1 / 44

Unit 11

Unit 11. Modern Atomic Theory. Review of the Discovery of the Atom. 1803 John Dalton discovered that elements are made of atoms. He thought that atoms were solid, like a marble . 1875 Crooks discovered the electron. The electron was found to have a mass of

bernad
Download Presentation

Unit 11

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Unit 11 Modern Atomic Theory

  2. Review of the Discovery of the Atom 1803 John Dalton discovered that elements are made of atoms. He thought that atoms were solid, like a marble. 1875 Crooks discovered the electron. The electron was found to have a mass of 9.1 x 10-28 g and have a negative (-) charge.

  3. 1907 J.J.Thomson finds the positive particle, the proton. Protons have a much larger mass of 1.67 x 10-24 g and a positive (+) charge. His model of the atom was called the plum pudding model, the p+ and e– were spread throughout the mass of the atom.

  4. 1911 Rutherford shoots a stream of helium nuclei at gold foil. He discovers that they bounce back after hitting a positive nucleus composed of p+ in the center of the atom.

  5. Rutherford stated that the p+ are found in the nucleus in the center of the atom. e- revolve around the nucleus.

  6. 1932 Chadwick discovers the neutral particle, the neutron. The neutron is found in the nucleus along with the p+. The e- revolve around the nucleus. Protons and Neutrons located in the nucleus

  7. Subatomic Particles

  8. Review of the Periodic Table 1829 Johann Dobereiner organized some elements into groups of three called triads. This was the first attempt at a periodic table. Example of triad: Cl, Br, I Other triads were: Li, Na, K Ca, Sr, Ba S, Se, Te

  9. 1869 Dmitri Mendeleev organized the elements according to their atomic mass. The vertical columns were families (or groups).

  10. Early 1900’s Henry Mosley organized the elements according to their atomic number (number of protons). This is the modern periodic table that we use.

  11. Elements with the atomic mass in ( ) means that very little of it exists at any given time. Scientists have not been able to attain an good average atomic mass. Atomic mass is the average number of p+ plus n found in the nucleus. Carbon 12.0111 Chemical symbol: 1 or 2 letters used to represent an element. The 1st letter is always capital. If there is a 2nd letter, it is lower case. C 6 Atomic number is the number of p+ found in the nucleus.

  12. Atomic mass is the average number of p+ and n found in the nucleus. All elements have Isotopes. Atoms of the same element with a different number of neutrons found in the nucleus, therefore they have different masses. Carbon 12.0111 Chemical symbol: 1 or 2 letters used to represent an element. The 1st letter is always capital. If there is a 2nd letter, it is lower case. C 6 Atomic number is the number of p+ found in the nucleus.

  13. Periodicity : regular, repeating properties; the elements when arranged in the order of their atomic numbers show a periodic variation in most of their properties. • Periods are the rows across the table. There are 7 periods. These also correspond to the energy levels that e- are found in. There is a variation in the properties of elements as you move from left to right, which repeats in the next period. • Families (or groups) are the columns going down the table. Elements in a family show very similar properties. • Go back to the last slide and check out the numbering of the periods and families.

  14. Metals : All elements to the left of the red line (metalloids) 1. conduct electricity 2. conduct heat 3. shiny silvery solids (except Au, Cu, Hg) 4. malleable and ductileNonmetals: All elements to the right of the red line (metalloids) 1. do not conduct electricity 2. do not conduct heat 3. not shiny silvery solids 4. not malleable or ductileMetalloids: All elements with a full side bordering the red line B, Si, Ge, As, Sb, Re, PoThese elements have some of the characteristics of metals and some characteristics of nonmetals.

  15. Alkali Metals : Li, Na, K, Rb, Cs, Fr • Alkali metals are all very reactive and cannot be used by themselves. • Melting and boiling temperatures decrease as you go down the family (or group). This is due to the increase in the size of the atoms. Melting: Li = 186 oC, Na=97oC, K=62 oC, Rb=38 oC, Cs=28 oC. • Ionization energy decreases as you move down the family. Again due to the increase in the size of the atoms. • All Alkali metals react with water. (M = alkali metal) 2 M (s) + 2 H2O (L) 2 M+1 (aq)+ 2 OH-1 (aq)+ H2(g) + Energy 2 Na (s) + 2 H2O (L) 2 Na+1 (aq)+ 2 OH-1 (aq)+ H2(g) + Energy • All form +1 ions

  16. Alkali Earth Metals: Be, Mg, Ca, Sr, Ba, Ra • Alkali Earth Metals are quite reactive. • Alkali Earth metals have high melting points. The melting point generally decreases as you go down the family because of the increase in size of the atom. Be=1280oC, Mg=650oC, Ca=839oC, Sr=769oC, Ba=725oC • These elements have many uses. Mg combines with other metals to make light weight alloys. Ca compounds include chalk and marble. Sr and Ba are used in fireworks. • Form +2 ions.

  17. Halogens: F, Cl, Br, I, At • Melting and boiling temperatures increase as you go down the family. • Ionization energy decreases as you move down the family. • These elements have many uses. Cl is used to kill bacteria. F is used to prevent tooth decay. I is used in medicine (thyroid hormone – thyroxine). Br is used to improve gasoline performance. • Make -1 ions. • The halogens all react with alkali metals. 2 M (s) + X2 2 MX(s) + energy M = alkali metal 2 K (s) + Cl2 (g)  2 KCl (s) + energy X2 = halogen

  18. Noble Gases: He, Ne, Ar, Kr, Xe, Rn • Noble gases do not chemically react. • As a family, the Noble gases have the lowest boiling temperatures known. The melting and boiling temperatures increase as you go down the family. Boiling points: He=-268.9oC, Ne=-229oC, Ar=- 185.7 oC, Kr=-156.6 oC, Xe=- 111.9 oC • Ionization energy decreases as you go down family due to atomic size increase. • Some very unstable compounds of Xenon and Krypton have been made. • Generally don’t make ions.

  19. Transition Metals: Element 21-30 and below • Many useful applications • Most can have more than one type of ionic charge. Example: Fe+2 and Fe+3 Rare Earth Elements: Lanthanide & Actinide Series • Silvery white or gray solids • Have common properties so they are hard to separate from each other. • Up to element 92 are naturally occurring together with other elements in rocks. Beyond element 92 are man made. • Usually have +3 ions

  20. Ion charge Halogens +1 +2 +3 + 4 -3 -2 -1 Alkali Earth metals Noble gases Alkali metals Transition metals 3 e- 4 e- 5 e- 6 e- 7 e- 8 e- 1 e- 2 e- # e - in outer energy level Rare Earth elements Lanthanide series Antinide series

  21. Relative atomic sizes for selected atoms.

  22. Rutherford’s model of the atom is incorrect! • The e- should move into the nucleus (unlike charges attract) • The nucleus should fly apart (like charges repel) Neils Bohr proposed a new model for the atom. The nucleus contains the p+ and N, and the e- occupy certain energy levels.

  23. The various energy levels can hold 2n2e-. 1st energy level = 2(1)2 = 2 e- 2nd energy level = 2(2)2 = 8 e- 3rd energy level = 2(3)2 = 18 e- 4th energy level = 2(4)2 = e- 5th energy level = 2(5)2 = e-

  24. The most recent model of the atom is the orbital model. It helps to explain magnetism, molecular geometry, and bonding energy. • Each energy level has the possibility of having the same number of sublevels as energy level number, with a maximum number of 4 sublevels. • The sublevels are given the letters s, p, d, f • Each sublevel has orbitals. An orbital is a region of space around a nucleus where e- are likely to be found. • Each orbital can hold a maximum of 2 e-with opposite spin.

  25. e- locations Energy level sublevels orbitals e- 1st s 1 2 = 2 e- 2nd s 1 2 p 3 6 = 8 e- 3rd s 1 2 p 3 6 d 5 10 = 18 e- 4th s 1 2 p 3 6 d 5 10 f 7 14 = 32 e-

  26. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

  27. The three 2p orbitals.

  28. Diagram of principal energy levels 1 and 2.

  29. Orbitals being filled for elements in various parts of the periodic table.

  30. Orbital Notation & Electron Configuration 1s 2s 2p H 1p+ 1 e- He 2p+ Li 3p+ Be 4p+ B 5p+ C 6p+ N 7p+ O 8p+ F 9p+ Ne 10p+

  31. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Filling order for placement of e- in the orbitals.

  32. Electromagnetic radiation travels on waves. All gaseous elements emit light of certain frequencies throughout the electromagnetic spectrum. Frequency is the number of waves passing a point per sec.

  33. High Energy - Low Above is the full electromagnetic spectrum. Humans see only a small portion of it, just the visible region.

  34. When an e- absorbs energy, it may move from a lower energy level (closer to the nucleus) to a higher energy level. The e- is not stable at these higher energy levels. They will fall back to the spot open in the lower energy level. When they fall, they give off the energy absorbed in the form of light. The amount of energy given off directly corresponds to a specific color (frequency) of light.

  35. The light given off when the e- fall back down to the lower energy levels, appear as bright lines with the rest of the spectrum black. Hydrogen atoms, give off 4 different bright lines of light. It includes: indigo at 410 nm, violet at 434 nm, blue at 486 nm, and red at 656 nm.

  36. M(g) + Ionization energy  M+1(g) + 1 e-1 Ionization Energy is the amount of energy required to remove one electron from a neutral atom.

  37. Ionization Energy 1 Hydrogen 1,312.1 Kj/mol 2 Helium 2,371.1 3 Lithium 520.1 4 Beryllium 899.1 5 Boron 780.0 6 Carbon 1,085.7 7 Nitrogen 1,401.6 8 Oxygen 1,313.0 9 Fluorine 1,679.9 10 Neon 2,079.4 11 Sodium 495.4 Kj/mol 12 Magnesium 733.0 13 Aluminum 577.0 14 Silicon 786.2 15 Phosphorus 1,011.3 16 Sulfur 999.1 17 Chlorine 1,255.2 18 Argon 1,519.6 19 Potassium 418.4 20 Calcium 589.9

  38. Ionization energy is the amount of energy required to remove one electron from a neutral atom. M(g) + Ionization energy  M+1(g) + 1 e-1

  39. Periodic Table Crossword Across 18. Mendeleyev 33. (end in –ide) Down 2. Affinity 8. Ramsay 31. Radii 34. (end in –ide) 35. Isotope

  40. Review • An atom has a nucleus with n and p+, and e- around it at specific distances from the nucleus. • An area around the nucleus where e- are found • 2 e- • they spin opposite directions • s:1, p:3, d:5, f:7 • When elements gain energy, e- jump to higher energy levels. When they fall back to normal spot, the energy comes off as light. • b (it is an alkali metal 3s1, count the e- = 11, Na) • a. O = 1s2 2s2 2p4b. V = [Ar] 4s2 3d3c. Ca = [Ar] 4s2 or 1s2 2s2 2p6 3s2 3p6 4s2 • have Mrs. W. put on board Part I: A = He, B = K, C = Ca, D = Br, E = S Part II: A = Rb, B = Br, C = Sr, D = Kr, E = Se

More Related