1. Electrons, Atoms, and the Periodic Table

2. ATOM Smallest particle of an element that retains its identity in a chemical reaction VERY SMALL!!! – A pure copper (Cu) penny contains 2.4x10 22 atoms – 100,000,000 copper atoms side by side = 1cm Made up of Subatomic Particles Mostly Empty Space Have no NET charge (neutral)

3. Early Models of the Atom Democritus (460BC – 370BC) – Atoms are indivisible and indestructible

4. Early Models of the Atom John Dalton (1766-1844) – Studied the ratios in which elements combine in chemical reactions Dalton’s Atomic Theory: 1.World was made of indivisible atoms 2.Atoms of the same element are identical 3.Atoms of different elements can physically mix and/or chemically combine 4.Chemical reactions occur when atoms separate, join, or are rearranged 5.Atoms are not changed into other atoms during a chemical reaction

5. Early Models of the Atom J. J. Thomson (1856-1940): – Discovered the electron (negatively charged subatomic particles)

6. Early Models of the Atom J. J. Thomson: – Plum-pudding model: electrons stuck in lump of positive charge

7. Early Models of the Atom Ernest Rutherford (1871-1937) – Gold Foil Experiment – Nuclear Atom: Atom is mostly empty space surrounding the area of positive charge (nucleus) that has most of the mass

8. Lab Purpose: To understand how Rutherford’s experiment lead to the development of our understanding of atomic structure Materials: Platfrom with object glued underneath, and marbles. Procedure: ? Data/Observations: (Maybe draw or explain your results) Conclusion: What did you learn? How does this experiment mirror Rutherford’s. What did you conclude, and what was Rutherford’s conclusion about the atom?

9. Bohr Model Neils Bohr (1885-1962) Electrons are found in certain circular orbits around the nucleus with fixed energy levels

10. Electron Cloud Model Erwin Schrodinger (1887-1961) – Electrons still have certain energies (like Bohr) – Electrons don’t have an EXACT path

11. Atomic Models Dalton - 1803 Thomson – 1897 “Plum-Pudding” Rutherford - 1911 Bohr - 1913Schrodinger – 1926 “Electron-Cloud”

12. Quantum Mechanical Model Determines: – Allowed energies an electron can have – Where an electron can be located (probability)

13. Quantum Numbers Think of these as the address of the electron… Quantum numbers are the “Allowed Energies” or “Energy Levels” Principal Quantum Number (n): 1, 2, 3… Higher number = Higher energy level = Larger orbital Electrons with the same value of “n” are in the same “shell”

14. Quantum Numbers Atomic Orbitals: Where electron is likely to be found s: spherical p: dumbbell d:cloverleaf f: complicated…

15. Principal Energy Level (Quantum Number) Number of SublevelsType of Sublevel n=111s (1 orbital) n=222s (1 orbital) 2p (3 orbitals) n=333s (1 orbital) 3p (3 orbitals) 3d (5 orbitals) n=444s (1 orbital) 4p (3 orbitals) 4d (5 orbitals) 4f (7 orbitals) For every energy level you add, you add 1 sublevel… Every sublevel has 2 more orbitals than the previous…. Ex: p has 2 more than s…

16. Electron Configuration Electron Configuration: The ways electrons are arranged in orbitals around the nucleus. Three rules tell you how to find the electron configuration: – Aufbau Principal – Pauli Exclusion Principal – Hund’s Rule

17. Aufbau Principle Electrons fill the orbitals of the lowest energy first – Which orbitals would be filled first: 1s or 2s? 2s or 2p? 3p or 3d? 3d or 4s?

18. Pauli Exclusion Principle An orbital has AT MOST 2 electrons. The pair will have opposite spins (clockwise and counterclockwise) – How many electrons are in the 1s orbital? – How many electrons are in a p orbital? The entire p sublevel?

19. Hund’s Rule All orbitals in a sublevel must have an electron before any orbital in that sublevel can have 2 electrons.

20. Electron Configurations What is the electron configuration for: H? He? Be? N? Electron Orbital Diagram

21. Find the electron configuration and draw the electron orbital diagrams for: NeNaBMg

22. Electron Configurations Valence Shell: The outermost electrons – have the highest principal quantum number full configurationvalence configuration Ex: O1s 2 2s 2 2p 4 2s 2 2p 4 Cl Nobel Gas Core: The full level under the valence shell – the last nobel gas before the element full configurationCore configuration O1s 2 2s 2 2p 4 [He] 2s 2 2p 4 Ca

23. Evolution of the Periodic Table Arranged the elements in order of increasing atomic mass (Mendeleev – 1800s) Periodic Law: Elements are arranged in order of increasing atomic number (modern)

24. Electron Configurations in Groups Where an element is on the periodic table is determined by its electron configuration (WHY?) – Hint: Valence Electrons

25. Electron Configuration Lab

26. Electron Configuration in Groups Noble Gases – Find the electron configuration of: He Ne Ar Kr

27. Electron Configuration and Groups Alkali Metals – Lithium 1s 2 2s 1 – Sodium 1s 2 2s 2 2p 6 3s 1 – Potassium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

28. Electron Configuration and Groups Group 6A – Oxygen 1s 2 2s 2 2p 4 – Sulfur 1s 2 2s 2 2p 6 3s 2 3p 4 – Selenium 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 4 * Note the number of valence electrons and the group number!

30. Periodic Trends – Atomic Size Atomic size: – increases down a column – decreases across a row (from left to right)

32. WHY?

33. Why the size difference? Increase down a column: Atomic #Element Electron Configuration 3Li1s 2 2s 1 11Na1s 2 2s 2 2p 6 3s 1 19K1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 37Rb1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1

34. Why the size difference Decrease across a row: – The atomic number is changing…. ElementAtomic #Configuration Al131s 2 2s 2 2p 6 3s 2 3p 1 Si141s 2 2s 2 2p 6 3s 2 3p 2 P151s 2 2s 2 2p 6 3s 2 3p 3 S161s 2 2s 2 2p 6 3s 2 3p 4 Cl171s 2 2s 2 2p 6 3s 2 3p 5 Ar181s 2 2s 2 2p 6 3s 2 3p 6

35. How is the size of a cation different from the size of the atom? How is the size of an anion different from the size of the atom?

36. How is the size of a cation different from the size of the atom? How is the size of an anion different from the size of the atom?

37. Periodic Trends – Ionization Energy Ionization Energy: energy required to remove an electron from an atom Ionization Energy: – Decreases down a column – Increases across each row (from left to right)

38. Periodic Trends - Electronegativity Electronegativity: The measure of how great an atom can attract electrons when it is in a compound Electronegativity: – Decreases down a column – Increases across a row (from left to right)