1. Atoms and the Periodic Table

2. Atomic Models  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

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

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

5. 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

6. 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.

7. 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

8. 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

9. 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

10. 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

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

12. 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

13. 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

14. Atomic Number

15. Mass Number (A)  The mass number is the sum of the total number of protons and neutrons in the atom  Mass # = # p + + # n 0  The mass number is not found on the Periodic Tablw

16. 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

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

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

19. 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

20. 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

21. Ions  Anions  Ions with more electrons than protons resulting in a negative charge  For each extra electron the negative charge increases by one  Cations  Ions with less electrons than protons resulting in a positive charge  For each missing electron the positive charge increases by one

22. Ions  Ions are symbolized 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 Magnitude Charge

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

24. Common Ions (Need to memorize these)  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

25. Average Atomic Mass  The weighted average of the naturally occurring isotopes of an element.  Found by averaging the natural abundances of its isotopes

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

27. 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?

28. 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?

29. Columbic Attraction  The attraction between oppositely charged particles  The greater the distance between the particles the weaker the attraction

30. Columbic Attraction in the Atom  The electrons in the atom are attracted to the protons  Electrons closest to the nucleus feel a stronger attraction force than electrons on the outermost energy level  As you move in a row from left to right on the Periodic Table the number of protons in an atom increases and so the attractive force on the outermost electrons increases

31. Bohr Model of the Atom  Niels Bohr found that electrons occupy distinct energy levels within the atom

32. Orbital Diagrams  Represent the electrons within energy levels and sublevels using arrows

33. Orbital Diagrams  Pauli Exclusion Principle  Each sublevel can hold TWO electrons with opposite spins.

34. Sub levels Each energy level has a different amount of sublevels 1  s 2  s, p 3  s, p, d 4  s, p, d, f 5  s, p, d, f  s has 1 sublevel  2 e -  p has 3 sublevels  6 e -  d has 5 sublevels  10 e -  f has 7 sublevels  14 e -

35. Orbital Diagrams  Aufbau Principle  Electrons fill the lowest energy level first. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p

36. RIGHT WRONG Orbital Diagrams  Hund’s Rule  Within a sublevel, place one e - per orbital before pairing them.  The last energy level being filled may not be completely filled with arrows

37.  Fluorine  Germanium B. Notation

38. Electron Configuration  An electron configuration is a shorthand description of how electrons are arranged around the nucleus of an atom.  Electrons within each energy level and sub level are represented with numbers in super script rather than arrows

39. S 16e - Valence Electrons Core Electrons 1s 2 2s 2 2p 6 3s 2 3p 4 Electron Configuration  Longhand Configuration

40.  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

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

42. 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.

43. Steps to Draw Lewis Structures