Chapter 4 “ Atomic Structure ”

1. Chapter 4“Atomic Structure”

2. Introduction to the Atom and Atomic Models

3. Democritus believed all things consisted of tiny indivisible units. He called these tiny units he called atomos. The Greek word for “can not be cut” or “indivisible” Ancient philosopher: Father of the Atom Democritus400 BC

4. John Dalton (1799)  Developed what is considered to be the 1st Atomic Theory  Was born into a modest Quaker family in England  Began lecturing in public at the age of 12

5. Dalton’s Model (1799) Dalton's model was that the atoms were tiny, indivisible, indestructible particles and that each one had a certain mass, size, and chemical behavior that was determined by what kind of element they were.

6. Dalton’s Model Dalton’s model of the atom was similar to a tiny billiard ball. Dalton’s model of the atom was solid and had no internal structure.

7. Dalton’s Atomic Theory • elements consisted of tiny particles called atoms. • all atoms of an element are identical • atoms of each element are different from one another; they have different masses. • compounds consisted of atoms of different elements combined together. • chemical reactions involved the rearrangement of combinations of those atoms.

8. Flaws in Dalton’s Model • Dalton’s falsely believed that the atom was the most fundamental particle. • We now know the atom is made up of even smaller particles we call the proton, neutron and electron. • Dalton’s theory could also not account for the formation of ions (charged particles)

9. Daltons Atomic Model Summary • Called: Billiard Ball Model • Could account for • Atoms of different atomic masses • Elements were tiny particles • Could NOT account for • Though atom was smallest particle • Did not have an internal structure • The formation of charged particles

10. John J. Thomson (1897) Discovered the electron using the Cathod Ray Tube (CRT) Thomson found that the beam of charge in the CRT was attracted to the positive end of a magnet and repelled by the negative end.

11. Thomson’s Hypothesis Concluded that the cathode beam was a stream of negative particles (electrons). He tested several cathode materials and found that all of them produced the same result. He also found that the charge to mass ratio was the same for all electrons regardless of the material used in the cathode or the gas in the tube. Thomson concluded that electrons must be part of all atoms.

12. Thomson’s atomic model Called the “plum-pudding” it was the most popular and most wildly accepted model of the time.

13. Thompsons atomic model could account for….. the atom having an internal structure Light given off by atoms Atom with different atomic masses

14. Thompsons atomic model could NOT account for….. Empty space (had atom filled with positive pudding) Formation of ions

15. Gold Foil Experiment Conducted by students of Rutherford. Proved that all atoms had a tiny, positively charged center. Confirmed that atom’s were mostly empty space.

16. Rutherford ~ early 1900s • α-particle interaction with matter studied in gold foil experiment

18. Rutherford's Nuclear Model • 1. The atom contains a tiny dense center • the volume is about 1/10 trillionth the volume of the atom • 2. The nucleus is essentially the entire mass of the atom • 3. The nucleus is positively charged • the amount of positive charge of the nucleus balances the negative charge of the electrons • 4. The electrons move around in the empty space of the atom surrounding the nucleus

19. Rutherford’s atomic model (1911) • Could account for: • Empty space • Ions • Internal structure • Light given off when heated to high temperature. • Could not account for: • Stability

20. Philipp Lenard (1903) Aluminum foil experiment Lenard found that a beam of electrons was able to pass through a sheet of Al foil with almost no deflection. Lenard correctly concluded that the majority of an atom’s volume is empty space.

21. Lenard’s atomic model Lenard’s model was composed of dynamids. Lenard calculated the size of a dynamid based on his experimental results and found it to be 1 billionth (1/1,000,000,000) the size of the atom.

22. Lenard’s model could account for: Different atomic masses (based on the number of “dynamids”). The internal structure of an atom The fact that most of the atom was empty space.

23. Flaws in Lenard’s model Formation of Ions (gaining or losing charge) The light given off by materials when heated to a high temperature.

24. Hantaro Nagaoka • First to present an atomic model close to the presently accepted model. • He came up with his model in 1903.

25. Nagoka’s Model • Nagoka’s model of the atom was unstable. • According to the laws of planetary motion, the atom would collapse over time.

26. Planetary model • Planetary model used to explain electrons moving around the tiny, but dense nucleus • Nucleus contains • Protons- existence proposed in 1900s • Neutrons- existence proposed in 1930s

27. Successes of Nagoka’s model Atoms were able to give off electrons to form ions. Accounted for the experimental fact that atom’s were mostly empty space Explained different atomic weights. Explained the light given off when heated to high temperatures.

28. Bohr • Questioned ‘planetary model’ of atom • Electrons located in specific levels from nucleus (discontinuous model) • Proposed electron cloud model based on evidence collected with H emission spectra

29. Bohr’s Atomic Model (1913) Bohr was a student of Rutherford. Improved Rutherford’s model by proposing electrons are found only in specific fixed orbits. These orbits have fixed levels of energy This explained how electrons could give off light (gain or lose energy)

30. BOHR MODEL 8e- 8e- Nucleus (p+ & n0) 2e- Electrons are placed in energy levels surrounding the nucleus

31. Bohr’s Atomic Model • Could account for • Internal structure • Atoms of different masses • Atom being mostly empty space • Light given off • Formation of positive ions • Flaws • Only really worked for Hydrogen

32. Chadwick (1932) Discovered the neutron by bombarding Be with beta radiation. Nuclear fission released a neutron. chart

33. Review • Describe each of the 6 different atomic models. Give the • Scientist Name • Name of model • What they could account for • What they could not account for (flaws)

34. Subatomic particle summary

35. Subatomic Particles Actual mass (g) Relative mass Name Symbol Charge Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron n0 0 1 1.67 x 10-24

36. Subatomic Particles (cont.) • All atoms of an element have the same # of protons protons identify an atom  atomic # • Atoms are electrically neutral #p = #e- • Only neutrons and protons contribute to an atoms mass #n + #p = atomic mass

37. ISOTOPES = atoms with the same number of protons but DIFFERENT numbers of neutrons Mass Number Atomic Number Element Symbol Ex. Na-23 or Sodium-23 or 23 Na or 23Na C-14 or Carbon-14 or 14C or 14C B-10 or Boron-10 or 10B or 10B 11 6 5

38. Isotope Practice Element has • 6 p+, 8 n and 6 e- • 6 p+, 6 n and 6 e- • 19 p+, 21 n and 10 e- • __p+, __ n and __ e- • __p+, __ n and __ e- Symbol • . • . • . • Cu-65 • H-2

39. Ions • Ion is an atom that has gained or lost one or more electrons and now has a charge • Metals always lose electrons to form POSITIVE ions • Non-metals always gain electrons to form NEGATIVE ions • The charge of the element is show on the right side of the symbol as a super script • Example: Na+1, Zn+1, S-2, N-3

40. Ions practice Element Ca+2 Br-1 Al+3 I-1 Number of electrons ? ? ? ?

41. D. Average Atomic Mass • weighted average of all isotopes • on the Periodic Table • round to 2 decimal places

42. D. Average Atomic Mass • EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.

43. D. Average Atomic Mass • EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37.