Chapter 5 Models of the Atom

1. Chapter 5Models of the Atom

2. Quick Review: What are the three major subatomic particles? • What is the charge on each of these particles? • Which subatomic particle identifies the atom and which subatomic particle tells the chemical properties of an atom. Proton Electron Neutron p+, no, e- p+ e-

5. What is the structure of the atom? • Small, dense nucleus. • Nucleus is composed of neutrons and protons. • Nucleus has a positive charge. • Most of the mass of an atomis found in the nucleus • Electrons surround the nucleusand occupy most of the volume

6. The Evolution of Atomic Models 1. John Dalton (1766 – 1844) : atomic theory 2. J.J. Thomson (1856 – 1940) : electron 3. Ernest Rutherford (1871 – 1937) : nucleus Aren’t they cute!!

7. Dalton’s Theory of an Atom: • All elements are composed of submicroscopic particles called atoms • Atoms are indivisible.

8. Thomson’s “Plum Pudding” Model: • Thomson discovered the electron in 1897 • Disproved Dalton’s theory of indivisible atoms which did not take electrons and protons into account. • He described the atom as a ball of positive charge containing a number of electrons embedded into it.

9. Rutherford Model • Discovered the nucleus in 1911. • Proposed that the electrons surround a dense nucleus and the rest of the atom is empty space. • Most of the atom’s mass is concentrated in a small region called the nucleus.

10. Along came Bohr . . . Bohr asked “What prevents the electrons from falling into the nucleus?” • Proposed in 1913 that electrons are arranged in circular paths or orbits around the nucleus. • The electrons have a fixed amount of energy, in a fixed path and will not fall into the nucleus. • It was a planetary model resembling how the planets orbit the sun.

11. Niels Bohr (1885 – 1962)The Bohr Model Bohr described electrons as existing in energy levels or regions around the nucleus where e- are likely to be moving Young BohrOld Bohr Wild Bohr To Ponder… Was Bohr “Bohring”?

12. Bohr thought that electrons move in circular paths around the nucleus and that they could “jump” from one energy level to another Quantum - the amount of energy needed to move an electron from one energy level to the next

13. Think of energy levels like steps.... • The steps, like the energy levels, are not equally spaced. • The higher the step, the more energy it has, and the farther away it is from the nucleus Energy Levels Sixth Fifth Forth Third Second First Nucleus

14. The Quantum Mechanical Model “Today’s Accepted Model of the Atom” • Derived by Erwin Schrodinger mathematically in 1926. • Does not define the exact path of the electron around the nucleus, but estimates the probability of finding an electron in a certain position or ORBITAL. • The model predicts the shapes of the various orbitals.

15. Erwin Schrodinger (1887 – 1961)The Quantum Mechanical Model The Quantum Mechanical model - a mathematical model which predicts the location area and energy of an electron. Electrons have specified energies and exist in energy levels. Electrons do not have definite energy paths. You can only calculate the probability of finding an electron in a certain area.

16. Quantum energy or quantumis the amount of energy,absorbed or emitted,required to move an electron from its present energy level to a different energy level. • The amount of energy gained or lost by every electron is NOT always the same. • The energy levels are NOT evenly spaced. • Energy levels are more closely spaced the farther they are from the nucleus.

17. The probability of finding an electron is portrayed as a fuzzy cloud in diagrams. • The more dense the cloud is, the greater the probability of finding the electron in that area. • The less dense the cloud is, the lower the probability of finding the electron in that area. • An orbital cloud is only a region of probability:90% of the time you will find an electron in that cloud area.

18. Schrödinger’s Cat A famous paradox in quantum theory A cat is placed in a box, together with a radioactive atom. If the atom decays,a hammer hits a flask of prussic acid, killing the cat. Is the cat alive or dead? The answer, according to quantum mechanics, is that it is 50% dead and 50% alive.

20. Schrödinger’s Cat illustrates that since we don’t know where exactly an electron is at any given moment, it is actually in all possible states simultaneously, as long as we don't look to check. It is the measurement itself that causes the object to be limited to a single possibility.

22. What is an Energy Level? • Energy levelis the distance from the nucleus where the electron is most likely to be moving. • Electrons cannot exist between energy levels. • To move from one energy level to another, an electron must gain or lose just the right amount of energy... a “quantum” of energy

23. That’s all for today!

24. n = 2 n = 1 n = 3 n = 4 Nucleus Atomic Orbitals Energy levels are in designated quantum #’s (n). • n = 1, 2, 3, 4, 5, 6 & 7...... • A Quantum number is equal to the period • Higher the quantum number, the greater average distance from the nucleus

25. Example: What is the quantum number (n) of: • Potassium • Helium • Lead 4 1 6

26. Energy levels are divided into sub-levels. Each energy sub-level corresponds to a specific cloud shape. Cloud shape can be calculated from the atomic orbital. Atomic orbital (different from Bohr’s atomic orbit) is a region in where there is a high probability of finding an electron. Letters represent the different orbitals : s, p, d, f. s – spherical p – dumb-bell shaped d – clover-leaf shaped f – complex shape

27. S P D F

28. Note the Principal energy levels and sublevels... Aren’t they fun?

30. The principal quantum number always equals the number of sublevels within that principal energy level.

31. Electrons fill the sublevels: • s sublevel has 1 atomic orbital, and can hold 2 e- • p sublevel has 3 atomic orbitals, and can hold 6 e- • d sublevel has 5 atomic orbitals, and can hold 10 e- • f sublevel has 7 atomic orbitals, and can hold 14 e- The max number of electrons that can occupy a given energy level is 2n2where n = energylevel.

32. http://www.youtube.com/watch?v=sMt5Dcex0kg

33. Shape of orbital # of orbital (s) spherical (p) dumbbell (d) clover (f) complex 3 orbitals 5 orbitals 7 orbitals 1 orbital Reviewing the Sublevels: 2 e-Max. 6 e-Max. 10 e-Max. 14 e-Max. # of e- http://cwx.prenhall.com/petrucci/medialib/media_portfolio/09.html

34. The maximum number of electrons allowed in each orbital is 2. • The number of orbitals increases as the distance from the nucleus increases.

35. Energy sublevels (s, p, d, f) Review of Inside an Atom: Atomic Orbitals (1, 3, 5, 7) Number of electrons (2, 6, 10, 14) Principal Energy levels (7 on P.T.)

36. The 2nd energy level has how many sublevels (different orbital types)? Two • What orbital types are within the sublevel? 2s and 2p • The 3rd energy level has how many sublevels and what orbital types are within the sublevel? 3s, 3p, 3d Three

37. Symbols Used in Writing Electron Configuration Example: This is how to write the p orbital of the third energy level when it is full. 3p6 Represents the energy level (n) the electrons are in. Represents the sublevel (orbital shape). Represents the # of electrons in the sub-level. 3 p 6

38. How many orbital TYPES are in the third energy level? Energy level = # of sublevels Examples: 3 orbital types • How many TOTAL orbitals are in the third energy level? one s, three p, five d orbitals = 9 How many electrons can be in the 3rd energy level? 2(3)2 18 e- 2n2,n=3

39. That’s all for today!

40. Examples of ground state electron configurations in the orbital box notation that shows electron spins.