Chapter 5
Periodic Table
Dmitri Mendeleev: Russian Chemist credited with the discovery of the
periodic table.
How did he organize the elements? According to similarities in their
chemical and physical properties. Arranged elements in order of increasing
atomic weight.
What did he notice? When elements are arranged in order of increasing
atomic weight, their chemical and physical properties repeated at regular
intervals.
What did he predict? The existence of elements that had specific
properties.
See Figure 2 on page 134. Compare to Modern Periodic Table.
1
How did Henry Moseley change the periodic table? He discovered atomic
number, and found that when elements are arranged in the periodic table,
their atomic number increases by 1. He then arranged the elements in order
of increasing atomic number, rather than atomic mass.
What is the periodic law? When elements are arranged in order of
increasing atomic number, their chemical and physical properties show a
repeating pattern.
Describe the modern periodic table: Elements with similar properties are
arranged in groups or families. (Columns)
Elements within a group have several things in common:
1. They have similar electron configurations.
2. They have the same number of valence electrons.
3. They form ions with the same charge.
4. They exhibit the same type of chemical behavior.
Valence Electrons: Electrons that are able to be gained, lost, or shared in
the formation of chemical compounds. The s and p electrons in the
outermost energy level.
Core Electrons: All of the other electrons.
Properties of the Representative Elements
S block
Group IA Family Name___Alkali metals__________________________
1. [ ] ns1
2. 1 valence e-
2
3. +1 (losing 1 electron) Na+, K+, Li+ etc.
4. Most Highly reactive metals on the periodic table. So reactive they are
never found in pure form.
Group II A Family Name___Alkaline Earth Metals______________________
1. [ ] ns2
2. 2 valence e3. lose 2 e- to form 2+ ions. Ca2+, Mg2+, Ba2+
4. Also very reactive, but not as much as group IA. Not found in nature in
pure form.
P block
Group IIIA Family Name____Aluminum Family________________
1. [ ] ns2 np1
2. 3 valence e3. lose 3 e- to form 3+ ions. Al3+
4. Not as highly metallic as group IA and IIA. This group also has a
metalloid (B) in it.
Group IVA Family Name_____Carbon Family_________
1. [ ] ns2 np2
2. 4 valence e3. lose or gain 4 e- to form 4+ or 4- ions. (Tends to share e- instead)
4. Lots of variety in this family – nonmetal (C) metalloids (Si, Ge) metals Pb,
Sn.
Group VA Family Name____Nitrogen Family__
1. [ ] ns2 np3
2. 5 valence e-
3
3. gains 3 e- to form 3- ions. N34. Nitrogen is most common element in atmosphere – air is 78% Nitrogen.
Group VIA Family Name_______Chalkogens____(Chalk-formers)___
1. [ ] ns2 np4
2. 6 valence e3. gains 2 e- to form 2- ions.
S2- , O2-
4. Form chalk like compounds (oxides) ex Calcium and magnesium oxides.
Group VIIA Family Name______Halogens _ (salt-formers)_______
1. [ ] ns2 np5
2. 7 valence e3. gain 1 e- to form 1- ions. F-, Cl-, Br4. Most active nonmetals. (They really want to GAIN electrons.) Form salt
compounds with active metals. NaCl. MgCl2 KCl CaCl2
Group VIIIA Family Name______Noble Gases___________
1. [ ] ns2 np6
2. 8 valence e3. No ions formed since configuration is stable as is.
4. Inert or nonreactive
Rule of Octet: Atoms tend to gain, lose, or share electrons in order to
acquire a complete set of valence electrons. (s2p6 =8) or just s2 if the atom
is in period 1 or 2.
4
Periodic Trends
Periodic Trends are: Systematic variations in properties of elements that change
in a predictable way as you move through the periodic table.
Periodic Trends we will study:
1. Atomic Radius: distance from the center of the nucleus to the outermost
valence e- of an atom.
2. Ionic Radius: distance from the center of the nucleus to the outermost
valence e- of an ion.
3. Ionization Energy: Energy required to remove the outermost electron. This
energy must always be put in to remove the electron. (Endothermic)
4. Electron Affinity: The energy change when an atom takes on an additional
electron, becoming an anion. This energy change can be an energy put in (if the
atom doesn’t want the electron- Endothermic) or Released- (exothermic if the
atom wants to take the electron.)
5. Electronegativity: The tendency for an atom to attract an electron in a
chemical bond.
1. Atomic Radius
http://2012books.lardbucket.org/books/beginning-chemistry/section_12/f07b1af3eb2bc0e5cffb096bdfa3c5dd.jpg
a. What is the trend as you move from left to right across a period?
b. Explain a: radius decreases because e- are added to same energy
level and there are more p+ pulling energy level in.
c. What is the trend as you move from top to bottom down a family?
d. Explain c: radius increases because there are more energy levels.
5
2. Ionic Radius
http://employees.csbsju.edu/cschaller/Principles%20Chem/ionics/PTionicradii.png
a. What is the trend as you move from left to right across a period?
Decreases until group IVA, then big increase. Then a decrease again.
b. Explain a: group IA to IV A are cations (they are smaller than the
atoms that form them) Groups VA to VIIA form anions (they are larger
than the atoms that form them)
c. What is the trend as you move from top to bottom down a family?
Increases, like atomic radii.
d. Explain c: more electron energy levels.
e. What does isoelectronic mean?
Same # of electrons (same electron configuration)
f. Suppose an anion and a cation are isoelectronic; which has a larger
radius? anion
g. Explain f: The higher the p+ : e- ratio, the SMALLER the radius.
h. Copper forms two different cations; Cu2+, and Cu+. Which has a larger
radius?
i. Explain h: 2+ has a greater P+ to e- ratio, so it is smaller
3. Ionization Energy
http://www.angelo.edu/faculty/kboudrea/periodic/trends_ionization_energy_fig3.gif
http://www.avon-chemistry.com/p_table_ionization.jpg
A. ionization energy can be thought of as a reflection of the attraction of
the nucleus for the outermost valence e-.
B. Atoms with High Ionization Energy hold on to their valence
Electrons very tightly.
C. Examples of groups/ families with high ionization energies:
VIIA, VIIIA
D. Atoms with low ionization energies are likely to : Give up valence e- s
easily
6
E. Examples of groups/ families with low ionization energy? Group IA, IIIA
(IIA has s2 and sort of wants to hang on to the filled s sublevel)
F. What is the trend in ionization energy as you move from left to right
across the periodic table? Increases; radius smaller, nucleus holding
tighter to valence e-s.
G. Successive Ionizations of Al
https://thomsonscience.files.wordpress.com/2011/09/screenshot-dataforionizationenergies-3rdperiod.png
1. First Ionization Energy: Energy required to remove outermost efrom neutral atom. Al + IE 1  1e- + Al+
IE 1 = 578 kJ/mol (3p1)
2. Second ionization Energy: Energy required to remove next (2nd)
e- from ion with a +1 charge...
Al+
+ IE 2  1e- + Al2+
IE 2 = 1817 kJ/mol (3s2)
3. Third Ionization Energy: Energy required to remove next (3rd) efrom ion with a +2 charge...
Al2+
+ IE 3  1e- + Al3+
IE 3 = 2750kJ/mol (3s1)
___________________________________________________________
__________________________________________________________
BIG JUMP IS HERE because the 4th ionization is taking a CORE eAl3+
+ IE 4  1e- + Al4+
IE 4= 11,600kJ/mol (2p6)
H. See table 3 on p. 155.
1. Where is the largest jump in successive ionization energies of
Aluminum? After the third ionization
2. Which electron is being removed in the first ionization of Al?
3p1
3. Write the reaction for 2: Al + IE1  1e- + Al+
4. Which electron is being removed in the 2nd ionization of Al? 3s2
The 3rd? 3s1 The fourth? 2p6
7
5.
Write the reactions for the 2nd, 3rd, and 4th ionizations:
2nd: Al+ + IE2  1e- + Al2+
3s2
rd
2+
3+
3 : Al + IE3  1e + Al
3s1
4th: Al3+ + IE4  1e- + Al4+
2p6
6. Summary
In successive ionizations of an element, the largest jump in
energy occurs after: the last valence e has been removed
The large energy jump indicates: the e- comes from the stable
p6 core.
7. Write the reactions for the successive ionizations of
Magnesium. See table 3 on p 155 for Ionization energies.
4. Electron Affinity (EA)
http://faculty.sdmiramar.edu/fgarces/zCourse/All_Year/Ch100_OL/aMy_FileLec/04OL_LecNotes_Ch100/03_AtomsEleme
nts/306_PeriodicTable/306_pic/electronaffinity.gif
A. EA values can be positive or negative.
1. If EA is positive: atom doesn’t want to take on the electron,
therefore energy must be provided to make it take the e-. Also, the
atom is going to a higher energy state (less stable)
2. What type of element has a positive EA? Elements that don’t
want to be anions. Metals, Noble gases.
3. If EA is negative: atom wants to take on the electron. Energy is
released as a result of the atom going to a more stable (lower
energy) state.
4. What kind of element has a negative EA? Elements that want
to be anions. Halogens, other nonmetals.
8
B.
In general, what is the trend in EA as you move down a
group? Gets more positive. The more metallic the element, the
more it wants to lose electrons, the less it wants to take them.
6. In general, what is the trend in EA as you move from left to right
across a period? Gets More negative. UNITL THE NOBLE
GASES! Then it gets + again.
7. What group/ family has the most highly negative EA?
Halogens. Why? They have p5 configuration.
5. Electronegativity
http://www.biog1445.org/media/Electronegativity_files/electroneg.increasing.gif
A. This property is different from the other periodic trends since it
focuses on the elements behavior in a chemical bond.
B. Scale of Electronegativity: Linus Pauling Scale: 0 - 4
1. High electronegativity: F = 4.0
2. Low electronegativity: 0.7 = Cs
c. What happens to electronegativity as you go from left to right across
the periodic table? Increases
D. What happens to electronegativity as you go down a group/ family?
decreases
E. Which element is the most electronegative element? F
F. Which element is the least electronegative? Cs
G. Memorize the electronegativity of second period of elements, plus H,
Cs, and Cl (3.0)
9
10