2. Atomic Structure Chapter 5

3. The Atomists: The first atomic theory 460 BCE: Greek Democritus suggested that matter is “ composed of minute, invisible, indestructible particles of pure matter which move about eternally in infinite empty” http://www.winneconne.k12.wi.us/middle_school/7th%20Grade/LENZ/history_of_atomic_theory.htm

4. Development of 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.

5. Law of Conservation of Mass 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.

6. Dalton’s Atomic Theory John Dalton (1766- 1844), an English schoolteacher and chemist, studied the results of experiments by Lavoisier and many other scientists. Dalton proposed his atomic theory of matter in 1803.

7. Dalton’s Atomic Theory (The main points) 1. All matter is made up of indivisible particles called atoms. 2. All atoms of one element are exactly alike, but are different from atoms of other elements. 3.Atoms can mix in simple whole number ratios to form compounds. 4.Chemical reactions happen when atoms are separated, joined, or rearranged.

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.

9. Cathode-Ray Tube Thomson knew that objects with like charges repel each other, and objects with unlike charges attract each other

10. The Electron Thomson named the negative particles found in his experiment electrons. Electrons (e - ): –Are the smallest subatomic particles –Have a charge of -1

11. Protons Atoms aren’t negative, so there must be something positive in the atom to balance the electrons. Protons (p + ): –Positively charged subatomic particles (+1) –Mass 1 proton = mass 1840 electrons

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

13. Neutrons Neutrons (n 0 ): Neutral charge Same mass as a proton Because of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass.

14. 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.

15. The Gold Foil Experiment

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

17. 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. Nucleus: Contains most of the mass of the atom (protons and neutrons) Is very small compared to the rest of the atom

18. You know you want to know how many of these particles are in an atom, right? Atomic Number : # of protons in the atom Mass number: # of protons + # neutrons

19. How to get the info from the periodic table H 1 1.0013 Atomic Number: Number of protons in the nucleus of the atom.

20. In this chapter you will fill out lots of tables… NameProtonsElectronsNeutrons Helium-4 Carbon-12 Carbon-13

21. How many subatomic particles? # protons = atomic number (always) For a NEUTRAL atom # electrons=# protons (this will be a bit different in upcoming chapters…) # neutrons = mass number - # protons (always)

22. In this chapter you will fill out lots of tables… NameProtonsElectronsNeutrons Helium-4 Carbon-12 Carbon-13 Carbon-12 and Carbon-13 are isotopes

23. Isotopes Atoms of an element that are chemically alike but differ in mass are called isotopes of the element. So why the difference in mass?

24. So what is the bottom number for? H 1 1.0013 Atomic Number: Number of protons in the nucleus of the atom.

25. The atomic mass unit (amu) Atoms don’t have much mass! (10 -23 g) –This is small, and a bit hard to work with Use carbon-12 isotope as a ref. and make 12 carbon atom= 1 amu On the periodic table the value is a weighted average of all isotopes of the element – that is why it isn’t a whole number (includes carbon-12, carbon-13, etc).

26. The Periodic Table Chemists kept discovering new elements and decided they needed a way to organize them. Dmitri Mendeleev constructed the 1 st periodic table –Grouped elements according to similar chemical properties –He was so smart he left spaces for elements that would be discovered later on.

27. The modern periodic table Henry Moseley tweaked the table and ordered by atomic number rather than atomic mass.

28. Period Group or family Alkali Metals Alkaline Earth Metals Transition Metals Inner Transition Metals Nonmetals Metals Halogens Noble Gases