1. Chapter 4 Atomic Structure
HONORS CHEMISTRY
Chapter 4
Atomic Structure

2. History of the Atomic Theory

3. DEMOCRITUS (400 BC) 1st atomic theory
“World is made of empty space & tiny particles called ‘atoms’.”
Atomos - Greek for indivisible
Smallest poss. particles of matter
“Diff. types of atoms for every type of matter”
General & not supported by experiment
Not accepted - Contradicted Aristotle

4. ARISTOTLE “Matter is continuous” - not made of smaller particles
“Hyle”
Accepted until 17th Century

5. Isaac Newton & Robt. Boyle
Published articles on belief in atomic nature of elements
No Proof
Attempted explanations, no predictions

6. John Dalton Logical hypothesis on existence of atoms
Studied & explained work of other scientists
Lavoisier - “In a closed syst., the mass of the reactants = the mass of the products”
LAW OF CONSERVATION OF MASS
Proust - “Specific substs. always contain elems. in the same ratio by mass”
LAW OF DEFINITE PROPORTIONS

7. Dalton’s Atomic Theory
Basis of modern atomic theory
1st atomic theory based on experimental evidence

8. Dalton’s Atomic Theory
Four important statements:
1. All matter is composed of indivisible atoms.
2. All atoms of the same elem. are identical.
3. Atoms of diff. elems. are not alike.
4. Atoms unite in simple ratios to form compounds.

9. Dalton’s Atomic Theory
Explains Law of Cons. of Mass
atoms are rearranged in a chem. rxn.
Explains Law of Definite Proportions
Not exactly correct

10. DALTON ALSO STATED: Law of Multiple Proportions
The ratio of masses of one element that combines w/ a constant mass of another elem. can be expressed in small whole numbers.

11. Other Scientists
Gay Lussac - “Under constant conditions, the volumes of reacting gases & gaseous products are in the ratio of small whole numbers.”
Avogadro explained this - “Equal volumes of gases, under the same conditions, contain the same # of molecs.”

12. Cathode Ray Tube Tube w/ charged metal electrodes in ea. end
Anode - Positive electrode
Cathode - Neg. electrode
Rays in tube seemed to travel from cathode to anode
Cathode Rays

13. J. J. Thomson Discovered electrons using cathode ray tube
Determined charge to mass ratio of e-

14. Robt. Millikan Oil Drop Experiment Measured the charge on an e-
std. unit of neg. charge (-1)
e- mass is 1/1837 mass of a H atom

15. J. J. Thomson Discovered electrons using cathode ray tube
Determined charge to mass ratio of e-
Discovered the proton using a modified cathode ray tube
same amt. of chg. as e- but positive
std. unit of (+) chg. = +1
Calculated mass of p+ (1836 X mass of e-)

16. Lord Rutherford
predicted 3rd particle

17. James Chadwick Discovered the neutron
high energy particle w/ no chg. & approx. same mass as p+

18. Dalton’s Theory was revised.
Subatomic particles had been discovered.

19. J. J. Thomson Discovered ISOTOPES
atoms of the same elem. that differ in mass
have same # of p+’s, but diff. # of no’s

20. Henry Mosely
using x-rays, found the number of p+’s in the nucleus of an atom is always the same for a given element
Atomic Number (Z) - # of p+’s in the nucleus
# p+’s = # e-’s in a neutral atom

21. The number of p+’s determines the identity of the elem
The number of p+’s determines the identity of the elem. and the # of no’s determines the particular isotope of the elem.

22. Dalton’s Theory revised again
Not all atoms of the same element are exactly alike.
Atoms are NOT indivisible!

23. Nucleons - particles that make up the atomic nucleus
Nuclide - a particular type of atom containing a definite # of p+’s & no’s
Nucleons - particles that make up the atomic nucleus
p+’s & no’s
Mass Number (A) - total # of nucleons in an atom
Number of no’s = A - Z
(mass # - atomic #)

24. Rutherford’s Gold Foil Experiment
led by Lord Rutherford, assisted by a team of physicists (Niels Bohr, Hans Geiger, & Ernest Marsden)
Procedure: shot (+) charged subatomic very thin sheet of gold foil.

25. Rutherford’s Gold Foil Experiment
Observations
1. Most particles passed straight thru foil.
2. Few particles were large angles.
3. Very few (1 in 8000) bounced almost straight back.
Conclusions:
1. Most of the atom is empty space.
2. + particles came close to “core” of atom which must have a + charge.
3. + particles almost hit core straight on.

26. Rutherford’s Gold Foil Experiment
Overall Conclusion
Atoms consist of (+) charged nucleus surrounded by e-’s

27. Diameter of an atom ~ 100-500 pm
Radii of nuclei of atoms vary between 1.2x10-3 and 7.5 x 10-3 pm
Nucleus is ~ 1 trillionth the vol. of the atom.

28. Henri Becquerel
1896 w/ Marie & Pierre Curie discovered Radioactive Substs.
When brought near charged electroscope, leaves become discharged

29. Radioactivity
Phenomenon of rays being produced spontaneously by unstable atomic nuclei
mixture of particles & energy given off by nuclei during spontaneous nuclear decay
amt. of energy very large - E = mc2
Half-Life - length of time needed for 1/2 an amt. of a radioactive nuclide to disintegrate.

30. Nuclear Force - force which holds p+’s and No’s together in nucleus
effective over very short distance

31. Scientists agree on:
1. Nucleons have a prop. that corresponds to spinning on an axis.
2. e-’s don’t exist in nucleus, but can be emitted from nucleus.

32. Star Trek Science

33. Subatomic particles - particles composing atoms 2 broad classes
Leptons - (light particles) - truly elementary
best known: electrons
Hadrons - appear to be made of smaller particles
best known: neutrons & protons

34. For every particle, a mirror image particle called an Antiparticle is believed to exist
antielectron is a positron
When particle & its antiparticle collide, both are destroyed & energy is produced.

35. Several Leptons electrons neutrinos - essentially massless
Muon - much more massive than e-
Tau - much more massive than e-

36. Hadrons divided into 2 groups
Mesons
Baryons
p+’s and no’s are baryons
Both made of Quarks
6 kinds of quarks
up, down, charmed, strange, top, bottom
ea. quark comes in 3 different “colors” - red, blue, green
ea. quark has antimatter counterpart - antiquark

37. If structure of nucleus is unstable, ejects particle or energy to become stable
Some nuclei naturally unstable, some artificially unstable

38. 3 forms of radiation from naturally radioactive nuclei
2 are particles
Alpha particle - 42He - helium nucleus a
Beta particle - 0-1e - an e- b
1 Form is energy
Gamma Rays- g - very high energy x-rays

39. Short hand to represent particles
Upper rt. “corner” - charge on ion
Lower rt. - # of atoms in formula unit
Upper left - mass #
Lower left - charge on nucleus or particle

40. Examples 3216S - Sulfur nucleus or atom 0-1e - electron
42He - alpha particle (helium nucleus)

41. Scientists create radioactive nuclides by bombarding stable nuclei w/ accelerated particles or w/ neutrons in nuclear reactor
Decay by emitting natural radiation & other methods.

42. Planetary Atomic Model
Proposed by Rutherford and Bohr
e-’s “orbit” around nucleus
H atom similar to solar syst. w/ 1 planet

43. Bohr exposed atoms to radiant energy
atoms absorb some energy
Excited Atoms
Excited atoms & molecs. produce energy changes
unique & can be used to identify particle
absorb and emit radiant energy

44. SPECTROSCOPY
Method of studying substs. exposed to exciting energy

45. SPECTRUM
Pattern of radiant energy studied in spectroscopy

46. ELECTROMAGNETIC (RADIANT) ENERGY
Visible light, radio, ultraviolet, infrared, etc.
Travels in waves
variations in elect. & magnetic fields taking place in regular repeating fashion
Frenquency - n- # of wave peaks that occur in a unit of time
meas. in hertz (Hz) = 1 peak or cycle per sec.

47. ELECTROMAGNETIC (RADIANT) ENERGY
speed of light (c)
3.00 x 108 m/s in vacuum
Wavelength - l - physical dist. betw. peaks
Related by c = ln
Amplitude - maximum displacement from zero

48. Excited atoms lose energy
Energy emitted by gaseous atoms can be spread into a spectrum.
Emission Spectrum - shows l’s of light given off by excited atoms
Absorption Spectrum - have lines missing from continuous spectrum showing which l’s of light have been absorbed

49. Lines missing in absorption spectrum are the same as lines shown in emission spectrum
unique to ea. elem.
used to identify elems.

50. Electromagnetic Spectrum
Radio waves - longest l’s
Gamma waves - shortest l’s
Visible light????

51. Planetary Model of Atom
Developed by Bohr to explain H spectrum
Used Quantum Theory - theory of energy emission stated by Max Planck

52. Quantum Theory
Planck assumed energy was emitted in packets or Quanta - not continuously
Quanta of radiant energy - Photons
Amt. of energy given off is directly related to frequency of light emitted
E = hn
h = Planck’s constant (6.63 x J/Hz)
E = energy in a quanta

53. Planck’s Hypothesis
Energy is given off in quanta instead of continuously

54. Bohr
“Absorption of light by definite l’s corresp. to definite changes in energy of e-.

55. Reasoned: 1. Orbits of e- around nucleus must have definite diameter.
2. e-’s can occupy only certain orbits.
3. Only orbits allowed - those w/ diff. in energy = energy absorbed when atom was excited
\ e-’s can absorb quantum of energy & move to larger orbit
Since quantum represents certain amt. of energy, next orbit must be definite dist. from 1st orbit.

56. When e- drops to lower orbit, energy is emitted (light).
Orbit represents definite energy \ def. amt. of energy is given off.
Ground State of e- is its smallest (lowest) orbit

57. Today’s atomic model differs from Bohr’s
Major diff. - e-’s do not move around nucleus like planets orbit sun
Idea of energy levels still basis of modern atomic theory.

58. Average Atomic Mass Mass of single atom too sm. to work with.
use mass of large group of atoms
Chemists meas. mass of single atoms in Atomic Mass Units.(amu or u)
C-12 nuclide chosen at std. - all other atoms compared to it
one C-12 atom defined as having a mass of amu

59. Average Atomic Mass
An Atomic Mass Unit - 1/12 the mass of a C-12 nuclide
e- = x 10-28g = u
p+ = x 10-24g = u
no = x 10-24g = u
Number in Periodic table based on “average atom” of the elem.
Ave. atomic mass used for calculations

60. Average Atomic Mass 2 ways of determining masses for atoms of elem:
1. Experimentation & calculation
2. Mass Spectrometer - meas. masses & relative amts. of nuclides for all isotopes of an elem.

61. Average Atomic Mass
If masses of isotopes & relative amts. are known, ave. atomic mass can be calculated
Atomic Mass of the elem.
Must use weighted average to find ave. atomic mass.