Plum Pudding Model • See section 4.5 in text. • In 1897, the English scientist J.J. Thomson proposed that atoms themselves are made of even smaller, sub-atomic, particles. • Thomson (among others) experimented with cathode ray tubes. When such a tube is evacuated of air and an electric current is passed through it, a glowing region (called cathode rays) was observed.

11. Cathode rays are so named because they are emitted at the negative (cathode) end and travel to the anode (positive) end of the tube. • Cathode rays are deflected towards a positively charged plate which showed they are composed of negatively-charged particles that have mass. • Regardless of the type of metal used at the cathodes and anodes of the vacuum tubes, the cathode rays were seen and had the same properties.

12. From these observations Thomson concluded that cathode rays were actually a stream of negatively charged particles that came from within the atom. • If particles smaller than the atom existed, then atoms themselves were divisible, not indivisible. • Thomson called these negatively-charged particles “corpuscles”. We, of course, know them as today as electrons. • Since atoms were neutral he reasoned that there must also be positively-charged parts in the atom, but he could never identify them. • http://www.youtube.com/watch?v=IdTxGJjA4Jw&list=PLD6A43875B9DEC57E&index=24

13. Thompson proposed a model of the atom that became popularly known as the “PlumPudding” model. • Atoms were made from a positively-charged substance with negatively-charged electrons scattered about, like raisins in a pudding. • The “Plum Pudding” model was also proposed around 1900 by William Thompson (better known as Lord Kelvin), who was no relation to J. J. Thompson.

14. 4: Rutherford’s Model • See sections 4.5, 10.1 in text • In 1910, New Zealand physicist Ernst Rutherford performed a series of experiments to study the structure of atoms.

15. Rutherford’s experiment (see next slide) involved firing a stream of positively (+) charged “bullets” – called alpha particles (actually, helium nuclei) -- at a thin sheet of gold foil about 2000 atoms thick. • Most of the positively charged alpha particle “bullets” passed nearly straight through the gold atoms in the sheet of foil. • However, some of the alpha particles bounced away from the gold sheet as if they had hit something solid. • How could something like this be explained given the current model of what atoms were?

17. Rutherford concluded that the gold atoms in the sheet were mostly open space, and not like a solid “pudding.” • He further concluded that both the mass and the positive charge of an atom are concentrated within a tiny fraction of the atom’s volume, which he called the nucleus. • The few alpha particle “bullets” that were deflected were bouncing off the tiny, dense nucleus at the center of the mostly empty atom. Those not deflected were passing through empty space!

18. - - - - - - - + + + + + + + + + + + + + + What Rutherford expected: alpha particle paths virtually uninterrupted as they went through the atom because at the time it was believed that atoms had the same consistency throughout.

19. - + - - - + + + + + + - - - What he observed: scattering of some alpha particles as they bounced off the dense nucleus at the center of the atom. Most of the alpha particles simply passed through the empty space of the atoms.

20. - - - + + + + - • These observations certainly changed the idea of how atoms could be pictured. • Rutherford’s atomic model contained positive charges in a very tiny, very dense nucleus. • The negatively charged electrons filled the remaining much larger volume of the atom and orbited outside the nucleus.

21. By 1920 Rutherford established that the nucleus of an atom consisted of positively charged particles, called protons, and, for stability, must contain neutral particles with the same mass as protons. • Hydrogen is the exception with a single proton in its nucleus. • The neutral particle, called the neutron, was eventually discovered in 1932 by James Chadwick, a student of Rutherford’s.

22. Another example of the Rutherford model of the atom in which electrons can orbit the atom in an infinite number of paths.

23. 5: Bohr Model • See section 10.5 in text. • Rutherford’s model, however, could not explain how negatively-charged electrons could maintain stable orbits around a positively-charged nucleus. Electrons should be drawn into the nucleus. But this is not observed. So what was keeping electrons in stable orbits – and atoms in existence? • In 1913, the Danish scientist Niels Bohr proposed that electrons are restricted to certain fixed (quantized) orbits around the nucleus.

24. These orbits, or energy levels (n=1, n=2, n=3, etc.), are only located at specific distances from the nucleus. • If electrons gain energy (become energized) they can jump to a higher fixed energy level. • When electrons lose energy they drop back down to a lower energy level. The energy lost is released (emitted) as a photon of light of a specific frequency (color).

25. A quantum (fixed amount) of energy is required to move electrons to the next highest level. Likewise, the same fixed amount of energy is emitted when an electron drops to a lower energy. • Bohr’s model explained the observed emission spectrum data for hydrogen (see fig 10.11, pg 286).

26. In Bohr’s model, the position of electrons around the nucleus is analogous to the steps of a stair (or rungs of a ladder). • Electrons cannot exist between energy levels, just like you can’t stand between steps on a stair (or rungs on ladder).

27. 6: The Wave Mechanical Model • See sections 10.6, 10.7, 10.8 in text. • Also called the quantum mechanical model, or cloud model. • By mid-1920s the Bohr model was shown to be incorrect because it only worked well for the hydrogen atom. • A new model suggested that electrons exhibit wave-like as well as particle-like behavior (just as photons do). • Electrons do not “orbit” the nucleus as in the Bohr model, but are located within regions of space around the nucleus.

28. Scientists responsible for the model: • Louis de Broglie was the first to support the idea that electrons exhibit wave characteristics. • Erwin Schrodinger developed a mathematical wave formula that described electrons as waves. • Werner Heisenberg showed that it is impossible to know both the exact location and the exact speed of an electron (the Uncertainty Principle).

29. Summary of wave-mechanical model: • The nucleus remains as defined by Rutherford. • Electron states are described as orbitals, not orbits. • Orbitals are regions around a nucleus where electrons have a probability of being found. The precise location and speed of electrons within orbitals cannot be accurately determined (see firefly example on pg 289). • Schrodinger’s wave equations describe orbitals of various shapes – such as spherical or dumbbell (see pg 291 in text).

30. Orbitals are also described as electron cloud regions. • Electron clouds are visual models that map the possible location of electrons in an atom. • The edge of an orbital or cloud is “fuzzy,” meaning it does not have an exact size. • Electron clouds are denser closer to the nucleus where electrons are more likely to be found. Electrons can only be present in certain regions around the nucleus, not just anywhere. not here here

31. Location of an electron depends upon how much energy the electron has. • Electrons closer to the nucleus have less energy than those further away. • Energy states of electrons correspond to orbitals with different shapes: s, p, d, f (see below).

32. Summary of Models • Greeks • Atoms are indivisible, infinite, moving, join together • Dalton • Atoms are indivisible, different for each element • Form compounds by joining together in fixed ratios • Thompson • Atoms are divisible. Consist of negatively-charged electrons and a particles with a positive charge. • Plum-Pudding model

33. Rutherford • Atoms consist of positively-charged particles called protons located in a tiny, very dense nucleus. Neutral particles called neutrons are also located within the nucleus for stability. • Negatively-charged electrons are arranged in orbits around the nucleus. • Atoms are mostly empty space. • Bohr • Retains Rutherford’s basic model except that electrons are restricted to certain orbits around nucleus. (Like planets around the sun.) • Orbits of electrons are quantized energy levels. • At best this model could only be applied to hydrogen. • Wave-Mechanical • Electrons occupy orbitals which are regions of space around the nucleus where they are likely to be found. They do not orbit the nucleus like planets around the sun. • This model worked for atoms of any element.

34. Some links to videos on the web • Atomic timeline video: http://www.youtube.com/watch?v=NSAgLvKOPLQ • Lots of interesting videos! Pick the ones that interest you. http://www.youtube.com/watch?v=bw5TE5o7JtE&list=PLD6A43875B9DEC57E&index=1