1. Atoms and the Periodic Table

2. Atoms  Atoms are the smallest pieces of matter that contain all the properties of a specific element  Each element contains only one type of atom

3. Atomic Models  Have been revised many times to explain new discoveries  Democritus (4 th century B.C.) thought all matter was made of particles he called the atom  Theory was modified when subatomic particles were discovered

4. Atomic Models  Plum Pudding Model  (1904) developed by J.J. Thomson  Planetary Model  (1911) developed by Ernest Rutherford

5. Newer Models  Bohr’s Model (1913)  Developed by Niels Bohr  Electron Cloud Model (1925) 

6. Inside the Atom  Atoms are made up of smaller particles  These particles are found in different regions of the atom

7. Protons (p + )  Positively charged particles found in nucleus of atom  Have an electrical charge of +1  Mass of 1 a.m.u.  Composed of quarks  Discovered by Ernest Rutherford using Gold Foil Experiment

8. Protons  The number of protons in an atom determines its identity  All oxygen atoms have 8 protons, all uranium atoms have 92 protons  If the number of protons change the identity of the atom changes.

9. Neutrons (n 0 )  Neutral particles found in nucleus of atom  Have no electrical charge  Mass of 1 a.m.u.  Composed of quarks  Discovered by James Chadwick

10. Nucleus  The nucleus is the positively charged dense core in the center of the atom  Houses protons and neutrons  Contains 99.9% of mass of atom  Extremely small compared to the entire size of the atom

11. Electrons (e - )  Negatively charged particles found in electron cloud  Have an electrical charge of -1  Constantly moving around outside nucleus  Have essentially no mass  Discovered by J.J. Thomson during cathode ray experiment

12. Electrons  The number of electrons in a neutral atom is equal to the number of protons  Neutral oxygen has 8 protons, therefore it has 8 electrons  Neutral lead has 82 protons, therefore, it has 82 electrons

13. Valence Electrons  Electrons in the outermost energy level of an atom are called valence electrons  These are the electrons furthest from the nucleus

14. Symbols  Elements are identified by their chemical symbols  Symbols are usually either one capital letter like C for Carbon, or one capital and one lowercase letter like Ne for Neon

15. Atomic Number (Z) Whole number shown on periodic table  Periodic table is arranged by atomic number Atomic Number = # of Protons *Also gives the number of electrons if the atom is neutral

16. Atomic Number

17. Mass Number (A)  The mass number is the sum of the total number of protons and neutrons in the atom  Mass # = # p + + # n 0

18. Atom Math Atomic Number Symbol Element Name Atomic Mass

19. Isotopes  Isotopes are atoms of the same element that have different numbers of neutrons  All atoms are isotopes  Each element has isotopes that are more common than others

20. Nuclear Symbol  Isotopes can be designated with their nuclear symbol

21. Hyphen Notation  Isotopes can also be designated using hyphen notation Carbon-16 Element NameMass Number

22. Write the Nuclear Symbol and Hyphen Notation for the Following Isotopes  Lithium isotope with 3 protons and 4 neutrons  Sulfur isotope with 17 neutrons  Lead with 122 neutrons

23. Ions  Ions are atoms or groups of atoms that have a net positive or negative charge  The charge results from an unequal number of electrons and protons within atom of group of atoms

24. Ions  Ions with more electrons than protons have a negative charge  For each extra electron the negative charge increases by one  Ions with less electrons than protons have a positive charge  For each missing electron the positive charge increases by one

25. Ions  The charge of an ion is indicated on its element symbol with a positive or negative sign on the upper right side and number equal to the magnitude of the charge  The number one is not included

26. Ions  F -  9 protons – 10 electrons = -1 charge  Ca 2+  20 protons – 18 electrons = -2 charge  P 3-  15 protons – 18 electrons = -3 charge

27. Common Ions (Need to memorize these)  Many of the main group elements in the same column form ions with same charge  Column 1 (Li, Na, K, Rb, Cs): +1  Column 2 (Be, Mg, Ca, Sr, Ba): +2  Column 13 (Al, Ga): +3  Column 15 (N, P, As): -3  Column 16 (O, S, Se, Te): -2  Column 17 (F, Cl, Br, I): -1

28. Atomic Mass  The average atomic mass is the number at the bottom of this square  Found by averaging the natural abundances of its isotopes

29. Calculating Average Atomic Mass (amu) If abundance is given as percent value:  If abundance is given as decimal value:

30. Average Atomic Mass Rubidium has two common isotopes, Rb-85 and Rb-87. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?

31. Uranium has three common isotopes. If the abundance of 234U is 0.0001, the abundance of 235U is 0.0071, and the abundance of 238U is.9928, what is the average atomic mass of uranium?

32. Electron Configuration  An electron configuration is a shorthand description of how electrons are arranged around the nucleus of an atom.  Electrons generally do not orbit the nucleus rather, they move in certain patterns based on three factors: (1) their energies, (2) their orientations in space and (3) their angular momentum

33. A. General Rules  Pauli Exclusion Principle  Each orbital can hold TWO electrons with opposite spins.

34. A. General Rules  Aufbau Principle  Electrons fill the lowest energy orbitals first.  “Lazy Tenant Rule”

35. O 8e -  Orbital Diagram  Electron Configuration 1s 2 2s 2 2p 4 B. Notation 1s 2s 2p

36. RIGHT WRONG A. General Rules  Hund’s Rule  Within a sublevel, place one e - per orbital before pairing them.  “Empty Bus Seat Rule”

37.  Shorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 B. Notation  Longhand Configuration

38. © 1998 by Harcourt Brace & Company s p d (n-1) f (n-2) 12345671234567 6767 C. Periodic Patterns

39.  Period #  energy level (subtract for d & f)  A/B Group #  total # of valence e -  Column within sublevel block  # of e - in sublevel

40. s-block1st Period 1s 1 1st column of s-block C. Periodic Patterns  Example - Hydrogen

41. C. Periodic Patterns  Shorthand Configuration  Core e - : Go up one row and over to the Noble Gas.  Valence e - : On the next row, fill in the # of e - in each sublevel.

42. [Ar]4s 2 3d 10 4p 2 C. Periodic Patterns  Example - Germanium

43.  Full energy level  Full sublevel (s, p, d, f)  Half-full sublevel D. Stability

44.  Electron Configuration Exceptions  Copper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10  Copper gains stability with a full d-sublevel. D. Stability

45.  Electron Configuration Exceptions  Chromium EXPECT :[Ar] 4s 2 3d 4 ACTUALLY :[Ar] 4s 1 3d 5  Chromium gains stability with a half-full d- sublevel. D. Stability

46.  Ion Formation  Atoms gain or lose electrons to become more stable.  Isoelectronic with the Noble Gases.

47. O 2- 10e - [He] 2s 2 2p 6 D. Stability  Ion Electron Configuration  Write the e - config for the closest Noble Gas  EX: Oxygen ion  O 2-  Ne

48. D. Lewis Diagrams  Also called electron dot diagrams  Dots represent the valence e -  Ex: Sodium  Ex: Chlorine Lewis Diagram for Oxygen

49. Steps to Draw Lewis Structures 1.Determine how many valence are in the element. 2.Staring on the Right side of the element draw a dot to represent a valence electron. 3.Place one dot on each side of the symbol. One electron must be drawn around each side of the element before a second electron can be added to any side.

50. Steps to Draw Lewis Structures