1. Atomic Structure Timeline
From this part of the lecture, you will be expected to…
draw the atomic models
match scientists to their experiments and discoveries
place the models in chronological order

2. Democritus (400 B.C.)
Proposed that matter was composed of tiny indivisible particles
Greek: atomos=small, solid, indestructible particles of different shapes and sizes
Not based on experimental data

3. Alchemy (next 2000 years) Mixture of science and mysticism.
Lab procedures were developed, but alchemists did not perform controlled experiments like true scientists.

4. John Dalton (1807) British Schoolteacher
based his theory on others’ experimental data
Billiard Ball Model
atom is a uniform, solid sphere

5. History of Atomic Models
5 Main Points of Dalton’s Atomic Theory
Elements are made of tiny, indivisible particles called atoms
All atoms of a given element are identical
Atoms of a given element are different than those of any other element
Atoms of one element can combine with atoms of other elements to give compounds
Atoms cannot be created, divided into smaller particles, nor destroyed in the chemical process; reactions change how atoms are grouped.

6. Problems with Dalton’s Atomic Theory?
1. Matter is composed, indivisible particles called atoms
Atoms Can Be Divided, but only in a nuclear reaction
2. All atoms of a particular element are identical
Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)!
3. Different elements have different atoms
YES!
4. Atoms combine in certain whole-number ratios
YES! Called the Law of Definite Proportions
5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements.
Yes, except for nuclear reactions that can change atoms of one element to a different element

7. Henri Becquerel (1896) Discovered radioactivity
spontaneous emission of radiation from the nucleus
Three types:
alpha () - positive
beta () - negative
gamma () - neutral

8. J. J. Thomson (1903) Cathode Ray Tube Experiments
beam of negative particles
Discovered Electrons
negative particles within the atom
Plum-pudding Model

9. J. J. Thomson (1903) Plum-pudding Model
positive sphere (pudding) with negative electrons (plums) dispersed throughout

10. Ernest Rutherford (1911) Gold Foil Experiment Discovered the nucleus
dense, positive charge in the center of the atom
Nuclear Model

11. Ernest Rutherford (1911) Nuclear Model
dense, positive nucleus surrounded by negative electrons

12. Niels Bohr (1913) Bright-Line Spectrum
tried to explain presence of specific colors in hydrogen’s spectrum
Energy Levels
electrons can only exist in specific energy states
Planetary Model

13. Niels Bohr (1913) Planetary Model
Bright-line spectrum
Planetary Model
electrons move in circular orbits within specific energy levels

14. Bohr’s Model
The electron must gain energy to move to a higher level and must lose energy to move to a lower level
Bohr said electrons act more like waves on a string than particles
Think of walking from floor to floor of the school building!

15. Erwin Schrödinger (1926) Quantum mechanics
electrons can only exist in specified energy states
Electron cloud model
orbital: region around the nucleus where e- are likely to be found

16. Electron Cloud Model (orbital)
Erwin Schrödinger (1926)
Electron Cloud Model (orbital)
dots represent probability of finding an e- not actual electrons

17. James Chadwick (1932) Discovered neutrons
neutral particles in the nucleus of an atom
Joliot-Curie Experiments
based his theory on their experimental evidence

18. revision of Rutherford’s Nuclear Model
James Chadwick (1932)
Neutron Model
revision of Rutherford’s Nuclear Model

19. Atomic Structure I. Structure of the Atom Chemical Symbols
Subatomic Particles
Atomic # and Atomic Mass
Isotopes
Average Atomic Mass
Mass of Compounds

20. Metal that forms bright blue solid compounds.
Chemical Symbols
Capitals matter!
Element symbols contain ONE capital letter followed by lowercase letter(s) if necessary.
Metal that forms bright blue solid compounds.
Co vs. CO
Poisonous gas.

21. Atomic Structure
An atom is the smallest unit of matter made of subatomic particles: protons, neutrons, and electrons

22. Subatomic Particles in a neutral atom Most of the atom’s mass.
NUCLEUS
ELECTRONS
in a neutral atom
PROTONS
NEUTRONS
NEGATIVE CHARGE
POSITIVE CHARGE
NEUTRAL CHARGE
Most of the atom’s mass.
Atomic Number
equals the # of...

23. Electrons
Bohr compared electrons to planets, saying electrons orbited the nucleus in specified and circular paths
However, an electrons exact location cannot be determined

24. Orbitals Electrons exist in energy levels called orbitals
The number of filled orbitals depends on how many electrons an atom has
Electrons occupy the orbitals that have the lowest energy (closest to the nucleus)
Four different kinds: s, p, d, f

25. Orbital Region where there is 90% probability of finding an electron.
Can’t pinpoint the location of an electron.
Density of dots represents degree of probability.

26. Orbtials s orbitals are spherical and can hold 2 electrons
p orbtials are dumbell shaped and can hold six electrons

27. Orbital
Orbitals have different shapes.

28. Valence Electrons
Electrons located in the outer most orbital are called valence electrons
These electrons determine the atom’s chemical properties and its abilities to form chemical bonds
Atoms with the same # of valence electrons have similar properties

29. Bohr Model Diagrams
Simplified energy levels using Bohr’s idea of circular orbits.
Can replace with: 3p 4n
Lithium
Atomic #: 3
Mass:
# of p: 3
# of e: 3
# of n: 4 e- e- p n
Maximum e-
Level 1 2e-
Level 2 8e-
Level 3 18e-
Level 4 32e- e-

30. Bohr Model Activity Choose a number between 1 & 18.
Find your element by the atomic number you picked.
Draw a Bohr Model diagram for your element on your marker board.
Round off the mass listed on the table and subtract the atomic # to find the # of neutrons.
Abbreviate the # of ‘p’ and ‘n’ in the nucleus.
Have a partner check your drawing.
Repeat with a new element.

31. III. Masses of Atoms Atomic Mass Mass Number Isotopes
Ch Atomic Structure
III. Masses of Atoms
Atomic Mass
Mass Number
Isotopes

32. Atomic Number Equals # of protons
Equals # of electrons in a NEUTRAL atom
Always a whole number

33. Atomic Mass atomic mass unit (amu) 1 amu = 1/12 the mass of a 12C atom
1 proton = 1 amu
1 neutron = 1 amu
1 electron=1/1822 amu
Lightest subatomic particle is the electron
1 amu = 1.67  g
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34. Mass Number Sum of the protons and neutrons in the nucleus of an atom.
© Addison-Wesley Publishing Company, Inc.
Always a whole number.

35. Mass Number # of protons + neutrons in atomic mass units (amu)
Isotopes - atoms of the same element with different masses
differ in number of neutrons
Examples: Carbon-13 & Carbon-14, Boron-10 & Boron-11
Element-mass#

37. Calculating # of neutrons
Mass # - # Protons
Example:
Aluminum
13 protons
27-13 = 14 neutrons

38. Isotopes Mass # Atomic #
Atoms of the same element with different mass numbers.
Nuclear symbol:
Mass #
Atomic #
Hyphen notation: carbon-12

39. Isotopes
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40. Isotopes Chlorine-37 atomic #: mass #: # of protons: # of electrons:
# of neutrons: 17 37 20

41. Relative Atomic Mass 12C atom = 1.992 × 10-23 g atomic mass unit (amu)
1 amu = 1/12 the mass of a 12C atom
1 p = amu 1 n = amu 1 e- = amu
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42. Average Atomic Mass Avg. Atomic Mass weighted average of all isotopes
on the Periodic Table
round to 2 decimal places
Avg.
Atomic
Mass

43. Average Atomic Mass
EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.
Avg.
Atomic
Mass
16.00
amu

44. Average Atomic Mass
EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37.
Avg.
Atomic
Mass
35.40 amu

45. Molar Mass of Compounds
Use the periodic table to find the mass of the elements you need.
Add up all the elements’ masses together appropriately.
Example: H2 gas
Mass of Hydrogen: 1.01 g/mole
Subscript says I have 2 hydrogen atoms
Mass of molecule = = 2.02 g/mole

46. Molar Mass of Compounds
Find the mass of ammonia, NH3
Find the mass of potassium sulfate, K2SO4

47. Ch. 10 - The Periodic Table II. Organization (p.286-291)
Metallic Character
Rows & Columns
Table Sections

48. A. Metallic Character
Metals
Nonmetals
Metalloids

49. B. Table Sections Representative Elements Transition Metals
Inner Transition Metals

50. B. Table Sections Overall Configuration Lanthanides - part of period 6
Actinides - part of period 7

51. C. Columns & Rows
Group (Family)
Period

52. III. Periodic Trends (p.288-291) Terms Periodic Trends Dot Diagrams
Ch The Periodic Table
III. Periodic Trends (p )
Terms
Periodic Trends
Dot Diagrams

53. A. Terms
Periodic Law
Properties of elements repeat periodically when the elements are arranged by increasing atomic number.

54. A. Terms First Ionization Energy
Valence Electrons
e- in the outermost energy level
Atomic Radius
First Ionization Energy
energy required to remove an e- from a neutral atom

55. B. Periodic Trends
Atomic Radius
Increases to the LEFT and DOWN.

56. B. Periodic Trends Increases to the RIGHT and UP.
First Ionization Energy
Increases to the RIGHT and UP.

57. Be or Ba Ca or Br Ba Ca B. Periodic Trends
Which atom has the larger radius?
Be or Ba
Ca or Br Ba Ca

58. N or Bi Ba or Ne N Ne B. Periodic Trends
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne N Ne

59. B. Periodic Trends Group # = # of valence e- (except He)
Families have similar reactivity.
Period # = # of energy levels 1A 2A
3A 4A 5A 6A 7A 8A

60. C. Dot Diagrams
Dots represent the valence e-.
EX: Sodium
EX: Chlorine