Periodic Properties
Chapter 7
Periodic Properties
Periodic Properties –depend on element’s position on table
Ex: Groups
H, Li, & Na all form similar oxides
(H2O, Li2O, Na2O)
Location gives you A LOT of information
Periodic Properties
2 Main Factors for Periodic Properties
1. Number of Levels – More levels, electrons
held less tightly
1. More shielding
2. More electron to electron repulsion (PUSH)
2. Number of Protons – More protons hold
electrons more tightly. Greater effective
nuclear charge (Zeff) (PULL)
Properties we will study: Periodic
1. Size of Atoms
2. Size of Ions
3. Ionization Energy
4. Multiple Ionization Energy
5. Electron Affinity
Properties
Size of Atoms
Atomic Radius
1. Measured in
picometers (1pm = 1 X 10-12 m) or
Angstroms (1 Å = 100 pm)
2. Average radius ~100 pm (1 Å)
Size of Atoms
3. Example: Bromine
1.14 Å X 100 pm = 114 pm
1Å
1.14 Å
Effective Nuclear Charge
Zeff = Z-S
Z = # protons
S = # core electron
What is the Zeff for chlorine (1s22s22p63s23p5)?
Zeff = Z-S
Zeff = 17- 10 = 7+
Size of Atoms
Down a group
– Atoms get larger (more levels)
– Shielding Effect – Core electrons shield the pull of the
nucleus (more electron repulsion) (PUSH wins)
H
1 level
Zeff = 1+
Li
2 levels Zeff = 1+
Na 3 levels Zeff = 1+
Size of Atoms
Size of Atoms
Across a period – atoms get smaller (same level),
greater Zeff (PULL wins)
Li
E config
Levels
Zeff
F
Si
Cl
Size of Atoms
Size of Ions
A. Positive Ions
1. Example:
Mg
E config
Levels
Zeff
Mg+
Size of Ions
Size of Ions
Positive ions always smaller
– Fewer electrons to control
– Less e- to e- repulsion
Mg
E config
Levels
Zeff
Mg+
Mg2+
Size of Ions
B. Negative Ions
1. Example:
O
E config
Levels
Zeff
O2-
O3-
Size of Ions
Size of Ions
Negative ions always larger
– More electrons to control
– More e- to e- repulsion
More levels
If same
Greater Zeff
(same levels, greater Zeff smaller)
If same
Ions
Positive = smaller(less electron repulsion)
Negative = Larger (more electron repulsion)
Size Review
Which is larger and why?
Li
or K
S
or S2+
Mg or S
O
or Te
Size Review
Kurveball
K
or K+
Ionization Energy
A.Ionization energy – The energy needed to remove
an electron from an isolated gaseous atom or ion
Na  Na+ + e-
Ionization Energy (kJ/mol)
2500
He
Ne
2000
Ionization Energy (kJ/mol)
Ar
1500
H
1000
500
Li
Na
K
0
0
5
10
15
Atomic Number
20
25
A low energy photon will excite an electron
BOHR MODEL
A high energy photon may ionize an atom (completely
remove the electron)
PHOTOELECTRIC EFFECT
Ionization Energy
B. Ionization energy is inversely proportional to
atomic radius
Examples:
Li (520 kJ/mol)
Na (496 kJ/mol)
F (1681)
Cl (1251)
Ionization Energy
Which has the higher Ionization Energy and why?
C or O
Na or Cl
C or Sn
Mg or Ra
Multiple Ionization Energy
Multiple Ionizations - Removing more than one
electron
1st
Mg  Mg+ + e738 kJ/mol
2nd
Mg+  Mg2+ + e1450 kJ/mol
3rd
Mg2+  Mg3+ + e- 7732 kJ/mol
There is a large jump once you reach Noble Gas
Configuration (Fewer levels, spike in Zeff)
Multiple Ionization Energy
Multiple Ionization Energy
1st
2nd
3rd
4th
Al  Al + + eAl +  Al 2+ + eAl 2+  Al 3+ + eAl3+  Al4+ + e-
577 kJ/mol
1816 kJ/mol
2744 kJ/mol
11580 kJ/mol
Multiple Ionization Energy
Examples:
a. Where will the large jump in I.E. occur for:
Be B
P
b. Element X has a large jump between its 4th
and 5th I.E. To what group does it belong?
Electron Affinity
• Energy change that occurs when an electron is
added to an isolated gaseous atom or ion
• Ease with which an atom gains an electron
• Cl(g) + e-  Cl-(g)
DE = -349 kJ/mol
• Positive for noble gases (don’t want electrons)
Properties of Metals
•
•
•
•
Malleable and ductile
Good conductors
Large radius/Low ionization energy
Form positive ions (+2 and +3 for transition
metals)
Metal Oxides
• Most metals oxidize easily in the atmosphere
2Ni(s) + O2(g)  2NiO(s)
• Metal oxides are bases
Metal Oxide + Water  Metal hydroxide (base)
Na2O(s) + H2O(l)  2NaOH(aq)
CaO(s) + H2O(l)  Ca(OH)2(aq)
Metal Oxide + Acid  Salt + Water
MgO(s) + HCl(aq)  MgCl2(aq) + H2O(l)
NiO(s) + H2SO4(aq)  NiSO4(aq) + H2O(l)
Al(s) + O2(g) 
Zn(s) + O2(g) 
CaO(s) + H2O(l) 
Li2O(s) + H2O(l) 
Al2O3(s) + HNO3(aq) 
CuO(s) + H2SO4(aq) 
Properties of Non-Metals
•
•
•
•
Dull
Poor conductors
Solids, liquids and gases
Non-metal oxides are acidic
Non-metal Reactions
Combination Reactions
2Al(s) + 3Br2(l)  2AlBr3(s)
Non-Metal Oxide + Water  Acid
CO2(g) + H2O(l)  H2CO3(aq)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
Non-Metal Oxide + Base  Salt + Water
CO2(g) + 2NaOH(aq)  Na2CO3(aq) + H2O(l)
SO3(s) + 2KOH(l)  K2SO4(aq) + H2O(l)
Predict the products in the following reactions:
Ba(s) + S8(g) 
Ga(s) + Se(g) 
SO2(s) + H2O(l) 
SeO2(s) + 2NaOH(l) 
SO3(s) + H2O(l) 
(acid rain reaction)
Group Names
Trends: Alkali Metals
• Soft, reactive (M  M+ + e-)
Reactions with water
Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)
Reactions with hydrogen (less electronegative than
hydrogen)
2Na(s) + H2(g)  2NaH(s) [H-, hydride ion]
Reactions of Na and K with oxygen (exceptions)
2Na(s) + O2(g)  Na2O2(s) [peroxide]
K(s) + O2(g)  KO2(s)
[superoxide]
Trends: Alkaline Earth Metals
• Become more reactive as you go down the table
(just like alkali metals)
• 99% of calcium in the body is in the skeleton
Trends: Transition and Post-Transition
Metals
• Often +2 and +3 ions
• Pt and Au are fairly unreactive
Trends: Hydrogen
• Non-metal
• Forms hydrides (H-) with alkali and alkaline earth
metals
2Na(s) + H2(g)  2NaH(s)
Ca(s) + H2(g)  CaH2(s)
Trends: Chalcogens
• Ozone = O3 (between 10-50 km in the
atmosphere, allotrope of oxygen)
• Hydrogen peroxide
2H2O2(aq)  2H2O(l) + O2(g)
-good oxidizing agent/disinfectant
-Can be decomposed by light
Trends: Halogens
•
•
•
•
“Salt formers”
Reactivity decreases down the table
Fluorine is most reactive
Chlorine as a bleaching agent
Cl2(g) + H2O(l)  HCl + HOCl (bleach)
NaOCl
(Clorox)
Ca(OCl)2 (swimming pool bleach)
Noble Gases
•Some do react
•XeF4 shown below
8. Tc
12. a) Effective nuclear charge is decreased by core electrons
b) 1s is less shielded
22. 1.33 Angstroms
26. a) Na<Ca<Ba
b) As<Sb<Sn
c) Be<Si<Al
28. a) Z = Constant, e- repulsion increases (I->I>I+)
b) Z = Constant, e- repulsion increases (Ca2+ > Mg2+>Be2+ )
c) Fe>Fe2+>Fe3+
30 Ca = Largest, Ca2+ middle, Mg2+ smallest
32. a) Cl-, Arb) Sc3+, Ar c) Fe2+, none
d) Zn2+, none e) Sn4+, none
34 a) K+ smaller b) Zeff = +7(Cl-) and +9(K+)
36 a) Se <Se2-<Te2- b) Co3+<Fe3+<Fe2+
c) Ti4+<Sc3+<Ca d) Be2+<Na+<Ne
38. LiCl (2.09 ionic, 2.05 covalent)
NaCl (2.83 ionic, 2.53 covalent)
KBr (3.34 ionic, 3.10 covalent)
RbI (3.72 ionic, 3.44 covalent)
40. Sn  Sn+ + eSn+  Sn2+ + eTi3+  Ti4+ + e42. a) Li smaller, higher IE
b) Sc higher b/c NGC
c) Li+ is higher b/c NGC
44. a) Inversely proportional b) same explanation
46.a) Ti b) Cu
c) Cl
d) Sb
48.a) [Ar]3d3
b) [He]2s22p6
c) [Ne]3s23p6
d) [Ar]3d9
e) [Xe]6s24f145d10 e) [Xe]4f145d10
74. 2Cs + 2H2O 2CsOH + H2
Sr + 2H2O  Sr(OH)2 + H2
2Na + O2  Na2O2
Ca + I2  CaI2
68.K2O + H2O  2KOH
P2O3 + 3H2O  2H3PO3
Cr2O3 + 6HCl  2CrCl3 + 3H2O
SeO2 + 2KOH  H2O + K2SeO3
74.2Cs + 2H2O 2CsOH + H2
Sr + 2H2O  Sr(OH)2 + H2
2Na + O2  Na2O2
Ca + I2  CaI2
82.Cl2 + H2O  HOCl + HCl Ba + H2  BaH2
2Li + S  Li2S
F2 + Mg  MgF2
a. Rank the following according to size (smallest to largest)
Be
F
FO2- O
Be2+
b. Why is Be3+ difficult to produce?
c. Which factor is most important in explaining the size of the O2ion compared to the O atom: Zeff or mutual electron repulsion?
d. Which factor is most important in explaining the size of the P
atom compared to the Cl atom: Zeff or mutual electron repulsion?
a. Rank the following according to size (smallest to largest)
S2K+
Ca2+ Clb. Rank the following in order of increasing first ionization energy:
Ne, Na, P, Ar, K
c. Write the electron configuration for:
Co3+
S2Fe2+
d. Would you expect an acidic, basic or neutral pH if Li2O is placed
in water? CO2?
SO2(s) + 2NaOH(l) 
Cu(s) + O2(g) 
Ba(s) + H2(g) 
SeO3(s) + H2O(l) 
Al(s) + F2(g) 
Ca(s) + O2(g) 
CO2(s) + H2O(l) 
CaO(s) + H2O(l) 
K2O(s) + H2O(l) 
CO2(s) + 2NaOH(l) 
Al2O3(s) + HCl(g) 
Ga(s) + Te(g) 
SO2(s) + H2O(g) 
Li(s) + H2O(l) 
Na(s) + O2(g) 
Ca(s) + S8(g) 
1. Make up a compound containing 3 elements
1. No subscripts >5
2. Use both even and odd numbers
3. Ex: C3O5F2
2. Calculate the percent composition of your
imaginary compound. Write it down on an
index card.
3. Give your card to another group. See if they can
determine the formula of your compound.