Chemistry 2: Periodic Table
Quantum Mechanical Model (6.5 to 6.6)
1. electron’s exact position or velocity is unknowable
a. Heisenberg uncertainty principle
b. orbital = 90 % probable location of an electron
2. each electron has four quantum numbers: n, l, ml, ms
a. principal energy levels (n)—defines orbital radius
b. sublevels (l)—defines orbital shape
1. l = 0, 1, 2, •••, (n - 1)
2. 0 = s, 1 = p, 2 = d, 3 = f
c. orbital (ml)—defines spatial orientation
1. ml = – l, ••• -1, 0, +1, •••, + l
2. number of orbitals: s (1), p (3), d (5), f (7)
d. spin (ms)—defines magnetic field +½ (), -½ ()
Pauli exclusion principle—no two electrons can
have the same spin in the same orbital
3. relationship among values of n, l, ml through n = 4
n Possible l Sublevel Designation
Possible ml
1s
1
0
0
2s
0
0
2
2p
1
-1, 0, 1
3s
0
0
3p
3
1
-1, 0, 1
3d
2
-2, -1, 0, 1, 2
4s
0
0
4p
1
-1, 0, 1
4
4d
2
-2, -1, 0, 1, 2
4f
3
-3, -2, -1, 0, 1, 2, 3
B. Electron Arrangements in Atoms and Ions (6.7 to 6.9)
1. electrons fill from low to high energy (same for all atoms)
a. n: 1 < 2 < 3 < 4 < 5 < 6 < 7
b. l: s < p < next energy level < d < f
c. (n, ml) equal energy = degenerate
2. outer (valence) electrons—interact with other atoms
a. highest occupied principle energy level
b. s and p sublevels only (maximum 8 electrons)
5s
4p
3d
4s
3p
Energy 
A.
3s
2p
2s
1s
3.
4.
1
2
3
4
5
6
7
inner (core) electrons + nucleus = core charge (+1 to +8)
organization of the periodic table
a. row (period)—same valence energy level (n)
b. column (group)—same # of valence electrons
1. main groups (1: alkali metals, 2: alkaline
earth metals, 17: halogens, 18: noble gases)
2. similar chemical properties
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1s
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
lanthanide
actinide
Name __________________________
5.
types of diagrams
a. electron configuration
1. n (#), l (letter), # of electrons (superscript)
Al: 1s22s22p63s23p1
2. abbreviated: replace inner (core) electrons with
noble gas symbol—Al: [Ne]3s23p1
b. orbital diagrams
1. electrons (arrows) fill specific orbital
2. Hund's rule: maximum # of electrons with
the same electron spin (maximum number of
half-filled degenerate orbitals)
Electron
Orbital Diagram
Element # eConfiguration
1s
2s
2p
3s
Li
3 




 1s2 2s1
Be
4 




 1s2 2s2
B
5 

 

 1s2 2s2 2p1
C
6 

  
 1s2 2s2 2p2
N
7 

  
 1s2 2s2 2p3
O
8 

  
 1s2 2s2 2p4
F
9 

  
 1s2 2s2 2p5
Ne
10 

  
 1s2 2s2 2p6
Na
11 

  
 1s2 2s2 2p6 3s1
c.
n
1
2
3
4
quantum numbers (arranged by periodic table
position)
l=0
ms (½, -½)
(s)
½ -½
½ -½
l = 2 (n = row # – 1)
(d)
½ -½
½ -½ ½ ½ ½ ½ ½ -½ -½ -½ -½ -½
0 0 -2 -1 0 1 2 -2 -1 0 1 2
ml
6.
7.
½
½
½
-1
½
½
½
0
½
½
½
1
-½
-½
-½
-1
-½
-½
-½
0
-½
-½
-½
1
column 6 and 11: an s electron moves to the d sublevel
to maximize half and/or full orbitals (lower energy state)
4s
3d
[Ar]
24Cr







[Ar]
29Cu


    
electron arrangements in monatomic ions
a. ions with noble gas structure
1. elements that are ± 3 of the noble gas lose or
gain electrons to reach noble gas electron
configuration
3210
1+
2+
3+
N3O2FNe
Na+ Mg2+ Al3+
2. ions with the same # of e: isoelectronic
b. transition metal ions
1. transition metal lose s electrons first
2. may lose d electrons if it eliminates sublevel
or reduces doubling up
4s
3d
+
[Ar] 
29Cu

    
[Ar] 






magnetic properties
a. element is magnetic (paramagnetic) if it contains
unpaired electrons, whose magnetic fields are
reinforcing  respond to external magnetic field
4s
3d
[Ar]  
26Fe
 



b. elements with all paired electrons (diamagnetic)
are unaffected by magnetic fields
(columns 2, 12, 18)
26Fe
8.
l=1
(p)
3+
C.
Periodic Properties—Main Groups (7.1 to 7.6)
1. effective nuclear charge (Zeff) or shielding effect
a. atom holds valence electrons because of
attraction between atom core and valence shell
b. atomic core = nucleus + core electrons
1. core charge = Zeff  (# p – # core e)
2. other valence electrons reduce Zeff because of
electron-electron repulsion
2. atomic size (radius)
a. minimum distance between two gas atoms or
distance between nuclei in diatomic molecule
b. group: increase as energy level increases
c. period: decrease as Zeff increases from +1 to + 8
d. transition metals
1. group: increase with energy level
2. period: no change (Zeff  constant)
3. ionic size (radius) compared to parent atom
a. smaller cations (lose energy level)
b. larger anions (more electron-electron repulsion)
c. isoelectronic series (± 3 from noble gas) largest
(fewest protons) to smallest (most protons)
N3- > O2- > F- > Ne > Na+ > Mg2+ > Al3+
4. ionization energy
a. E to remove electron from a gaseous atom
1. X(g)  X+(g) + 1e2. all +E (greater value = harder to ionize)
b. inversely proportional to atomic radius (E  1/r)
1. weaker hold on distant electron  less
energy to remove electron
2. anomalies: ionized electron comes from a
relatively less negative energy level  less
energy to reach zero energy (ionization)
a. 13: ionized electron comes from a higher
energy sublevel p vs. s
b. 16: ionized electron comes from a full
orbital (higher energy than ½-filled)
c. successive ionization energies
Successive Ionization Energies in kJ/mol
Element
I1
I2
I3
I4
Na
496
4562***
(core electrons)
Mg
738
1451*
7733***
Al
578
1817**
2745*
11,577***
1. *small increases within a sublevel
2. **greater increase between sublevels
3. ***greatest increase between energy levels
5. electron affinity
a. E to add electron to gaseous atom
1. X(g) + 1 e-  X-(g)
2. E = EX X- + Eea. EX  X- > 0 (added electron increases
atom's overall energy level)
b. Ee- < 0 (added electron's energy
decreases from zero)
c. –E when added electron enters
relatively low-energy orbital (stable anion)
d. +E when added electron enters relatively
high-energy orbital (unstable anion)
b. electron affinity becomes more negative from left to
right because receiving orbital energy decreases
c. anomalies: receiving orbital energy is relatively high

2: p orbital energy > s orbital energy

15: full orbital energy > ½-filled orbital energy
(electron-electron repulsion raises level)

18: next higher energy level > valence level
d. group: little change for next energy level because
greater orbital volume reduces electron-electron
repulsion, which counterbalances reduced core
attraction
6.
metals, nonmetals and metalloids
a. metals—left side of stair step
1. shiny, conduct heat and electricity, malleable
and ductile, mostly solids (except Hg)
2. form ionic compounds with nonmetals
3. small positive ionization energy
4. positive or small negative electron affinity
5. lose electrons during reactions
(alkali metals are most reactive)
b. nonmetals—right side of stair step and H
1. opposite properties to metals
2. form molecules in addition to ionic compounds
3. large positive ionization energy
4. large negative electron affinity except
columns 15 and 18
5. gain or share electrons during reactions except
noble gases (halogens are most reactive)
c. metalloids—touch stair step except Al
1. intermediate properties depending on physical
and chemical conditions
2. Po and At classification unresolved
Experiments
1.
Halogen Lab (Wear Goggles)—Mix halogen molecules
with halide ions to determine if an electron is transferred
and in doing so, rank Cl, Br and I from most negative
electron affinity to least negative.
Add 20 drops of bromine water to three small test tubes.
Add 20 drops of hexanes (HEX) to each test tube. Stopper
the test tubes and shake until the bromine color is mostly in
the HEX layer. (AVOID BREATHING OR TOUCHING
HALOGENS). Record the color of the top (HEX) layer. Add
20 drops of 0.1 M NaCl to the first tube, 20 drops of 0.1 M
NaBr to the second tube, and 20 drops of 0.1 M NaI to the
third. Stopper and shake each tube. Record the color of the
HEX layer. Repeat with chlorine water and iodine water.
a. How do you know if the halogen molecule took an
electron from the halide ion?
b.
Record the color of the hex layer after adding the
halogen, and after adding the halide ions. Highlight the
box where a reaction occurred.
HEX layer Color
Halogen
HEX
water
Color
With Br
With ClWith I-
2.
Fill in the quantum numbers for the first three principle
energy levels, and then answer the questions below.
n
1
2
3
l
ml
ms
a.
b.
3.
4.
How many electrons have the quantum numbers?
2,1, __, __
3, __, __, +½
B. Electron Arrangements in Atoms and Ions
Write the order that electrons fill sublevels from 1s to 7p.
Complete the chart for each element.
Period Group
Symbol
Group Name
#
#
5
Halogen
4
Cl2
I2
Highlight which reaction (1 or 2) occurred for each pair.
Pair
Reaction 1
Reaction 2
Br2 & I- Br2 + 2 I-  2 Br- + I2
I2 + 2 Br-  2 I- + Br2
5.
c.
6.
Cl2 & Br- Cl2 + 2 Br-  2 Cl- + Br2 Br2 + 2 Cl-  2 Br- + Cl2
Cl2 & ICl2 + 2 I-  2 Cl- + I2
I2 + 2 Cl-  2 I- + Cl2
d. The halogen with the more negative electron affinity will
take an electron from the halogen with the less negative
electron affinity. Rank the halides from most negative
electron affinity to least negative electron affinity.
1
What is the correlation between electron affinity and
atomic radius? Give a reason for this correlation.
3
noble gas
In the electron configuration (1s2) what does each indicate?
2
s
1
Write the electron configuration for
magnesium
silicon
Iron
7.
What does the symbol [Ar] represent in [Ar]4s23d8?
8.
Write the two ways to show the abbreviated configuration
for iodine.
emphasize filling order
emphasize valence shell
How do the results compare to the electron affinity data?
9.
Cl
Br
I
-349 kJ/mol
-325 kJ/mol
-295 kJ/mol
f.
Metal or
Nonmetal
Be
Br2
e.
Give the four quantum numbers for the two electrons
in the 3d sublevel where ml = 0.
Write the orbital diagram for the following.
1s 2s
2p
3s
3p
4s
S           
Co











3d








10. What is the abbreviated electron configuration for each?
Cr
Cu
Mo
Ag
11. Write the quantum numbers for the boxed electrons.
1s 2s
2p
3s
3p
4s
3d
              
Practice Problems
1.
A. Quantum Mechanical Model
Complete the number of electrons in each sublevel and
total for the following energy levels.
Maximum in sublevels, 2(2l +1)
Level
Total
n
2n2
s (l = 0) p (l = 1) d (l = 2) f (l = 3)
1
2
3
4
12. Write the set of quantum numbers for the electron position
(a-d) on the periodic table below.
1
2
a
3
4
b
d
5
c
a
b
c
d
C. Periodic Properties—Main Groups
13. List the ions with Ar noble gas structure.
cations
anions
b. Ionization Energy (kJ/mol)
Group #
1
2
13
14
15
16
17
18
Period 2 520 899 801 1086 1402 1314 1681 2081
Period 3 496 738 578 786 1012 1000 1251 1521
14. Write the abbreviated electron configurations for each ion.
Sc3+
Ag+
Ru3+
Zn2+
2100
15. Label the following as paramagnetic or diamagnetic.
Cu
Mg
Cu+
1700
16. Use the relative sizes of atoms and their common ion for
columns 1, 2, 3, 16 and 17 to answer the questions.
1300
1900
1500
1100
900
700
500
c. Electron Affinity (kJ/mol)
Group #
1
2
13
14
Period 2 -60
>0
-27 -122
Period 3 -53
>0
-43 -134
15
>0
-72
16
-141
-200
17
-328
-349
18
>0
>0
>0
0
-100
-200
a.
Indicate whether the trend is increase or decrease for.
Relative Size
Increase Decrease
Atoms from top to bottom
Cations from top to bottom
Anions from top to bottom
Atoms from left to right
Cations from left to right
Anions from left to right
Isoelectric ions from low Z to high Z
Size of cation compared to atom
Size of anion compared to atom
b. Explain why cations are smaller than their atoms.
c.
17. Which period 3 element has the following successive
ionization energies?
First
Second
Third
Fourth
Fifth
786
1577
3232
4356
16,091
16
0.73
1.02
17
0.71
0.99
19. Use the graphed data above to complete the table.
Atomic Radius
Period Pattern left to right
Anomaly group numbers
Ionization Energy
Period Pattern left to right
Anomaly group numbers
Electron Affinity
Period Pattern left to right
Anomaly group numbers
20. Explain the following observations.
a. Atomic radii of Li = 1.34 Å and Na = 1.54 Å.
Explain why anions are larger than their atom.
18. Graph the data on the grids below.
a. Atomic Radius (x 10-10 m)
Group #
1
2
13
14
15
Period 2 1.34 0.90 0.82 0.77 0.75
Period 3 1.54 1.30 1.18 1.11 1.06
-300
18
0.69
0.97
b.
Atomic radii of Al = 1.18 Å and Si = 1.11 Å.
c.
First ionization energies for B = 801 kJ/mol and Al =
578 kJ/mol.
d.
First ionization energies for Si = 786 kJ/mol and P =
1012 kJ/mol.
1.5
1.3
1.1
0.9
0.7
e.
f.
g.
h.
i.
First ionization energies for Mg = 738 kJ/mol and Al =
578 kJ/mol.
First ionization energies for P = 1012 kJ/mol and S =
1000 kJ/mol.
For Na, the first ionization energy I1 = 495 kJ/mol and
second ionization energy I2 = 4562 kJ/mol.
The gap between the first and second ionization
energies is greater for Al (I1 = 578 kJ/mol and I2 = 1817
kJ/mol) than Si (I1 = 786 kJ/mol and I2 = 1577 kJ/mol).
The common ion for magnesium is Mg2+, where I1 =
738 kJ/mol, I2 = 1451 kJ/mol and I3 = 7733 kJ/mol.
j.
The electron affinities for Mg > 0 and Na = -53 kJ/mol.
k.
The electron affinities for Si = -134 kJ/mol and
P = -72 kJ/mol.
2.
States that an orbital can hold no more than two electrons.
3.
Predicts that it is impossible to determine simultaneously
the exact position and the exact velocity of an electron.
4.
Which set of quantum numbers describes the highest energy
valence electron in a ground-state gallium atom (Z = 31)?
(A) 4,0,0,½ (B) 4,0,1,½ (C) 4,1,1,½ (D) 4,1,2,½
5.
Which has the outer electronic configuration, s2p3?
(A) Si
(B) Cl
(C) Se
(D) As
Questions 6-9 refer to atoms with the atomic orbitals shown.
(A) 1s 2s
(B) [He] 2s2p 
(C) [He] 2s2p (D) [Ar] 4s3d
6. Represents an atom that is chemically unreactive.
7.
Represents an atom in an excited state.
8.
Represents an atom that has four valence electrons.
9.
Represents an atom of a transition metal.
Questions 10-12
(A) 1s2 2s22p5 3s23p5
(B) 1s2 2s22p6 3s23p6
(C) 1s2 2s22p62d10 3s23p6
(D) 1s2 2s22p6 3s23p63d3 4s2
10. An impossible electronic configuration
11. The ground-state configuration for a halogen anion
l.
The electron affinities for S = -200 kJ/mol and
Cl = -349 kJ/mol.
m. The electron affinities for Cl = -349 kJ/mol and Ar > 0.
Practice Multiple Choice
Briefly explain why the answer is correct in the space provided.
Questions 1-3
(A) Heisenberg uncertainty principle
(B) Pauli exclusion principle
(C) Hund's rule
(D) Shielding effect
1. Can be used to predict that a gaseous carbon atom in its
ground state is paramagnetic.
12. The ground-state configuration for an alkaline earth cation
13. Which represents the ground state for the Mn3+ ion?
(A) 1s2 2s22p6 3s23p63d4
(B) 1s2 2s22p6 3s23p63d5 4s2
(C) 1s2 2s22p6 3s23p63d2 4s2
(D) 1s2 2s22p6 3s23p63d8 4s2
14. In which group are the three species isoelectronic?
(A) S2-, K+, Ca2+
(B) Sc, Ti, V2+
(C) O2-, S2-, CI(D) Mg2+, Ca2+, Sr2+
15. The ionization energies for element X are listed in the table
below. On the basis of the data, element X is most likely
Ionization Energies for element X (kJ mol-1)
First
Second
Third
Fourth
Fifth
580
1,815
2,740
11,600
14,800
(A) Na
(B) Mg
(C) Al
(D) Si
16. In the periodic table, as the atomic number increases from
11 to 17, what happens to the atomic radius?
(A) It remains constant. (B) It increases only.
(C) It decreases only.
(D) It increases, then decreases.
b.
c. Complete the following.
Electron configuration of element 3
Common ion charge of element 2
17. Which elements have most nearly the same atomic radius?
(A) Be, B, C, N
(B) Ne, Ar, Kr, Xe
(C) Mg, Ca, Sr, Ba
(D) Cr, Mn, Fe, Co
Questions 18-19 Use the following options.
(A) O
(B) Rb
(C) N
(D) Mg
18. What element has the most negative electron affinity?
Identify element 3. Explain.
Chemical symbol of element 2
Element with the smallest atomic radius
Chemical symbol for element 4
3. Use the principles of atomic structure to explain each.
a. The atomic radius of Li is larger than that of Be.
19. Which of the elements above has the smallest ionic radius
for its most commonly found ion?
b.
The electron affinity for K is less than Ca.
c.
The first ionization energy of Se is less than As.
20. Which gaseous atoms (Ca, V, Co, Zn, As) are paramagnetic?
(A) Ca and As only
(B) Zn and As only
(C) Ca, V, and Co only
(D) V, Co, and As only
21. Which property generally decreases across the periodic
table from sodium to chlorine?
(A) 1st ionization energy (B) Atomic mass
(C) Ionic radius
(D) Atomic radius
4.
Questions 22-23 Consider a ground state atom of each element.
(A) S
(B) Ca
(C) Ga
(D) Sb
22. The atom that contains exactly two unpaired electrons
Consider the element strontium (Sr). Justify each answer.
a. What is the outer electron configuration of Sr?
b.
How does the atomic radius of Sr compare to Rb?
c.
How does the atomic radius of Sr compare to Ca?
d.
Compare Sr to Ca and Rb in first ionization energy.
e.
How does the Sr2+ ion compare in size to the Sr atom?
f.
How does the Sr2+ ion compare in size to the Br- ion?
g.
As successive electrons are removed from the Sr
atom, where does the largest jump in ionization
energy occur?
h.
Is strontium diamagnetic or paramagnetic?
i.
Highlight the correct option when completing the
following sentence: compared to Br, Sr is shinier/duller
and a better/worst conductor.
23. The atom that contains only one electron in the highest
occupied energy sublevel
Practice Free Response
1.
For the following elements:
a. Write the abbreviated electron configuration
b. Write the abbreviated orbital diagram
c. Circle the electron in (b) that has the quantum numbers.
d. Write the abbreviated electron configuration for the ion
Mg
Cu
P
a
b
c
d
2.
(3, 0, 0, -½)
Mg2+
(3, 2, -2, +½)
Cu+
(3, 1, 0, ½)
P3-
The table shows the first three ionization energies in
kJ/mol for third period elements, which are numbered
randomly. Use the information to answer the questions.
Element
First
Second
Third
1
1,251
2,300
3,820
2
496
4,562
6,910
3
738
1,451
7,733
4
578
1817
2745
a. Which element is most metallic in character? Explain.