Atom Unit Chapters 5, 13, & 14

1. Atom UnitChapters 5, 13, & 14 Atoms and the Periodic Table

2. Atom: smallest particle of an element that retains the properties of that element

3. Atomic Bombs • Fission Bomb: the splitting of a nucleus into smaller fragments, accompanied by the release of neutrons and a large amount of energy • Fusion Bomb: a rxn in which 2 light nuclei combine to produce a nucleus of heavier mass, accompanied by the release of a large amount of energy • Example: H-Bombs

4. Fission BombsMaterials • The core of a fission bomb is either plutonium (Pu) or highly enriched uranium (U). These are the only materials that can achieve a self-sustaining chain reaction • Pu and U atoms are both heavy, meaning they have a large number of protons and neutrons in the nucleus. Fission of a heavy nucleus can be spontaneous or induced by the absorption of a neutron. During fission, when the heavy nucleus splits into two smaller nuclei, extra neutrons are released. If these neutrons are absorbed by other nuclei, they in turn could split, also releasing neutrons, thus creating a chain rxn

5. Fusion Bombs • In fusion bombs, deuterium (2H) and tritium (3H) - two isotopes of hydrogen - are fused together to create heavier atoms • This is the same reaction as occurs in the center of the sun and other stars. Fusion can only happen at very high temperatures and pressures • In a nuclear weapon, these are created through using a fission explosion (i.e. an atom bomb) to trigger the fusion reaction. There is no theoretical limit to the explosive force of a fusion weapon • Typically, fusion weapons are 10 - 100 times as explosive as the fission bombs which nearly destroyed Hiroshima and Nagasaki

6. Atomic Structure • Democritus is an old Greek guy (4th century BC) who suggested that the universe was made of invisible units called atoms • The word atom means “unable to be divided” • He was unable to provide evidence that atoms existed

7. In 1808 John Dalton proposed that: • Every element is made of tiny, unique particles called atoms that cannot be subdivided • Atoms of the same element are exactly alike • Atoms of different elements can join to form molecules

8. Electron Cloud Model Dalton’s Atomic Model: Spherical Bohr’s Atomic Model

9. Thompson Atomic Model • A tiny ball of positive charge containing a number of electrons

10. Rutherford’s Atomic Model • Most of the atom’s mass is concentrated in a small, positively charged region called the “nucleus” • He determined this by Rutherford’s Gold Foil Experiment

11. Gold Foil Experiment

12. Quantum Mechanical Model“Electron Cloud” • Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space

13. Atomic Orbitals • Electrons are found in orbitals with distinct energy levels • Examples of orbitals are the “S”, “P”, “D”, and “F” orbitals • They all have specific shapes and paths that electrons follow

15. Principal Energy LevelsSee chart on page 364: Table 13.1 • Electrons will occupy the lowest energy levels in the atom first • They always occupy the “s” orbital first, then “p”, “d”, and “f” (which has the most energy) • Every atom has one or more valence electrons: these are the electrons in the outer most energy level

16. Electron Configuration Rules • Aufbau Principle: electrons enter orbitals of lowest energy first! (SPDF) • Pauli Exclusion Principle: an atomic orbital may hold at most 2 electrons. Two electrons occupy the same orbital and have opposite spin. • Hund’s Rule: when electrons occupy orbitals of equal energy one electron enters each orbital until all the orbitals contain one electron with parallel spins.

18. Remember: Read the table from left to right NOT down the columns!!

19. Examples: • H • Li • N • Na • Sr • Cu • Br 1S1 1S2 2S1 1S2 2S2 2p3 1S2 2S2 2P6 3S1 1S2 2S2 2P6 3S2 3P6 4S2 3D10 4P6 5S2 1S2 2S2 2P6 3S2 3P6 4S2 3D9 1S2 2S2 2P6 3S2 3P6 4S2 3D10 4P5

20. Orbital Diagrams

21. How to Draw Orbital Diagrams • Find the element on the periodic table • Get the atomic number (this is the number of electrons that you will be using) • Get electron configuration • Place up arrows in each sublevel until full then place down arrows to fill level with max of two electrons 1S2 2S2 2P4

23. The Periodic Table • Structure of the Periodic Table • Periodic means “repeated in a pattern” • Mendeleev, Russian chemist, discovered there was a pattern to elements put onto a table. He searched for a way to organize elements. He arranged them in order of increasing atomic mass. • The periodic table of elements is the arrangement of elements according to repeated changes in properties HOTTIE!

24. Groups of elements: the vertical columns in the periodic table are called “groups” or families • Ex: Group 1 contains Hydrogen, Lithium, Sodium, Potassium, etc… • Elements in the same group have similar properties because they have the same electron arrangements. Elements in group 1 have 1 electron in their outermost shell • It is the number of electrons in the outer energy level, or valence electrons, that determines the chemical properties • Dot Diagrams: the symbol of the element and dots to represent the electrons in the outer shell (valence e-) • Ex: dot diagram of sodium • Ex: dot diagram of oxygen

25. GROUPS PERIODS

27. Group Properties

28. The Alkali Metals • Highly reactive metals • The alkali metals are: Li, Na, K, Rb, Cs, & Fr • These elements are found in nature only in compounds

29. Alkali Metal Reactions in Water http://www.youtube.com/watch?v=eCk0lYB_8c0&feature=related

30. Each of the alkali metals has one electron in its outer energy level: they are in group ONE! • The electron is given up when an alkali metal combines with a nonmetal (to form an ionic compound). • The attraction is very strong… • http://www.youtube.com/watch?v=Mx5JJWI2aaw

31. Alkaline Earth Metals • The alkaline earth metals make up group 2 and have TWO e- in valence shell • They are shiny, malleable, ductile • So reactive that they are also found only in compounds in nature • The elements in this group are: Be, Mg, Ca, Sr, Ba, & Ra

32. http://www.youtube.com/watch?v=d8hpUtRnsYc&feature=related Metallic Flame Test

33. Periodic Trends • Atomic Radius: ½ the distance between the nuclei of two of the same atoms in a diatomic molecule • Atomic radius indicates it’s relative size

35. Ionization Energy • When an atom gains or loses an electron, it becomes an ion • Ionization Energy: the energy required to overcome the attraction of the protons in the nucleus to remove an electron from a gaseous atom

37. Ionic Size • Atoms of metallic elements have low ionization energies • Thus they form positive ions easily • Atoms of nonmetallic elements readily form negative ions • Positive atoms are smaller than their neutral atom because the loss of electrons causes a decrease in radius • The opposite is true for negative ions

38. Anions Decrease Cations Decrease Genera l l y Increase

39. Trends in Electronegativity • Electronegativity: tendency for the atoms of an element to attract electrons when they are chemically combined with atoms of another element

41. Flourine has an electronegativity value of 4.0, it is the strongest electronegative value • Flourine has a strong tendency to attract electrons when it is chemically bonded to any other element

42. See Figure 14.16 on pg. 406

43. Trends Down a Group Across a Period Atomic radius decreases Ionic size decreases Ionization E increases Electronegativity increases • Atomic radius increases • Ionic size increases • Ionization E decreases • Electronegativity decreases

44. THE END!