1. Unit 01 (Chp 6,7): Atoms and Periodic Properties
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Unit 01 (Chp 6,7): Atoms and Periodic Properties
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.

2. Development of the Atomic Model
empty space
Indivisible
Identical
React in fixed ratios
+ stuff
– electrons +
nucleus

3. Development of the Atomic Model
Rutherford’s atomic model didn’t explain
properties of matter (color, reactivity, …) Li Na Cu

4. Atomic Emission Spectra
white light  continuous spectrum
prism
Atomic Emission Spectra
elements  discrete lines of E & f
helium (He)
lamp
prism
(only specific colors of energy & frequency)

5. Hydrogen Emission Spectrum
A mystery for Niels Bohr.

6. Bohr’s Shell Model + ∆E EXCITED state
(1913–Niels Bohr)
EXCITED state
e–’s emit (–) energy, move back to inner levels
(n=5 to n=2)
e–’s absorb (+) energy, move to outer levels
(n=2 to n=5) +
GROUNDstate 5 4 3 2 ∆E 2 2
Which transition shows a light wave of the greatest energy?
n=5 to n=2

7. Photon Energy as Light Waves
Distance between same point on adjacent waves is the _______________. (m)
Number of Waves passing a given point per unit time is the ______________. (Hz)(s–1)
wavelength ()
frequency ()
lambda nu
 and  are
_________ proportional
inversely

8. All light waves move at the same speed, so which color has more energy?

9. Electromagnetic (EM) SpectrumY G B I V
Low Energy
High Energy
Electromagnetic (EM) Spectrum
Low Frequency
High Frequency
(higher E) (higher ) (shorter )
All EM waves travel the same speed:
the speed of light (c),  108 m/s.
c = 

10. Photon (Light) Calculations
Given wavelength () of light, one can
calculate the energy (E) of 1 photon of that light:
Speed of light:
Plank’s constant:
2.998  108 m/s
(constants)
6.626  10–34 J•s
(given on Exam)
c = 
E = h
 ,  (inverse) E ,  (direct)
Avogadro’s number:
6.022  1023 particles/mole
 ↔ E
 ↔ 
HW p. 253
#14,25ab,26,34

11. Quantum Mechanical Model
(1926–Schrodinger, Plank, de Broglie, etc. )
Heisenberg Uncertainty Principle:
The more precisely a particle’s motion is known,…
…the less precisely its
position is known.
(wave)
(E , l , n)
(particle)
(probable locations)
Schrödinger Wave Equation:
s , p , d , f
3-D regions of probability (ORBITALS) in sublevels
of each fixed energy level which better explains reactivity.

12. Quantum Mechanical Model
Electrons as Waves (instead of particles)
electrons occupy only specific levels (shells)
of “quantized” energy
(& wavelength & frequency)
nucleus
“quantized” into specific multiples of
wavelengths,
but none in between.

13. Development of Atomic Models
1803 Dalton
Atomic Theory
1904 Thomson
Plum Pudding + –
1911 Rutherford
Nuclear Model – – – + – –
1913 Bohr
Shell Model
These illustrations show how the atomic model has changed as scientists learned more about the atom’s structure.
1926 Quantum Mechanical Model

14. Where are the electrons really?
(Shell) principle energy level (n) (1,2,3,4 …)
(Sub-shell) shape
(Orbital) 3-D arranged
(Electron) spin up/down
(not rings)
s (1)
p (3)
d (5)
f (7) x y z
HW p. 255
#57ac, 60

15. Electron Configuration (arrangement)
Orbital Notation +8
# of e–’s in each sublevel
Oxygen (O)
1s2 2s2 2p4
1s2 2s2 2p4
1s2 2s2 2p4
1s2 2s2 2p4
energy level (shell, n)
sublevelshape (s,p,d,f)

16. Electron Configuration (arrangement)+8
Oxygen (O)
1s2 2s2 2p4
1s2 2s2 2p4 6
E-Config?
Element?
How many valence e–’s? Na
____________
1s2 2s2 2p6 3s1
(outer level) __ Al
1s2 2s2 2p6 3s2 3p1 Cl __
[Ne] 3s2 3p5
(noble gas core configuration)

17. d orbital e–’s are core e–’s …NOT valence e–’s
2p6
2p3
3s2
3p6
4s2
3d10
4p2
Hund:
1 e– in equal orbitals before pairing
()
(3d fills after 4s) ?
Pauli Exclusion:
-Where should we start placing electrons first? -Opposite spin alleviates repulsion. -All single before double.
no e–’s same props
(opp. spin) (↑↓)
nucleus
Aufbau:
fill lowest energy
orbitals first. +

18. Electron Configuration of Ions
Ion E-Con (i) F– (ii) Ca2+ (iii) S2– (iv) Na+ (v) Al3+
1s2 2s2 2p6 [Ne]
1s2 2s2 2p6 3s2 3p6 [Ar]
1s2 2s2 2p6 3s2 3p6 [Ar]
1s2 2s2 2p6 [Ne]
Which ions are isoelectronic?
F– , Na+ , Al3+
Ca2+ , S2–
List 3 species isoelectronic with Ca2+ & S2–.
P3– , Cl– , Ar, K+ , Sc3+ , Ti4+, V5+, Cr6+, Mn7+

19. Other Aspects of Electron Configs
Paramagnetic:
species attracted by a magnet
(caused by unpaired electrons).
Fe: [Ar] ↑↓ ↑↓ ↑ ↑ ↑ ↑
4s d
Diamagnetic:
species repelled by magnets
(caused by all paired electrons)
Zn: [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
(“di-” is 2)

20. Other Aspects of Electron Configs
d block metals lose their outer s electrons before any core d electrons to form ions.
Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Fe2+ 1s2 2s2 2p6 3s2 3p6 3d6
HW p.255 #74
Fe3+ 1s2 2s2 2p6 3s2 3p6 3d5
video
d block (trans. metals) have colored ions b/c light excites e– transitions in d orbitals
(Video – V5+  V4+  V3+  V2+ with different colors) (4 min) (1:50 - 4:36)

21. SPECTROSCOPIC TECHNIQUE
Spectroscopy
SPECTROSCOPIC TECHNIQUE
EM REGION
APPLICATION
Molecular Structure by molecular Rotation
Microwave
Microwave
Types of bonds by bond Vibration IR
Infrared
Vis/UV Atomic Emission Spectra (lines of frequencies/colors)
Visible & Ultraviolet
Transition of e–’s between energy levels
Ionization of e–’s shows e– configuration
PES (Photoelectron Spectroscopy)
X-ray
WATCH this 6 min Video Explanation of PES at HOME.

22. Photoelectron Spectroscopy (PES)
Which peak is H and which is He?
higher peak = more e–’s
1s2
Relative # of e–’s He
1s1 H
Binding Energy
...or Ionization Energy
(required to remove e–’s)
(MJ/mol)
 further left = more energy required
(stronger attraction
due to more protons)

23. Photoelectron Spectroscopy (PES)
Which peak is H and which is He?
2p6
higher peak = more e–’s ? Ne
1s2
Identify the element
& e-config
Relative # of e–’s He
1s1
1s2
2s2 H
Binding Energy
...or Ionization Energy
(required to remove e–’s)
(MJ/mol)
 further left = more energy required
(stronger attraction
due to more protons)

24. PES (A) PES (B) 3d10 2p6 3p6 1s2 2s2 3s2 4s2 4p2 Ge n = 1 n = 2 n = 3
Identify element (A)
3d10
2p6
3p6
1s2
2s2
3s2
4s2
4p2 Ge
n = 1
n = 2
n = 3
n = 4
PES (B)
Identify element (B)
2p6
3p6
4s1
1s2
2s2
3s2 ? K

25. Write the complete electron configuration of element (X), and identify the element.
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1 Ga
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1
WS 3a
PES (X)
3d10
2p6
3p6
4s2
1s2
2s2
4p1
3s2

26. Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Unit 1 (Chp 7): Periodic Properties …or… Periodicity of Trends in Atomic Properties
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.

27. Periodic Trends We will explain observed trends in Zeff & shielding
Atomic (and Ionic) Radius
Ionization energy
Electronegativity
size
lose e–
attract e–
Zeff & shielding
(explains ALL periodic trends and properties)

28. Zeff & Shielding attraction shielding Zeff Na atom
effective nuclear charge, (Zeff):
Zeff = Z − S
Z = nuclear charge (+proton’s)
S = shielding
(core e–’s)
attraction
shielding
Zeff
shielding, (S):
inner core e–’s shield valence e–’s from nuclear attraction.
Z = +11
+11
Zeff = +1
Na atom

29. decreases across a period
Atomic Radius
att.
=shield
Zeff
decreases across a period
-due to increasing Zeff
(more protons)
-due to
increasing shielding
(more energy levels)
increases down a group
att.
shield
=Zeff

30. Ionic Radius Na+ Cations are smaller than neutral atoms. e– e–
outermost electron(s) are removed and loses a shell
core shell closer to nucleus
inner e–’s shielded (Zeff)

31. Ca2+ < K+ < Ar < Cl– < S2–
Ionic Radius e– e– e– HW
p. 292
#13,28
Anions are larger than their parent atoms.
electrons are added and repulsions are increased
(=Zeff & =shielding)
Arrange the following species by
increasing size: Ar, K+, Ca2+, S2–, Cl–
Ca2+ < K+ < Ar < Cl– < S2–

32. Ionization Energy (IE)
energy required to remove an electron
more energy to remove next electron
IE1 < IE2 < IE3, …
look for a huge jump in IE
once all valence e–’s are removed, the next e– is on an inner level with attraction (shielding & Zeff).
huge jump in IE4 b/c 4th e– on inner level
(must have 3 valence e–’s)

33. increases across a period
Trends in First IE
att.
=shield
Zeff
increases across a period
-due to increasing Zeff
(more protons)
-due to
increasing shielding
(more energy levels)
decreases down a group
att.
shield
=Zeff

34. 5B & 8O exceptions to trend.
Does this graph support your understanding of IE1 and the Periodic Table?
Figure: 07-10
Title:
First ionization energy versus atomic number.
Caption:
The red dots mark the beginning of a period (alkali metals), the blue dots mark the end of a period (noble gases), and the black dots indicate other representative elements. Green dots are used for the transition metals.
5B & 8O exceptions to trend.
Why?

35. Exceptions to 1st IE Trend
1st IE tends to…
increase across period
(Zeff , =shielding) ↑↓ 2s ↑ 2p ↑↓ 2s B Be
1st IE of B < Be b/c…
The e– in 2p orbital of B is higher energy than the e– in 2s orbital of Be ; less energy needed to remove 1st e– in B.

36. Exceptions to 1st IE Trend
1st IE tends to…
increase across period
(Zeff , =shielding) ↑↓ 2s ↑ ↑ 2p ↑ ↑↓ 2s ↑ ↑ 2p ↑ O N
1st IE of O < N b/c…
The paired e– in 2p orbital of O experiences e–---e– repulsion requiring less energy to remove 1st e– in O.
HW p. 293 #38,46

37. Trends in Electronegativity (EN)
-ability of an atom to attract electrons when bonded (sharing e–’s) with another atom.
att.
=shield
Zeff
increases across a period
-due to increasing Zeff
(more protons)
-due to
increasing shielding
decreases down a group
att.
shield
=Zeff

38. Periodic Table Elements arranged by… atomic #
How are the elements arranged? Used to be by at. mass and repeated properties were seen, now better by at. #.

39. Periodic Table
Metals on the left
(80% of all elements)

40. Periodic Table
Nonmetals on the right
(except H)

41. Metalloids border the stair-step
Periodic Table
Metalloids border the stair-step
(Al is metal)
Why called metalloids, what are their props (like both, discuss Si as semiconductor)

42. Periodic Table Rows on the periodic chart are called _______. periods
Columns are _______.
Elements in the same ______ have similar _________________.
periods
Why do they have similar props? (same val #) (electrons are the deciding factor of reactivity)
(shell, n)
(energy level)
groups
(# val. e–’s)
group
chemical properties

43. Group Names (1, 2, 17, 18) 1 2 17 18

44. Group 1: Alkali Metals lowest IE’s (lose e–’s easily) Zeff
more reactive down a group b/c…
shielding causes
att. & IE,
easier to lose e–
soft, metallic solids
Why low IE’s? (video clip -
video clip

45. Group 2: Alkaline Earth Metals
low IE’s, but not as low as alkali metals.
less reactive than alkali metals (Zeff , att. & IE),
but more reactive down the group.
(shielding causes att. & IE)

46. Group 17: Halogens high IE’s (don’t lose e–’s easily) (Zeff , att.)
large EN (attract e–) (Zeff , att.)
more reactive at top of a group b/c…
shielding causes att. & EN,
easier for nonmetals to attract e–

47. Group 18: Noble Gases UNREACTIVE (mostly) b/c… HUGE IE’s b/c……
Monatomic gases
Zeff , att. (no lose e–), and filled valence shell (no gain e–)

48. Metals vs. Nonmetals Si Table 7.3 p. 277 (in book) Metalloids:
characteristics of metals & nonmetals.
Silicon is shiny, but brittle and is a semi-conductor. Si

49. Periodicity: –repeating pattern of properties nonreactive nonmetals
soft highly reactive metals
highly reactive nonmetals
harder less reactive metals
Reactivities are based on valence electrons!

50. Periodic Trends (Summary)
Electronegativity
Can you explain all of this in terms of p’s and e’s?
Zeff & shielding
Electronegativity
WS Periodicity
WS 7a
Atomic radius