Chemistry Science 09 Review!!

1. Chemistry Science 09 Review!! Science 10 CT00D01

2. Topics from Science 09 • Atomic Theory • Subatomic Composition • The Periodic Table • Chemical Bonding • Nomenclature and Chemical Reactions

3. Atomic Theory Major Contributors: Dalton, Thomson, Rutherford, Bohr

4. John Dalton • England, 1808 • Atomic Theory which states: • An atom cannot be broken down into smaller parts. • Atoms of the same element are exactly alike. • Atoms of different elements are different. • Atoms can combine in definite whole-number ratios to form compounds.

5. Dalton Model Diagram: Shading is used in the sphere to show that it’s solid throughout. Every substance is made up of atoms Atoms are indestructible and indivisible Atoms can neither be created or destroyed

6. J.J. Thomson • England, 1897 • Discovery • Atoms contain electrons • Thomson found the charge-to-mass ratio of an electron.

7. - - + + - + + - + - - + + + - - + + - - + + - + - Thomson Model Diagram • This model is referred to as the ‘Plum Pudding model.’ Thomson believed that the atom consisted of a positive sphere (the pudding) with electrons embedded in it (the raisins).

8. Ernest Rutherford • New Zealand, 1911 • Discovery • Atoms are made up of mostly empty space • Atoms contain a small, dense, positively-charged nucleus. • Gold Foil Experiment by Ernest Rutherford • Positively charged radiation directed towards a thin sheet of gold foil • Odd results: most radiation went through the foil, some scattered at wide angles, and some shot back at him!

9. - - - - - - - + Rutherford’s Model Diagram Dense, positively charged nucleus + Empty Space

10. Niels Bohr • Denmark, 1913 • Discovery: • Electrons contain specific amounts of energy and orbit the nucleus in specific paths, call energy levels • Electrons must gain energy to move to a higher energy level or lose energy to move to a lower energy level

11. e- e- e- Bohr Model Diagram • The small particles in the center represent the protons and neutrons in the nucleus. Nucleus contains protons (+) and neutrons (o) Electrons exist in energy levels, 2-8-8, then….

12. Subatomic Composition Atoms consist of: Protons, neutrons, electrons Isotopes, ions, anions, cations

13. Three charges of an atom… Mass of p & n are ~1000x greater than mass of e (making it insignificant)

14. First look at the Periodic Table: What do the numbers mean? Atomic Number (p) Atomic Symbol Ar (Relative At. Mass) For most common isotope, round it off. (For H would be 1 = 1p & 0n)

15. How many of each are there? • Protons: always the atomic # • Electrons: • same as p if neutral • If -1, has one more e than p • If +1, has one less e than p • Neutrons: • Depends on isotope • For most common, round off the atomic mass to find the mass number, then use the formula Ar = p + n Mass Number = protons + neutrons

16. Atomic Number Mass Number • Number of protons and neutrons in an atom (p + n) • Can change for various isotopes • Number of protons contained in an atom (p) • Determines the element • Every element of that kind has the same number of protons

17. Ions Isotopes • Differing number of neutrons (n) • Changes Mass Number • Differing number electrons (e) • Changes overall charge • Cations (positive) • Anions(negative)

18. Isotopes • An isotope is an element that contains a different number of neutrons and protons. • Affects the molar mass of the element • Longhand notation • Element – Mass Number • Carbon – 12 • Carbon – 14 • Shorthand notation Charge Mass# X Atomic #

19. The Periodic Table Layout: families, groups, etc Trends: Atomic number, mass, valence, charge Electronegativity, ionization energy, atomic radius, reactivity, metallic property

20. PERIODS - Similarities: The number of outer electron shells.

21. GROUPS – Similarities: The number of electrons in the outer shell. Common reactivity, bonding, chemical and physical properties.

22. Halogens Non-metals Alkali Metals Metaloids Alkaline Earths Weak/Poor Metals Transition Metals Noble Gases Lanthanides Actinides

23. METALIC PROPERTIES Similarities: An elements relative ability to conduct energy in the form of heat or electricity. Non metals the “stair” Metals Metaloids

24. Valence Electrons 1 2 The number of electrons in the outer shell of an atom 1 2 3 4 5 6 7 8 2 2 2 2 2 2 2 2 2 3 4 5 6 7 8 1 8 8 8 8 8 8 8 8 2 2 2 2 2 2 2 2

25. Valence Electrons

26. Periodic Law • When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

27. Atomic Radius A measure of the distance from the center of the nucleus to the outer-most electron (Bottom Left, Fr has the largest Atomic Radius)

28. Electronegativity An atoms ability or affinity to gain another electron. Trend (Top right, excluding the noble gasses, F is the most electronegative element)

29. Table of Electronegativities

30. Ionization Energy The energy required to lose the outer-most electron from an element. (Top right, He has the greatest ionization energy)

31. An atoms general ability to undergo a chemical reaction. The Alkali Metals are the most electronegative family/group. (Bottom Left, Fr is the most reactive) Reactivity • Highly reactive – readily participates in chemical reactions • Un-reactive – does not readily participate in chemical reactions http://teachertube.com/viewVideo.php?video_id=41344&title=Alkali_Metals___Brainiac

32. Trends of the Periodic Table Atomic Radius: A measure of the distance from the center of the nucleus to the outer-most electron Electronegativity: An atoms ability or affinity to gain another electron. Ionization Energy: The energy required to lose the outer-most electron from an element. Reactivity: An atoms general ability to undergo a chemical reaction.

33. 4 major trends: top right or bottom left? Electronegativity (F) & Ionization Energy (He) Atomic Radius (Fr) & Reactivity (Fr)

34. Chemical Bonding Inter- vs Intra-molecular forces Ionic vs Covalent

35. Molecular Interactions • Inter-molecular Forces • Interaction between molecules that hold it together in a network. • Intra-molecular Forces • Forces that hold groups of atoms together and make them function as a unit

36. Intra-molecular Forces: Bonding • Forces that hold groups of atoms • together and make them function • as a unit. • Ionic bonds – transfer of electrons • Covalent bonds – sharing of electrons • Metallic Bonding– sea of electrons

37. Which elements form which type of bond? • Metal / nonmetal = Ionic • NaCl, MgBr2 • Nonmetal / nonmetal = covalent • CO2, CH4, H2O, NO • Group like metals = metalic • Fe, Ti, Mg

38. Review: Why do atoms bond? • To satisfy the octet rule? • Yes, but to be more specific, atoms share electrons in order to complete their outer electron shell making them more stable as they are then in a lower state of energy. • But how do we know what type of bonding will occur between two atoms?

39. Lewis Dot Diagrams I II III IV V VI VII VIII Transition metals • Lewis Dot Diagrams are used in both ionic and covalent bonding Metalloids Nonmetals Metal

40. Simple Rules for Bonding • Duet Rule: • H and He require 2 electrons for stability • Octet Rule: • All other elements will have 8 valence electrons for stability • In order to achieve this elements can steal (ionic) or share (covalent) electrons

41. Ionic Bonds Ions form when atoms lose or gain electrons. Atoms with few valence electrons tend to lose them to form cations. Atoms with many valence electrons tend to gain electrons to form anions N O F Ne Na Mg Na+ Mg2+ N3- O2- F- Cations Anions 41

42. Ionic Bonds Ionic bonds result from the attractions between positive and negative ions. Ionic bonding involves 3 aspects: • loss of an electron(s) by one element, • gain of electron(s) by a second element, • attraction between positive and negative 42

43. Ionic Bonds • Electrons are transferred • Electronegativity differences are • Greater than for covalent • The formation of ionic bonds is • always exothermic!

44. Covalent Bonds Why should two atoms share electrons? + 7e- 7e- 8e- 8e- F F F F F F F F lone pairs lone pairs single covalent bond single covalent bond lone pairs lone pairs A covalent/molecular bond is a chemical bond in which two or more electrons are shared by two atoms. Lewis structure of F2

45. Covalent Bonds single covalent bonds H H H H or H H O 2e- 2e- O 8e- O C O C O O double bonds 8e- 8e- 8e- double bonds O N N triple bond N N triple bond 8e- 8e- + + Lewis structure of water Double bond – two atoms share two pairs of electrons or Triple bond – two atoms share three pairs of electrons or

46. Covalent Bonds 180o Linear

47. Covalent Bonds 120o Trigonal Planar

48. Covalent Bonds 109.5o Tetrahedral

49. Covalent Bonds 107o Pyramidal (Tetrahedral)

50. Covalent Bonds 104.5o Bent (Tetrahedral)