1. ATOMIC CONCEPTS

2. EVOLUTION OF MODERN ATOMIC THEORY
 Chemistry and the Greeks
 Matter composed of the four elements
 Earth, Air, Fire, Water

3. DALTON’S ATOMIC THEORY (1803)
 Atom – The basic unit of matter.
 All elements are composed of indivisible atoms.
 All atoms of a given element are identical.
 Atoms of different elements are different (Have different
masses)
 Compounds are formed by the combination of different
elements.
Dalton’s Atomic Model

4. J.J. THOMPSON AND THE ELECTRON
 Experimental studies of the atom soon showed that
it is NOT indivisible but in fact made up of smaller
particles.
 J.J. Thompson used a cathode ray tube to discover one
such particle.

5. J.J. THOMPSON AND THE ELECTRON
 Beam emitted by the cathode would respond in
different ways when exposed to a magnetic field.
When the positive end of When the negative end of
a magnet was held near a magnet was held near
the beam it would move the beam it would move
toward the magnet. away from the magnet.
(attracted) (repelled)
 Thompson surmised that the beam was composed
of negatively charged particles which he called
Electrons

6. J.J. THOMPSON AND THE ELECTRON
 Electron (e-) – negatively charged subatomic
particles that part of an atom.
 Thompson’s “Plum Pudding Model” (1897) of the atom
visualized electrons as being embedded within the
atom.
The mass of the rest of the was evenly
distributed and positively charged.

7. RUTHERFORD’S GOLD FOIL EXPERIMENT
 Used alpha particles directed at a thin piece of gold
foil which led to the discovery of the Nucleus.
Alpha particles are
positively charged
particles that are much
smaller than the atom.

8. RUTHERFORD’S GOLD FOIL EXPERIMENT
 If Thompson’s plum pudding model was correct the alpha particles
would simply pass through the foil with just a few being slightly
deflected.
 Rutherford discovered that while most of the alpha particles did
indeed pass through the foil, some were greatly deflected, and some
even bounced back.

9. RUTHERFORD’S GOLD FOIL EXPERIMENT
 Rutherford theorized that the atom must then be
composed of mostly empty space, with a dense
positively charged core that he called the nucleus.
Rutherford’s Atomic Model
(1909)

10. PROTONS AND NEUTRONS
 Atoms electrically neutral so there must be particles
to offset the electro-magnetic charge of the
negative electrons.
 Protons (p+) – tiny positively charged particles
found within the nucleus of the atom.
 Neutrons (n0) – tiny particles found within the
nucleus of the atom having no electro-magnetic
charge

11. MODERN ATOMIC THEORY (BOHR MODEL)
 Niels Bohr’s “Planetary Model” (1913) of the atom.
 Nucleus (protons and neutrons) in the center.
 Electrons shown in concentric circles or shells around
the nucleus.
 Designated by letters K, L, M, N, O, P, Q or the numbers 1
through 7.

12. MODERN ATOM THEORY
(WAVE MECHANICAL MODEL)
 Dual Nature of Matter
 Energy viewed as waves and matter as particles.
 Electrons exhibit a dual nature in which they not only
have mass but possess wavelike properties as well.
 Wave Mechanical Model
 Dense centrally located positive nucleus
 Electron no longer pictured in fixed orbits but as regions
of differing energy levels where they are most likely to
be found called Orbitals

13. SUBATOMIC PARTICLES
 An atom is the smallest unit of an element. It
consists of three major particles.
a.m.u. = atomic mass unit
a.m.u. = 1/12 the mass of a C-12 atom, or, 1.66x10-24 grams.

14. ATOMIC SYMBOLS
 Written in a shortened form as….
Atomic Mass rounded to the
closest whole number.

15. PRACTICE
 Write a short form for each atomic symbol.

16. DIFFERENCES BETWEEN ATOMS
 Atomic Number: The number of protons in the nucleus of an
atom.
 It is also the number of electrons in an electrically neutral atom
 Atomic Mass (Mass Number):

17. SAMPLE PROBLEM
 Find the number of neutrons in an atom of ?

18. SAMPLE PROBLEM
 Find the number of neutrons in an atom of ?

19. DIFFERENCES BETWEEN ATOMS
 Why are there fractional mass numbers (atomic
masses) on the periodic table?

20. DIFFERENCES BETWEEN ATOMS
 Why are there fractional mass numbers (atomic
masses) on the periodic table?
 Answer: Because of the existence of Isotopes

21. ISOTOPES
 Atoms having the same number of protons but
different number of neutrons.
 Example;
Average ≈ 22.98977
*Average is based on various isotopic masses and the
relative abundances of each.

22. SAMPLE PROBLEM
 Atomic masses can be calculated from the mass and the
abundance of naturally occurring isotopes. Carbon has two
naturally occurring stable isotopes. Most carbon atoms
(99.89%) are C-12, while the remaining 1.108% are C-13.
What is the atomic number of carbon?

23. SAMPLE PROBLEM
 Atomic masses can be calculated from the mass and the
abundance of naturally occurring isotopes. Carbon has two
naturally occurring stable isotopes. Most carbon atoms
(99.89%) are C-12, while the remaining 1.108% are C-13.
What is the atomic number of carbon?

24. SAMPLE PROBLEM
 Element X has two naturally occurring isotopes. If 72.0% of
the element has an isotopic mass of 84.9 amu and the 28.0%
has an isotopic mass of 87.0 amu, the average atomic mass
of element X is?

25. SAMPLE PROBLEM
 The average isotopic mass of chlorine is 35.5 amu.
Which mixture of isotopes (shown as percents)
produces this mass?
1. 50% C-12 and 50% C-13
2. 50% Cl-35 and 50% Cl-37
3. 75% Cl-35 and 25% Cl-37
4. 75% C-12 and 25% C-13

26. DIFFERENCES BETWEEN ATOMS

27. DIFFERENCES BETWEEN ATOMS

28. IONS
 Atoms of the same element having the same # of
protons, but different # of electrons.
 No longer electrically neutral, Ions are charged particles
 Example;

29. IONS
 Atoms of the same element having the same # of
protons, but different # of electrons.
 No longer electrically neutral, Ions are charged particles
 Example;
6 p+
6 e-

30. IONS
 Atoms of the same element having the same # of
protons, but different # of electrons.
 No longer electrically neutral, Ions are charged particles
 Example;
6 p+ 6 p+ 6 p+ 6 p+
6 e- 10 e- 4 e- 2 e-

31. QUIZ
 Identify the number of protons, neutrons, and
electrons for each element.

32. ATOMIC MODELS

33. ATOMIC MODELS
 Bohr Model
 K-shell = max. of 2 e-
 L-shell = max. of 8 e-
 M-shell = max of 18 e-
 N-shell = max. of 32 e-

34. BOHR MODEL

35. BOHR MODEL

36. BOHR MODEL

37. BOHR MODEL PRACTICE

38. BOHR MODEL PRACTICE

39. QUIZ
 Copy the picture and label the following, Atomic Mass,
Atomic Number, Electron Configuration, Selected
Oxidation State.
 Atomic ___________indicates the number of
__________ within the nucleus of the atom.
 Atomic __________ is equal to the number of
_________ plus the number of __________ within the
nucleus of the atom.
 Selected Oxidation states indicate the most common
__________ for a particular element.
 C-12, C-13 are examples of _____________.

40. ATOMIC MODELS –
IMPORTANT DEFINITIONS
 Principal Energy Level: Region around the nucleus in which electron
can be found.
 Designated by letters K, L, M, N, O, P, Q or the numbers 1 through 7.
 The closer to the nucleus the lower the energy.
 Quanta: Small amount of energy that an electron can release or absorb
as it moves through principle energy levels.
 Ground State: All electrons fill lowest energy levels before higher
energy levels are filled.
 Excited State: one or more electrons absorb energy (quanta) and
occupy a higher principle energy level than
 Spectral Lines: As electrons at principle higher energy levels (excited
state) fall back to their normal principle energy levels (ground state) they
emit that extra energy in the form of light.
 Visible Spectrum – ROY G BIV

41. EMISSION SPECTRA

42. EMISSION SPECTRA OF THE SUN

43. ORBITAL MODEL
(WAVE-MECHANICAL MODEL)
 Principal Energy (Quantum) Level represents the
level in which electrons are found.
 These correlate with period number on the periodic
table.
 These are also your K, L, M, N….
 Sublevels are represented by s, p , d, f
 Number of sublevels = Principal Energy Level
 Principal Energy Level 1 has one sub level (s)
 Principal Energy Level 2 has two sub level (s,p)
 An Orbital is an exact region in which electrons
within a principal energy level are most likely to be
found.
 The maximum number of electron in any orbital is 2

44. Note: The principal energy level is represented by n. The
number of Orbitals per level would be n2, and the maximum
number of electrons per level would be 2n2.

45. SUBLEVELS

46. WRITING ELECTRON CONFIGURATIONS
I
n
c
r
e
a
s
i
n
g
E
n
e
r
g
y
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

47. WRITING ELECTRON CONFIGURATIONS

48. EXAMPLES OF ELECTRON CONFIGURATIONS
Na
S
Kr

49. WRITE THE ELECTRON CONFIGURATIONS OF
THE FOLLOWING IONS
Na+1
S-2
Kr+2

50. WRITING ELECTRON CONFIGURATIONS FROM THE
INFORMATION ON THE PERIODIC TABLE OF
ELEMENTS.

51. WRITING ELECTRON CONFIGURATIONS FROM THE
INFORMATION ON THE PERIODIC TABLE OF
ELEMENTS.

52. WRITING ELECTRON CONFIGURATIONS FROM THE
INFORMATION ON THE PERIODIC TABLE OF
ELEMENTS.

53. EXCITED STATE ELECTRON CONFIGURATIONS
 A phosphorus atom has an electron configuration of 1s22s22p63s13p4 . Is the
atom in its ground state, or is it in an excited state?
P

54. EXCITED STATE ELECTRON CONFIGURATIONS
 Which electron configuration represents an atom in
the excited state?
1. 1s22s22p63s2
2. 1s22s22p63s1
3. 1s22s22p6
4. 1s22s22p53s2

55. VALENCE ELECTRON & LEWIS DOT DIAGRAMS
 Valence electrons are the electrons that fill the
outermost prinicpal energy level of an atom.
 Example;
 Mg 2-8-2 has 2 valence electrons.
2 2 6
 Ne 1s 2s 2p has 8 valence electrons.
 Valence electrons are largely responsible for an
elemement’s chemical and physical properties.

56. DO NOW: HOW MANY VALENCE ELECTRONS DO
EACH OF THE FOLLOWING ATOMS HAVE?
Na Al Cl
Na+1 Al+3 Cl-1
Na Si S-2
Mg+2 Si-4 Ar

57. DO NOW:
 What is the most common isotope of the element
Bromine?
 How many protons, neutrons and electrons does it
have?
 How many valence electrons does it have?
 What are the most common ions for the element
bromine? (Hint: There are three)
 How many protons, neutrons and electrons does each
ion have?
 How many valence electrons does each ion have?

58. VALENCE ELECTRON & LEWIS DOT DIAGRAMS
 The term Kernel refers to all of the non-valence electrons as
well as the nucleus (p+ & n0) of the atom. The Kernel is
represented by the element’s symbol.
 Valence electrons are represented by dots.
Na Ne N
2-8-1 2-8-8 2-5

59. MORE ON ELECTRONS
 Similarly, shell electrons may be represented by arrows
pointing in opposite directions (up & down) occupying their
perspective orbitals.
1. Write the orbital notation for the outermost principal level for the
following elements.
Na
2-8-1
P
2-8-5
Cl
2-8-7

60. MORE ON ELECTRONS
2. Which is the correct orbital notation of a lithium atom
in its ground state

61. MORE ON ELECTRONS
3. Which orbital notation correctly represents a noble gas
in the ground state?

62. TABLE S
 Ionization Energy: the amount of energy needed to
remove the most loosely held electron from the
valence shell of an atom in the ground state.
 Low Ionization energy = EASY to remove e-’s (Fr)
 High Ionization Energy = DIFFICULT to remove e-’s (F)
 The lowest Ionization Energies are found in the lower left
corner of the periodic table (metals)

63. TABLE S
 Electronegativity: The affinity (attractiveness or pull
for) of electrons by an atom.
 High Electronegativity = Atoms most likely to gain e- (F)
 Low Electronegativity = atoms most likely to lose e- (Fr)
 The highest electronegativity (an arbitrary value of 4.0) can
be found in the upper right corner of the periodic table (non-
metals)
 Does not include the Noble Gases (Group 18)

64. ELECTRONEGATIVITY
 As atom gain or lose e- they become Ions.
0 -1
F F
0 +1
Fr Fr

65. ELECTRONEGATIVITY
 As atom gain or lose e- they become Ions.
High Ionization Energy (-) gains
0 High Electronegativity -1
F F one
electron
2-7 2-8
Low Ionization Energy (+) loses
0 +1
Fr Low Electronegativity
Fr one
electron
2-8-18-32-18-8-1 2-8-18-32-18-8

66. ELECTRONEGATIVITY
 As atom gain or lose e- they become Ions.
0 -1
F F
2-7 2-8
1s22s22p5 1s22s22p6
0
Fr
2-8-18-32-18-8-1
1s22s22p63s23p63d104s24p64d104f145s25p65d105f146s26p67s1
+1
Fr
2-8-18-32-18-8
1s22s22p63s23p63d104s24p64d104f145s25p65d105f146s26p67s1

67. ELECTRONEGATIVITY
 The Noble Gases are not assigned electronegativity values. This is due to
the fact that they tend not to gain or lose valence electrons because they
have complete octets (complete outer principal energy levels).
Helium
Neon
Argon
Krypton
Xenon
Radon

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