Atomic Theory: Why do we believe in atoms?

1. ID1004: Approaches to Reality Atomic Theory:Why do we believe in atoms? Dr John Mitchell

2. For Years People Wondered …

3. What is the World Made of?

4. What is Matter?

5. Is Water Made of the Same Stuff as Rock?

6. Is Matter Continuous?

7. Is Matter Continuous?

8. Is Matter Continuous?

9. Is Matter Continuous? Indivisible, no microscopic structure.

10. Is Matter Discrete?

11. Is Matter Discrete?

12. Is Matter Discrete? Composed of microscopic particles, probably with void between them.

13. Matter: Continuous or Discrete? Related questions: What makes rock rocky? What makes water watery? Do materials have some essence, independent of objects formed from them?

14. Ancient Times: Philosophy, not Experiment In ancient times, the question of the continuity or discreteness of matter was not amenable to experimental investigation … … it was a matter for philosophical speculation, though the Greek expertise in geometry and simple observations of the world could also be relevant.

15. Leucippus and his pupil Democritus Leucippus Democritus Time: approximately mid- to-late 5th century BC. Place: Greece.

16. Leucippus &Democritus’s Beliefs Matter is composed of indivisible, indestructible atoms, in continuous motion through an empty void. There are many different kinds of atom, varying in shape and size. Atoms can be connected together. Shape, size and connections relate to the observed properties of materials.

17. Leucippus &Democritus’s Beliefs Salt atoms: pointed, sharp (presumably relates to taste). Water atoms: smooth (flow freely). Air atoms: light and rapidly moving. Iron atoms: tightly bound together by connections (hence a strong solid material).

18. Reasonable Reasons for their Beliefs Movement requires a void. Materials have an indivisible ‘essence’. A vase is still clay, even if it is broken.

19. Reasonable Reasons for their Beliefs Not proofs! Movement requires a void. Movement was surmised, not observed. Materials have an indivisible ‘essence’. A vase is still clay, even if it is broken. Logically possible for various materials all to be continuous, but made of ‘different stuff’.

20. Plato’s Perfect Solids as Atoms Plato The cube is the only Platonic solid to tessellate without gaps – appropriate for solidity of earth.

21. The Rise of Experimental Science

22. The Rise of Experimental Science By the late 18th century, alchemy was giving way to chemistry. The existence of chemical elements was recognised. Numerous elements were being discovered (about 20 new ones in the 18th century, at least 50 more in 19th century). This allowed compounds to be analysed to obtain compositions in terms of component elements.

23. The Rise of Experimental Science Some key principles emerged that gave an experimental basis for discussions of the nature of matter. Atomic theory was to gain ground gradually, rather than through one dramatic discovery, because the ideas of atoms and molecules would allow chemists to make sense of the chemical world.

24. Elements Combine in Fixed Proportions The key idea underlying the acceptance of atomic theory is that elements combine in fixed and simple proportions to form compounds. This is easily explained if we believe in atoms, much harder to understand otherwise. However, because different atoms have different masses, correctly identifying these ratios was not trivial.

25. 1. Law of Conservation of Mass Antoine Lavoisier (1743-1794) found that the total mass of reactants is equal to the total mass of products in a chemical reaction.

26. Lavoisier had worked as a tax collector and as an aristocrat found himself on the wrong side in the French revolution.

27. Lavoisier had worked as a tax collector and as an aristocrat found himself on the wrong side in the French revolution.

28. 2. Law of Definite Proportions Joseph Proust (1754-1826) stated in 1806 that every sample of a given chemical compound has an identical elemental composition, in terms of percentages by mass. Jons Jacob Berzelius (1779-1848) later confirmed and popularised this finding. Example: All samples of mercury oxide are 92.6% mercury by mass and 7.4% oxygen.

29. 3. Law of Multiple Proportions Manchester chemist John Dalton (1766-1844) stated in 1803 that whenever elements combine to form more than one compound, the relative amounts of two given elements in each compound will be in ratios of small whole numbers. Example: combine tin and oxygen. One compound has 88.1% tin and 11.9% oxygen by mass (ratio 7.4 : 1). The other compound has 78.8% tin and 21.2% oxygen (ratio 3.7 : 1). Ratio of ratios = 7.4 : 3.7 = 2 : 1, a simple whole number ratio. Put another way, 100g of tin combines either with 13.5g or 27g of oxygen – a factor of two.

30. Dalton’s Atomic Theory John Dalton saw that these observed laws made sense if the elements were formed of atoms which could combine in simple ratios to form compounds.

31. Dalton’s Atomic Theory Dalton believed that … Elements consist of tiny indivisible atoms. All atoms of the same element are identical. Atoms of different elements have different atomic weights. Atoms of different elements combine to form compounds of fixed stoichiometry.

32. Dalton’s Atomic Weights Dalton also believed that stoichiometry followed a Rule of Greatest Simplicity … … which meant that he assumed water to be HO and ammonia to be HN. Since the correct stoichiometry of compounds was unknown, early attempts to compile atomic weights were error-prone. Incorrect assumptions about atomic weights no doubt delayed the acceptance of atomic theory by impairing its ability to explain the observations.

33. 4. Law of Combining Volumes Joseph Gay Lussac (1778-1850) found in 1808 that the ratios of the volumes of each reactant and product gas in a chemical reaction are ratios of simple small whole numbers. 3 litres of hydrogen gas + 1 litre of nitrogen gas  2 litres of ammonia gas. 2 litres of hydrogen gas + 1 litre of oxygen gas  2 litres of water vapour.

34. 5. Avogadro’s Law Amedeo Avogadro (1776-1856) made two significant advances in atomic theory. Firstly, he clarified the distinction between atoms and molecules. Oxygen, hydrogen and nitrogen, though elements, form diatomic molecules (O2, H2, N2). Secondly, around 1811, he interpreted the Law of Combining Volumes in terms of atomic, or molecular, theory. Equal volumes of any gas, at the same temperature and pressure, contain equal numbers of molecules.

35. The Composition of Water 2 litres of hydrogen gas + 1 litre of oxygen gas  2 litres of water vapour. How can we explain this ratio using Avogadro’s Law? 2H2 + O2  2H2O The formula of water is H2O, not HO as previously assumed by Dalton. Thus, the atomic weight of oxygen is 16 times that of hydrogen (not 8 times).

36. Cannizzaro and the Karlsruhe Congress StanislaoCannizzaro (1826-1910) gave a paper at the 1860 Karlsruhe congress, in which he advocated Avogadro’s approach. Cannizzaro is credited with persuading the somewhat sceptical chemical community to adopt a recognisably modern view of atoms, molecules and atomic weights. This included the reformed atomic weights H=1, C=12, O=16.

37. Structure in Chemistry Over the second half of the 19th century it became apparent that atomic and molecular theory could both make sense of the disparate body of knowledge about substances and reactions that chemistry had become and form the basis for reliable predictions. The key idea was that molecules each had a defined structure based on connections between atoms. The atoms themselves each had a valency, a set number of connections they could make.

38. Couper and Molecular Structure Scottish chemist Archibald Scott Couper (1831-1892) was a pioneer of the concept of valency and of the use of simple structural formulae to represent molecules. Because of the (incorrect) atomic weights in use at the time, these structures have too many oxygen atoms. Sadly, Couper’s career and health were broken by a spectacular row with his Professor when he was beaten into print by …

39. August Kekulé August Kekulé (1829-1896), who is most famous for proposing the ring structure of benzene in 1865. Equally important was his 1857 proposal of tetravalent carbon.

40. Crum Brown & Molecular Structural Formulae Edinburgh chemist Alexander Crum Brown (1838-1922) developed the diagrammatic representation of molecules from 1864, his diagrams clearly showing single and double bonded connections between atoms. He also discovered the double bond in ethene. His diagrams were topological and deliberately two dimensional: “I do not mean to indicate the physical, but merely the chemical position of the atoms.”

41. Structural Formulae

42. The Periodic Table Dmitri Mendeleev (1834-1907), possibly the greatest of all chemists, had created his periodic table on the basis of periodic chemical properties and atomic weights that increased down a group and, broadly speaking, increased along a period.

43. The Periodic Table Dmitri Mendeleev (1834-1907), possibly the greatest of all chemists, had created his periodic table on the basis of periodic chemical properties and atomic weights that increased down a group and, broadly speaking, increased along a period.

44. Arguments over Atomic Theory Although most chemists by then accepted the reality of atoms, there was still significant opposition as late as 1900. Amongst notable opponents was physical chemist Wilhelm Ostwald (1853-1932, Nobel Chemistry Laureate 1909).

45. A Useful Model? Ludwig Boltzmann (1844-1906) developed the kinetic theory of gases which seemed only to make sense if atoms or molecules really existed. He was a strong advocate of atomic theory, but prepared to accept a compromise whereby atomic theory was agreed to be a useful model, but no absolute belief in its reality was required. Frustration at the lack of acceptance of his theories, as well as probable psychiatric illness, may well have contributed to his suicide.

46. Atomic Structure: A Surprise Ernest Rutherford’s 1911 paper interpreting the results of Geiger and Marsden (firing alpha particles at gold foil) brought a surprise … … almost all the mass of the atom was concentrated in a tiny region in the centre, the nucleus. Ernest Rutherford, NZ-British physicist, 1871-1937

47. Atomic Number as an Observable Manchester physicist Henry Moseley (1887-1915) found that the frequencies of the X-rays emitted by elements depended directly on their atomic numbers. This meant that atomic number was an experimentally observable quantity. Moseley saw this as confirmation of van den Broek’s hypothesis that the atomic number was the positive charge of the nucleus. Antonius van den Broek, Dutch lawyer & physicist (1870-1926)

48. Atomic Number as an Observable Moseley’s observations also justified the need to leave gaps in the periodic table, as Mendeleev had famously done, which would be necessary if no known element had a given value of atomic number (as then was true for 43, 61, 72, and 75; Tc, Pm, Hf, Re). Moseley also showed that the ordering of elements by atomic number was sometimes different from that by their atomic weights (examples are Co, Ni and Ar, K).

49. Moseley’s death at Gallipoli in 1915 is widely credited with changing British policy towards sending outstanding scientists to the front line. “In view of what he [Moseley] might still have accomplished ... his death might well have been the most costly single death of the War to mankind generally.” (Isaac Asimov)

50. How Big is the Atom? Michael Faraday (1791-1867) Robert Millikan (1868-1953) Avogadro’s number is the number of atoms in 12g of carbon, or molecules in 18g of water, or formula units in 58.5g of NaCletc. The electric charge of a mole (Avogadro’s number) of electrons was known to be equal to a constant found by Michael Faraday. When Robert Millikan measured the charge on the electron in 1910, Avogadro’s number was now known: F = 96,500 C mol-1 e = 1.610-19C NA= F/e = 6.02 1023 mol-1