Development of the Modern Atomic Theory

1. Chapter 3: Atoms: building Blocks of Matter Development of the Modern Atomic Theory • In 1782, a French chemist, Antoine Lavoisier (1743-1794), made measurements of chemical change in a sealed container. • He observed that the mass of reactants in the container before a chemical reaction was equal to the mass of the products after the reaction.

2. Development of the Modern Atomic Theory • Lavoisier concluded that when a chemical reaction occurs, mass is neither created nor destroyed but only changed. • Lavoisier’s conclusion became known as the law of conservation of mass. Click box to view movie clip.

3. Development of the Modern Atomic Theory • In 1799, another French chemist, Joseph Proust, observed that the composition of water is always 11 percent hydrogen and 89 percent oxygen by mass. • Regardless of the source of the water, it always contains these same percentages of hydrogen and oxygen. H H O

4. Development of the Modern Atomic Theory • Proust studied many other compounds and observed that the elements that composed the compounds were always in a certain proportion by mass. This principle is now referred to as the law of definite proportions.

5. Dalton’s Atomic Theory • John Dalton (1766-1844), an English schoolteacher and chemist, studied the results of experiments by Lavoisier, Proust, and many other scientists.

6. Dalton’s Atomic Theory • Dalton proposed his atomic theory of matter in 1803. • Although his theory has been modified slightly to accommodate new discoveries, Dalton’s theory was so insightful that it has remained essentially intact up to the present time.

7. Dalton’s Atomic Theory • The following statements are the main points of Dalton’s atomic theory. 1. All matter is made up of atoms. 2. Atoms are indestructible and cannot be divided into smaller particles. (Atoms are indivisible.) 3. All atoms of one element are exactly alike, but are different from atoms of other elements.

8. The Electron • Because of Dalton’s atomic theory, most scientists in the 1800s believed that the atom was like a tiny solid ball that could not be broken up into parts. • In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model was not accurate. • Thomson’s experiments used a vacuum tube.

9. The Electron • A vacuum tube has had all gases pumped out of it. • At each end of the tube is a metal piece called an electrode, which is connected through the glass to a metal terminal outside the tube. • These electrodes become electrically charged when they are connected to a high-voltage electrical source.

10. Cathode-Ray Tube • When the electrodes are charged, rays travel in the tube from the negative electrode, which is the cathode, to the positive electrode, the anode. • Because these rays originate at the cathode, they are called cathode rays.

11. Cathode-Ray Tube • Thomson found that the rays bent toward a positively charged plate and away from a negatively charged plate. • He knew that objects with like charges repel each other, and objects with unlike charges attract each other.

12. Cathode-Ray Tube • Thomson concluded that cathode rays are made up of invisible, negatively charged particles referred to as electrons. • These electrons had to come from the matter (atoms) of the negative electrode.

13. Cathode-Ray Tube • From Thomson’s experiments, scientists had to conclude that atoms were not just neutral spheres, but somehow were composed of electrically charged particles. • Reason should tell you that there must be a lot more to the atom than electrons. • Matter is not negatively charged, so atoms can’t be negatively charged either.

14. Cathode-Ray Tube • If atoms contained extremely light, negatively charged particles, then they must also contain positively charged particles—probably with a much greater mass than electrons.

15. Protons • In 1886, scientists discovered that a cathode-ray tube emitted rays not only from the cathode but also from the positively charged anode. • These rays travel in a direction opposite to that of cathode rays.

16. Protons • Like cathode rays, they are deflected by electrical and magnetic fields, but in directions opposite to the way cathode rays are deflected. • Thomson was able to show that these rays had a positive electrical charge. • Years later, scientists determined that the rays were composed of positively charged subatomic particles calledprotons.

17. However, in 1910, Thomson discovered that neon consisted of atoms of two different masses. Protons • At this point, it seemed that atoms were made up of equal numbers of electrons and protons.

18. Protons • Atoms of an element that are chemically alike but differ in mass are called isotopes of the element. • Today, chemists know that neon consists of three naturally occurring isotopes. • The third was too scarce for Thomson to detect.

19. Neutrons • Because of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass. • Calculations showed that such a particle should have a mass equal to that of a proton but no electrical charge. • The existence of this neutral particle, called a neutron, was confirmed in the early 1930s.

20. Rutherford’s Gold Foil Experiment • In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important experiments that revealed an arrangement far different from the cookie-dough model of the atom.

21. Rutherford’s Gold Foil Experiment • The experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively charged subatomic particles through a small hole.

22. Rutherford’s Gold Foil Experiment • Today, we know that the particles of the beam consisted of clusters containing two protons and two neutrons and are called alpha particles. • The sheet of gold foil was surrounded by a screen coated with zinc sulfide, which glows when struck by the positively charged particles of the beam.

23. The Gold Foil Experiment

24. The Nuclear Model of the Atom • To explain the results of the experiment, Rutherford’s team proposed a new model of the atom. • Because most of the particles passed through the foil, they concluded that the atom is nearly all empty space.

25. The Nuclear Model of the Atom • Because so few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a nucleus. Thompson 1906 Rutherford 1913 Bohr 1924

26. The Nuclear Model of the Atom • The new model of the atom as pictured by Rutherford’s group in 1911 is shown below.

27. Atomic Numbers • The atomicnumber of an element is the number of protons in the nucleus of an atom of that element. • It is the number of protons that determines the identity of an element, as well as many of its chemical and physical properties.

28. Atomic Numbers • Because atoms have no overall electrical charge, an atom must have as many electrons as there are protons in its nucleus. • Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element.

29. Masses • The mass of a neutron is almost the same as the mass of a proton. • The sum of the protons and neutrons in the nucleus is the mass number of that particular atom.

30. Masses • Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

31. Atomic Mass • In order to have a simpler way of comparing the masses of individual atoms, chemists have devised a different unit of mass called an atomic mass unit, which is given the symbol u. • An atom of the carbon-12 isotope contains six protons and six neutrons and has a mass number of 12.

32. Atomic Mass • Chemists have defined the carbon-12 atom as having a mass of 12 atomic mass units. • Therefore, 1 u = 1/12 the mass of a carbon-12 atom. • 1 u is approximately the mass of a single proton or neutron.

33. Information in the Periodic Table • The number at the bottom of each box is the average atomic mass of that element. • This number is the weighted average mass of all the naturally occurring isotopes of that element.

34. Question 1 How does the atomic number of an element differ from the element’s mass number? Answer The atomic number of an element is the number of protons in the nucleus. The mass number is the sum of the number of protons and neutrons.

35. Calculating Atomic Mass

36. Calculating Atomic Mass • Copper exists as a mixture of two isotopes. • The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms. • The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

37. Calculating Atomic Mass • The atomic mass of Cu-63 is 62.930 amu, and the atomic mass of Cu-65 is 64.928 amu. • Use the data above to compute the atomic mass of copper.

38. Calculating Atomic Mass • First, calculate the contribution of each isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

39. Calculating Atomic Mass • The average atomic mass of the element is the sum of the mass contributions of each isotope.

40. Assessment Questions Question 3 Calculate the atomic mass of germanium. Answer 72.59 amu