Chapter 4. Atomic Structure

Transcription

1 Chapter 4 Atomic Structure

2 History: Early 1800 s Lots of introductory work with electricity Atoms are electrical in nature Have parts with + charge Have parts with charge Technology was limited until late 1800 s, early 1900 s

3 Cathode Ray Tube Tube containing two metal electrodes/plates Under VERY low pressure (vacuum) Glows / green beam (cathode ray)

4 So what are Cathode Rays??? John Thompson (1897) Able to deflect a cathode ray using an electric field Beam attracted to + plate, deflected from plate THUS- cathode ray is composed of charged particles DISCOVERY OF ELECTRONS!!!

5 John Thompson, cont d DISCOVERED ELECTRONS Also used magnetic fields to deflect cathode ray Allowed him to calculate ratio of electron mass to charge 1906 Nobel Prize

6 Canal Rays Eugen Goldstein (1886) Used a discharge tube with holes in cathode + ions shot in opposite direction of cathode rays (electrons) == CANAL RAYS Some went through the holes repeated in presence of magnetic field--- found the + particles had varying massess

7 Millikan Oil-Drop Experiment Robert A. Millikan-(1909) Discovered charge of an electron Placed electrically charged oil drops in an electric field Oil drops sprayed in upper chamber Some gain charge (static electricity) Calculate mass by measuring how fast oil drops fall though hole (gravity)

8 Millikan Oil-Drop Experiment, cont d Then Millikan adjusted voltage on plates in bottom chamber - particles attracted to upper plate to offset pull of gravity---suspended drop in space Calculate mass: charge ratio Know mass from first exps, so can calculate charge! Smallest change between two drops taken to be charge of an electron 1923 Nobel Prize

9 Millikan Oil-Drop Experiment, cont d Mass of an electron: 9.1X10-28 g Charge of an electron: - 1

10 Radioactivity Antione Bequerel Placed a sample of uranium on covered photographic paper (NO light exposure) Fogged the photographic paper!! Student = Marie (Sklodowska) Curie called it radioactivity Bequerel, Pierre Curie, Marie Curie shared Nobel Prize in Physics Marie Curie won Nobel Prize in chemistry for additional work on radioactivity

11 Radioactivity, cont d 3 types of radioactivity (Ernest Rutherford) Alpha (α) Beta (β) Gamma (γ)

12 Radioactivity, cont d Alpha Particles: Deflected magnetic field like + particles Mass = 4 X that of hydrogen atom Charge: 2 X that of an electron, but + charged

13 Radioactivity, cont d Beta particles: Electrons / cathode rays Charge = -1 Mass = 1/1837 of a hydrogen atom Gamma Particles: NOT deflected in a magnetic fields Very high energy

14 Radioactivity, cont d particle Mass (u) charge Penetrating ability alpha 4 +2 Low (paper shield) beta 1/ Intermediate (plastic shield) gamma 0 0 High (lead shield)

15 Gold Foil Experiment Ernest Rutherford Bombard gold foil with alpha particles

16 Gold Foil Experiment Many α particles passed through gold foil as expected, but SOME DEFLECTED Concluded that + charge and most of mass of an atom is contained in a tiny core (NUCLEUS)

17 Models of Atomic Structure By early 1910 s, scientists knew about electrons and protons, and that the majority of the mass (an + charge) was in the nucleus Began to develop models of atomic structure

18 Rutherford s Plum Pudding Model Atoms are like plum pudding Lots of empty space (pudding), but a central core (plum) every so often, where most of the mass is located Nucleus: concentrated, tiny volume where vast majority of the mass is located, + charges there Electrons: surround nucleus, but occupy almost the entire volume of an atom

19 Rutherford s Plum Pudding Model or Football Analogy Football Stadium

20 Subatomic Particles 3 main ones (there are others, but that s for physics..) Neutrons Protons Electrons

21 Electrons JJ Thompson, RA Millikan 1st discovered late 1800 s, early 1900 s cathode rays Negatively (-) charged particles Mass = X g Charge = -1 = X10-19 coulombs Occupy majority of space in an atom

22 Protons Eugen Goldstein (1886) canal rays Mass = X10-24 g 1800 X as heavy as an electron! Charge = +1 = X10-19 coulombs Located in tiny, dense nucleus

23 Neutrons James Chadwick (1932) Provide mass, but no charge Mass = X g (nearly the same as a proton) Located in the nucleus

24 Nuclear Symbols X = element symbol Z = mass # = # protons + # neutrons A = atomic # = # protons (gives an element its identity!!) So # neutrons = Z-A Z A X charge

25 Atomic # # protons Gives an element it s identity If 6 P.C 7 P.N 1 P.H 53 P I NOTE: # above element symbol in periodic table = # protons = Atomic #

26 Electrons and atom charge If # P = # electrons (e-) Then atom is NOT CHARGED Charge = #P - #e- If #P > #e- Then + charged (cation) If # P < # e- Then charged (anion)

27 ISOTOPES Atoms with same # Protons, but different # neutrons Example 1 H, 2 H, 3 H All are H (have 1 proton) Z is different!!! Isotope #P #N 1 H = (1+0) 2 H = (1+1) 3 H = (1 +2)

28 How many P, N, e-?? Au # P = 79 # N = (197-79) = 118 # e- = Na +1 # P = 11 # N = (23-11) = 12 #e- = = #p - #e- +1 = 11- # e-

29 Write the nuclear symbol for An atom containing 18 e-, 16 P, and 17 N. What element is it?? Atomic # = 16 P = Sulfur (S) Z = #N + # P = 16+17=33 Charge = #p -#e- = 16 (18) = S 16

30 Which of the following represent isotopes of the same element? 16 8 X 16 7 X 14 7 X 14 6 X 12 6 X Look for same atomic #: (same # protons) 16 7 X and 14 7 X both have 7 protons, so are isotopes of Nitrogen X and 12 6 X both have 6 protons, so are isotopes of C.

31 The Periodic Table LOTS on INFORMATION!!! Columns = Groups = Families = up and down, # 1-18 Main groups : 1, 2, Transition metals: 3-10 Inner Transition metals: at bottom of chart Members of the same group have similar physical and chemical properties Rows = Periods

32 Main groups Transition metals Inner transition metals

33 Metals Gps 1,2,3-10, some in 14 and 15 (blue) Maleable, ductile, conductive, Usually form cations

34 Metaloids Green in picture Elements on step between metals and non-metals Intermediate properties (semi-conductors)

35 Non-metals Rest of chart (yellow) Nonconductive usually form anions

36 Gp 1- alkalai metals Very soft not found in nature in uncharged atom form (find as cation) VERY reactive with nonmetals and water

37 Gp 2- alkali earth metals Harder, more dense vs gp 1 Less reactive than gp 1 All react with oxygen to form metal oxides

38 Gp 17- Halogens salt-producing React with metals to form salts Exist as diatomic molecules in nature (F 2, Cl 2, etc.)

39 Gp 18- Noble Gases Colorless gases Exist as single atoms Very Unreactive Some limited reactivity of Xe, Kr, Rn

40 What is it about elements that explain the reactivity and other properties? -arranged by atomic # (# protons), for UNCHARGED elements (#p = #e-) So as #P increases, #e- increases Thus reactivity based on ARRANGEMENT OF ELECTRONS around the nucleus

41 Electron Organization Look at how electrons are positioned around the nucleus Not in any one place, but fill different regions of space around the nucleus

42 Continuous Line Spectra When white light (sun) is separated through a prism, see that colors blend into one another Spectrum is continuous Each color represents a different wavelength of light

43 Emission Line Spectra Take pure element in gas phase and put in cathode ray tube, apply high voltage Atoms excited as they absorb energy Atoms emit energy as LIGHT

44 Emission Line Spectra When pass light through prism, get bands of color at specific wavelengths NOT continuous like white light! Each element has a signature ELS

45 Bohr Model of Atomic Structure (1913) Based on Emission line spectra of Hydrogen Electrons can have only a specific amount of energy = QUANTA Energy values = energy levels Quanta = difference between any two energy levels

46 Bohr Model- Cont d 6 to 2, 5 to 2, 4 to 2 and 3 to 2 transitions emit energy of visible wavelengths Analogy- you can stand on the rungs of a ladder, but not in between

47 Bohr Model- Planetary Model Energy levels are in rings/ orbitals around the nucleus Each orbital (ring) is assigned an integer value n = principle quantum # Radius of orbit increases as n increases n=1 is the lowest possible level, where e- in H usually resides

48 Bohr Model- Planetary Model Ground state = lowest energy level an e- occupies in the normal state Excited state = when an e- is in an energy level greater than the ground state Takes energy to move e- from ground to excited state (move away from + pull of nucleus) Excited state is unstable, e- wants to go back to normal state e- releases energy (light) as if falls back to the ground state This is the light in the Emission Line Spectra

49 Bohr Model- Planetary Model Bohr won 1922 Nobel Prize for model For atoms with more than 1 electron, the Bohr model suggests that the energy levels can hold only a certain # of e- #e- = 2n 2 (n= orbital #) N = #e- =

50 Bohr Model- Planetary Model BUT the outermost energy level can only hold 8 e-. 1 st row on PT 2 nd row on PT 3 rd row on PT 1 st level now filled 2n 2 = 2(1) 2 = 2 2 nd level now filled 2(2) 2 = 8e- 3 rd level maxed with 8e-; but can hold 18

51 Bohr Model- Planetary Model After this, things get more complicated Bohr s model begins to fail.

52 Erwin Schrodinger s Treat electrons as both waves and particles (Bohr treated only as particles..) Devised UGLY mathematical equations that state the PROBABILITY of finding an electron in a given region around the nucleus Cannot tell exact location, just the likelihood of where an electron will be located!

53 Quantum Mechanics Schrodinger model subdivides Bohrs energy levels Orbitals = regions of space (replaced Bohr s planetary orbits) Camera analogy- if you could look at a timelaspe photo of an electron, the area exposed is the orbital

54 Quantum Mechanics Each orbital can hold 2 electrons Some energy levels hold more than 1 orbital # orbitals = n (row #, principle quantum #) n # orbitals names 1 1 s 2 2 s, p 3 3 s, p, d 4 4 s, p, d, f

55 Quantum Mechanics Each orbital has a distinct shape S = sphere P=dumb bell D, f = very complicated

56 Quantum Mechanics

57 Inverted Triangle Apartment Building Floor = n Apt = orbital (s, p, d, f) s = 1 bedroom = 2 e- p = 3 bedroom = 6 e- d = 5 bedroom = 10 e- f = 7 bedroom = 14 e- (2 e- per room)

58 Fits Periodic Table!

59 Order of Filling 10 e- 14 e- Follow atomic # on Periodic Table 6e- 2 e-

60 Order of Filling

61

62 Practice: Give the electron configurations for each of the following elements Na O Mn Still have only 8e- max in outer most shell??

63 Don t forget the f-block! Ba La Ce Lu Hf Hg

64 Identify the following elements given the electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 1s 2 2s 2 2p 6 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 3

65 Octet Rule No more than 8 e- in the outermost shell!! outershell e- = valence electrons Outershell = greatest n!!! General Rule: Atoms combine with other atoms (form bonds) in order to have a COMPLETE octet (8 outershell electrons) Share e- (covalent bond) Gain/lose electrons (forming IONS) Want to have electron configuration of Nobel gases! (fill a p6 sublevel!)

66 IONS Atom that has either a + or a charge, obtained by either by gaining an e- (extra vs +) so charge anion Lose an e- (less vs +) So + charge.cation

67 Anions Gained an e-.now more e- than protons(+) Net charge Ex Oxygen Atom : 1s 2 2s 2 2p 4 Needs 2 e- to fill up p level (be 2p6), like Neon So will GAIN 2 e- Ion: O 2- : 1s 2 2s 2 2p 6 (put charge in upper right)

68 Anions, cont d Chlorine: atom: 1s 2 2s 2 2p 6 3s 2 3p 5 needs 1 e- to get configuration like Argon (..3p 6 ) (isoelectric) Picks up 1 e- Cl - : 1s 2 2s 2 2p 6 3s 2 3p 6

69 Cations Lose an e-.now fewer e- than protons(+) Net + charge Ex Sodium Atom : 1s 2 2s 2 2p 2 3s 1 One electron, all alone on 3 rd floor jumps off So will LOSE 1 e- Ion: Na + : 1s 2 2s 2 2p 6 (put charge in upper right)

70 Cations, cont d Calcium: Atom: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Lose the two outer shell (n=4) electrons Ion: Ca 2+ : 1s 2 2s 2 2p 6 3s 2 3p 6 (like Argon)

71 Ions in Summary Gp 1: (ns 1 ) : lose 1 e- Gp 2: (ns 2 ) : lose 2 e- Gp 16: (ns 2 np 4 ) gain 2 e- Gp 17: (ns 2 np5) gain 1 e- +1 cations +2 cations -2 anions -1 anion

72 Periodic Trends Certain trends based on electron configurations / organizations Atom size Parent atom vs Ion size Reactivity Electronegativity

73 Atom Size

74 Atom Size What is happening? 1. More orbitals (energy levels, apt floors), the larger the atom So INCREASE in size DOWN a group Li < Na < K < Rb < Cs < Fr n = 1 < 2 < 3 < 4 < 5 < 6

75 Atom Size

76 Atom Size 2. As e- fill in the same energy level /floor (across from left to right) get more protons (+) in the nucleus while e- filling same level So more + to attract e- in towards nucleus Trend: DECREASE atom size ACROSS (L to R) a row Li > Be > B > C > N > O > F > Ne

77 Atom Size

78 Example Arrange the atoms by increasing size (smallest to largest): Mg Al Na S Ar Find location in PT: note all in n = 3 so Left = largest and right = smallest Ar < S < Al < Mg < Na

79 Ion size vs. Atom size What happens to radius when an atom gains an e-? When it loses an e-?

80 Anions: Ion size vs Parent Atom add e-: now # e- > # P (+) each electron now has less than a full +1 pull from the nucleus e- not held as tightly Radius increases / larger in anion vs parent atom

81 Ion size vs Parent Atom Cations: lose e-: now # e- < # P (+) each electron now has more than a full +1 pull from the nucleus e- held more tightly Radius decreases / smaller in cation vs parent atom

82 Trends in Reactivity Metals: the larger the atom, the farther away the valence e- is from the + nucleus Lose the electron more easily (more reactive) as size increases Ex: Li vs Na vs K : which is more reactive with water? K is largest, so more reactive

83 Trends in Reactivity Non-metals: - want to gain e- (Opposite trend of metals) - Closer the new e- will be to the nucleus, the greater the attraction or pull to get that electron to the new atom - So Reactivity of non-metals increases as atomic radius decreases

84 Electronegativity desire to gain an electron F is most EN atom All increases to F Fr is Least EN atom

85 Why is F the most EN atom? F REALLY wants an additional electron to fill its p level The F p-level (n=2) is the closest p-level to the nucleus, so there is a VERY strong, unshielded nuclear pull So an F atom will pull as many electrons to it as it can (more on this in the next chapter )

86 Chapter Objectives 1. Know how electricity played an important role in unraveling the structure of the atom. Define: cathode ray tube, cathode rays, canal rays. Be able to describe the role cathode ray tubes played in the discovery of electrons and protons.

87 2. Be familiar with the key experiments performed by Thomson (properties of cathode rays), Goldstein (canal rays), Milliken (charge and mass of electrons), and Chadwick (neutrons). Be able to describe Thomson s plum pudding model of the atom. 3. Know the brief history of radioactivity and how work with radioactive elements led to an understanding of the existence of subatomic particles. Be familiar with the accidental discovery of radioactivity by Becquerel and the contributions of the Curies.

88 4. Know three types of radioactivity: alpha, beta, and gamma emissions as defined by Rutherford. Be able to distinguish them by mass and charge. 5. Discuss Rutherford s gold foil experiment, which led to the discovery of the atom s nucleus. Contrast Rutherford s atomic model with Thomson s model and current models.

89 6. Be able to define and understand the following terms: atomic number, mass number, and isotope. Know how to find the atomic number on the periodic table and know its significance in identifying an element. Understand that the mass number of an isotope can not be found on the periodic table. 7. Be able to use the atomic number and mass number of any element to determine the number of electrons, protons, and neutrons in a neutral atom. Also, given the number of protons and neutrons, be able to determine atomic number and mass number and be able to draw the notation for an isotope.

90 8 8. Know the difference between continuous spectra and line spectra. Be able to explain how the line spectra of hydrogen can be explained by the Bohr model of the atom. Define the terms ground state and excited state of an electron as defined by the Bohr model. 9. Be able to explain the Schroedinger model of the atom based on probabilities of electron location and pathway. Define the term electron orbital and know the shape of the s and p orbitals. Know how electron configurations for neutral atoms are derived according to the Schroedinger model using s, p, d, and f orbitals.

91 10. Be familiar with the organization of the periodic table. Know the difference between a group (column) and a period (row). Know and locate the Group A elements. Be able to derive electron configurations for main group elements. Given an electron configuration, be able to locate the element on the periodic table. 11. Using the periodic table be able to point out the alkali metals, alkaline earth metals, halogens, and noble gases. Be able to find transition metals and inner transition metals.

92 12. Know area on the periodic table that correspond to the filling of s, p, d, and f orbitals. Be able to use the periodic table to do electron configurations. 13. Define valence electrons. Be able to relate group A numbers to number of valence electrons as derived from electron configurations. Know the significance of the number of valence electron s of group VIIIA elements (ie, noble gases) and how this might relate to chemical reactivity of other main group elements.

93 14. Be able to distinguish general physical properties of metals and nonmetals; eg, malleability, ductility, ability to conduct heat and current. Be familiar with general chemical properties of metals and nonmetals; eg, metals tend to give up electrons to form positive ions when reacting with nonmetals, nonmetals tend to gain electrons when reacting with metals. Understand why this might occur. 15. Predict trends in sizes of atoms and trends in chemical reactivity going down groups and going across periods of the periodic table. Be able to understand reasons for these trends.