The History of the Atom

1. The History of the Atom

2. Presentation Outline • Introduction • Ancient Atomic theory • The Modern Atomic Theory • Rutherford's Experiment • Bohr’s Model • Quantum Theory of the Atom

3. Timeline of Atomic Models ?

4. Introduction Atomic theory first originated with Greek philosophers around 2500 years ago. This basic theory remained unchanged until the 19th century when it first became possible to test the theory with more sophisticated experiments. As science has rapidly advanced over the past few centuries the atomic theory has been refined in accordance with the accepted scientific principles and theories of the time. Advances in technology and theory have allowed experiments to be designed that enable us to probe matter to microscopic scales.

5. Ancient Atomic Theory   The atomic theory of matter was first proposed by Leucippus, a Greek philosopher who lived at around 400BC. At this time the Greeks were trying to understand the way matter is made. According to Anassagora, another Greek philospher ,it is possible to subdivide matter in smaller and smaller parts, and he proposed that this process can be continued with no limit. But according to Leucippus, eventually you arrive at small particles which can not be further subdivided. Leucippus called these indivisible particles atoms (from the Greek word atomos, meaning “indivisible”). Leucippus

6. Ancient Atomic Theory continued.. Leucippus's atomic theory was further developed by his disciple, Democritus. Democritus suggested the atomic theory, explaining that all things are "composed of minute, invisible, indestructible particles of pure matter which move about eternally in infinite empty.". If a sample of a pure element was divided into smaller and smaller parts, eventually a point would be reached at which no further cutting would be possible—this was the atom of that element. Democritus According to the ancient Greeks, atoms were all made of the same basic material, but atoms of different elements had different sizes and shapes. The sizes, shapes, and arrangements of a material’s atoms determined the material’s properties. It was believed that there were four elements that all things were mare from; Earth, Air, Fire and Water.

7. Dalton’s Atomic Theory For centuries scientists did not have the methods or technology to test their theories about the basic structure of matter, so people accepted the ancient Greek view. In the earliest days of chemistry the chief ‘chemical’ occupations were held by the alchemists. During the middle ages, alchemists tried to transform various metals into gold. They also produced metals from their ores, made glasses and enamels, and dyed fabrics. However, there was no understanding of what happens in these processes. It wasn’t until 1803, however, when an English school teacher, John Dalton, was able to propose the very first atomic theory. This theory was the result of much experimentation into the nature of matter. It paved the way towards a deeper understanding of what chemicals are and what happens when they react. John Dalton

8. The Laws John Dalton based his atomic theory upon the following laws. • Law of Conservation of Mass (1782) (demo time!) • Law of Definite Proportions (1797) • Law of Multiple Proportions (1803)

9. Law of Conservation of Mass

10. The Law of Constant Composition • The law of definite proportions states that a chemical compound contains the same elements in the same proportions, regardless of the size of the sample. “Any given compound always consists of the same atoms and the same ratio of atoms. For example, water always consists of oxygen and hydrogen atoms, and it is always 89 percent oxygen by mass and 11 percent hydrogen by mass” H2 + O2 = H2 O 89 gm + 11 gm = 100 gm 45.0 gm + 5.0 gm = 50 gm

11. The Law of Multiple Proportions • The law of multiple proportions states that if two or more different compounds consist of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small, whole numbers. “If two elements combine to form more than one compound,the masses of one of the elements that can combine with a given mass of the other element are related by factors of small wholenumbers” • For example, water has an oxygen-to-hydrogen mass ratio of 8:1. Hydrogen peroxide, another compound consisting of oxygen and hydrogen, has an oxygen-to-hydrogen mass ratio of 16:1. The ratio of these two ratios gives a small whole number.

12. The Dalton’s Atomic Theorycontinued… 1. All elements are composed of atoms, which are indivisible and indestructible particles. 2. All atoms of the same element are exactly alike; in particular, they all have the same mass. 3. Atoms of different elements are different; in particular, they have different masses. 4. Compounds are formed by the joining of atoms of two or more elements. They are joined in a definite whole-number ratio, such as 1 to 1, 2 to 1, 3 to 2, etc. 5. Chemical reactions involve the rearrangement of atoms to make new compounds. Although the first two theories , Dalton contributed greatly to the advances of atomic theory, and would greatly influence J.J. Thompson in his own discoveries, they would later be proved incorrect.

13. Dalton’s Atomic Theory • What is wrong with Dalton’s Theory? • The atom is not the smallest particle of matter. I bet you already know that there are three sub atomic particles that make up the atom. Next we will learn about their discovery. • Also all carbon atoms are not exactly the same. We will learn about this also.

14. Problems with Dalton’s Atomic Theory? 1. matter is composed, indivisible particles Atoms Can Be Divided, but only in a nuclear reaction 2. all atoms of a particular element are identical Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)! 3. different elements have different atoms YES! 4. atoms combine in certain whole-number ratios YES! Called the Law of Definite Proportions 5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements. Yes, except for nuclear reactions that can change atoms of one element to a different element

15. Expanding the Dalton’s Atomic Theory J.J. Thompson is the person who is credited for discovering the electron. Thompson created a tube that had a positively charged anode on one side and a negatively charged cathode on the other side. Thompson then applied a magnet to the middle of the tube and discovered that particles were emanating towards the positive magnetic field. From this, Thompson concluded that particles were negatively charged and came from inside the atom, that is, the atom is not indivisible and there are smaller particles than the atom. J.J. Thompson Thompson then created the Plum Pudding model, which suggested that electrons were randomly placed throughout the atom in a positively charged “pudding”. The “Plum Pudding”

16. J.J. Thomson’s Experiment J.J. Thomson was experimenting with electricity using a gas discharge tube like that shown below. No air was present in the tube. Thomson discovered that an electric current flowed from the cathode to the anode by means of a ray of some kind. These rays later became known as electrons. Your instructor may give you more information about Thomson’s experiments.

17. Take the link below to see the cathode ray tube experiment • http://www.chem.uiuc.edu/clcwebsite/cathode.html • We can see this happen as a real demo in class!

18. FYI (For Your Information) William Crookes was the first to discover these negative rays in 1879. Thomson was able to conduct experiments that proved these rays were actually negatively charged particles. Thomson is generally given credit for the discovery of the electron. But…. The plum pudding model did not explain Rutherford's observations that showed that a positive charge was not found throughout the atom and electrons were not randomly distributed either.

19. Rutherford's Experiment In 1911 British scientist Ernest Rutherford set out to test Thomson’s proposal by firing a beam of charged particles at atoms. Alpha particles are heavy particles with twice the positive charge of a proton. Alpha particles are now known to be the nuclei of helium atoms, which contain two protons and two neutrons. Ernest Rutherford's experiment was to emit alpha particles towards a thin gold sheet. Rutherford would then determine where the deflections of the alpha particles would go, and therefore be able to theorize what kind of placement protons and electrons had. Ernest Rutherford

20. Rutherford's Experiment continued…

21. Results of foil experiment if Plum Pudding model had been correct.

22. What Actually Happened

23. Take the links below to see videos of the Gold Foil Experiment • http://www.youtube.com/watch?v=Q8RuO2ekNGw&feature=related • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf

24. Rutherford's Experiment continued… Rutherford observed that most of the alpha particles went strait through the foil. However a large proportion were deflected through small angles and some (though very few) deflected straight back. Rutherford then theorized that there was something called a nucleus, which contained a high density of positively charged particles. Rutherford was able to say there was a nucleus because alpha particles that deflected right back must have hit something more massive and with a strong positive charge. This led Rutherford to propose a very different model for the atom. Instead of supposing that the positive charge and mass were spread throughout the volume of the atom, he theorized that it was concentrated in the center of the atom. Rutherford called this concentrated region of electric charge the nucleus of the atom. What was wrong with his theory?

25. Electron ( - charge) Nucleus Protons (+ charge) Great distance The Model and It’s Flaws Click until buttons appear. Flaws 1. Protons only accounted for half the mass of the atom. 2. Electrons should attract to the nucleus and collapse the atom. Opposite charges attract!

26. Bohr’s Model Danish physicist Niels Bohr used new knowledge about the radiation emitted from atoms to develop a model of the atom significantly different from Rutherford’s model. Scientists of the 19th century discovered that when an electrical discharge passes through a small quantity of a gas in a glass tube, the atoms in the gas emit light. This radiation occurs only at certain discrete wavelengths, and different elements and compounds emit different wavelengths.

27. He concluded that because atoms emit light only at discrete wavelengths (certain colors), electrons could only orbit at certain designated radii, and light could be emitted only when an electron jumped from one of these designated orbits to another. This fixed amount of energy kept the electron from collapsing into the nucleus.

28. Bohr’s Quantum Model of the Atom Electrons can only be found orbiting the nucleus at fixed distances. When the atom absorbs energy the electron can jump to a higher path. When the atom emits (gives off) light energy the electron is jumping down to an orbit closer to the nucleus.

29. Quantum Theory of the Atom To make his theory work, Bohr had to propose special rules that violated the rules of classical physics. He concluded that, on the atomic scale, certain preferred states of motion were especially stable. In these states of motion an orbiting electron (contrary to the laws of electromagnetism) would not radiate energy. The quantum mechanical view of atomic structure is that the nucleus is at the center of the atom and provides the electrical attraction that binds the electrons to the atom. Contrary to Bohr’s theory, however, the electrons do not circulate in definite planet-like orbits. Due to the wavelike character of electrons and provides the framework for viewing the electrons as fuzzy clouds of negative charge. We will study this in detail at another time……

30. Modern Atomic Theory • 1. Atoms are not indivisible. They are made up of protons, electrons, and neutrons. • 2. Atoms of the same element can, and do have different masses. These atoms are called isotopes. Isotopes have the same number of protons but a different number of neutrons. • 3. Atoms of different elements are different. They differ in their physical and chemical properties. • 4. Compounds are formed by the joining of atoms of two or more elements. They are joined in a definite whole-number ratio, such as 1 to 1, 2 to 1, 3 to 2, etc. • 5. Chemical reactions involve the rearrangement of atoms to make new compounds.

31. THE END

32. Continued… Electrons still have assigned states of motion, but these states of motion do not correspond to fixed orbits. Instead, they tell us something about the geometry of the electron cloud—its size and shape and whether it is spherical or bunched in lobes like a figure eight. Physicists called these states of motion orbitals. The way electrons fill up orbitals determines the number of electrons that end up in the atom’s valence shell. This in turn determines an atom’s chemical and physical properties, such as how it reacts with other atoms and how well it conducts electricity. It is through today’s understanding of the nature of matter that advancement of technology is able to continue at an increasing rate.

33. Bohr developed a theory by which he could predict the same wavelengths scientists had measured radiating from atoms with a single electron. He concluded that because atoms emit light only at discrete wavelengths, electrons could only orbit at certain designated radii, and light could be emitted only when an electron jumped from one of these designated orbits to another. This fixed amount of energy kept the electron form collapsing into the nucleus.

34. References • encarta.msn.com/encyclopedia • www.funsci.com/fun3_en/democritus/democritus • www.aare.edu.au • www.absoluteastronomy.com/encyclopedia • encarta.msn.com/encyclopedia • www.funsci.com/fun3_en/democritus/democritus • www.aare.edu.au • www.absoluteastronomy.com/encyclopedia • encarta.msn.com/encyclopedia • www.funsci.com/fun3_en/democritus/democritus • www.aare.edu.au • www.absoluteastronomy.com/encyclopedia