View
219
Download
0
Category
Tags:
Preview:
Citation preview
2
3
Interpreting the Periodic Table
4
Interpreting the Periodic Table
1. Typically they have a shiny luster.
2. Relatively high density.
3. Malleable ( they can be hammered into thin sheets).
4. Ductile (they can be drawn into a wire).
5. Good conductors of electricity.
6. Reflect light and heat.
7. High melting and boiling points so they are solids at room temperature (except Hg).
8. Lose electron(s) forming positive ions (cations).
9. Combine with non-metals.
10. Do not readily combine with each other.
General Properties of Metallic Elements:
Examples of Metals
Potassium, K reacts with water and must be stored in kerosene
Zinc, Zn, is more stable than potassium
Copper, Cu, is a relatively soft metal, and a very good electrical conductor.
Mercury, Hg, is the only metal that exists as a liquid at room temperature
11 p11 e
11 p10 e
• Metals lose electrons forming positive ions.
• The radius of the cation is always smaller than the radius of the parent atom.
1. Poor conductors of heat and electricity.
2. Not malleable or ductile; fragile.
3. Low densities.
4. Low melting and boiling points so they can be gases, liquids, or solids.
5. Gain electron(s) forming negative ions (anions).
6. Combine with metals.
7. Combine with each other to a limited extent.
General Properties of Non-metallic Elements:
17 p17 e
17 p18 e
• Non metals gain electrons forming negative ions.
• The radius of the anion is always larger than the radius of the parent atom.
Metalloids: elements that lie in the colored stair step line of the Periodic Table. They have both metallic and non-metallic properties.
These elements are weak conductors of electricity, which makes them useful semiconductors in the integrated circuits of computers.
11
Physical state of elements
12
Chemical Periodicity
13
Periodic Properties of the ElementsAtomic Radii
Atomic radii describes the relative sizes of atoms. It is understood as the distance from the nucleus to the outermost occupied energy level.Atomic radii increase within a column going from the top to the bottom of the periodic table.Atomic radii decrease within a row going from left to right on the periodic table.
14
Atomic Radii
Example: Arrange these elements based on their atomic radii. Se, S, O, Te
15
Atomic Radii
Example: Arrange these elements based on their atomic radii. P, Cl, S, Si
16
Atomic Radii
Example: Arrange these elements based on their atomic radii. Ga, F, S, As
17
Ionization Energy
First ionization energy (IE1) The minimum amount of energy required to remove the
most loosely bound electron from an isolated gaseous atom to form a 1+ ion.
Symbolically:Atom(g) + energy ion+
(g) + e-
Mg(g) + 738kJ/mol Mg+ + e-
18
Ionization Energy
Second ionization energy (IE2) The amount of energy required to remove the
second electron from a gaseous 1+ ion.
Symbolically: ion+ + energy ion2+ + e-
Mg+ + 1451 kJ/mol Mg2+ + e-
•Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies.
19
Ionization EnergyPeriodic trends for Ionization Energy:
1. IE2 > IE1
It always takes more
energy to remove a second electron from an ion than from a neutral atom.
2. IE1 generally increases moving from left to right in the same period.
3. IE1 generally decreases moving down a group.
IE1 for Li > IE1 for Na, etc.
20
Ionization Energy
Example: Arrange these elements based on their first ionization energies. Sr, Be, Ca, Mg
21
Ionization Energy
Example: Arrange these elements based on their first ionization energies. Al, Cl, Na, P
22
Ionization Energy
Example: Arrange these elements based on their first ionization energies. B, O, Be, N
23
Ionization EnergyGroup
and element
IANa
IIAMg
IIIAAl
IVASi
IE1 (kJ/mol)
496 738 578 786
IE2
(kJ/mol)4562 1451 1817 1577
IE3
(kJ/mol)6912 7733 2745 3232
IE4
(kJ/mol)9540 10,550 11,580 4356
24
Ionization Energy
The reason Na forms Na+ and not Na2+ is that the energy difference between IE1 and IE2 is so large. Requires more than 9 times more energy to
remove the second electron than the first one.
The same trend is persistent throughout the series. Thus Mg forms Mg2+ and not Mg3+. Al forms Al3+.
25
Ionization Energy
Example: What charge ion would be expected for an element that has these ionization energies?
IE1 (kJ/mol) 1680
IE2 (kJ/mol) 3370
IE3 (kJ/mol) 6050
IE4 (kJ/mol) 8410
IE5 (kJ/mol) 11020
IE6 (kJ/mol) 15160
IE7 (kJ/mol) 17870
IE8 (kJ/mol) 92040
Notice that the largest increase in ionization energies occurs between IE7 and IE8. Thus this element would form a 1- ion.
26
Ionization Energy
Example: What charge ion would be expected for an element that has these ionization energies?
IE1 (kJ/mol) 1680
IE2 (kJ/mol) 3370
IE3 (kJ/mol) 11586
IE4 (kJ/mol) 12410
IE5 (kJ/mol) 13020
IE6 (kJ/mol) 15160
IE7 (kJ/mol) 17870
IE8 (kJ/mol) 18040
Notice that the largest increase in ionization energies occurs between IE2 and IE3. Thus this element would form a 2+ ion.
27
Ionization Energy
Example: What charge ion would be expected for an element that has these ionization energies?
IE1 (kJ/mol) 1680
IE2 (kJ/mol) 3370
IE3 (kJ/mol) 4586
IE4 (kJ/mol) 5410
IE5 (kJ/mol) 6020
IE6 (kJ/mol) 7160
IE7 (kJ/mol) 17870
IE8 (kJ/mol) 18040
Notice that the largest increase in ionization energies occurs between IE6 and IE7. Thus this element would form a 2- ion.
28
Ionic Radii
Cations (positive ions) are always smaller than their respective neutral atoms.Element Li Be
Atomic Radius (Å)
1.52 1.12
Ion Li+ Be2+
Ionic Radius (Å)
0.90 0.59
29
Ionic Radii
Cations (positive ions) are always smaller than their respective neutral atoms.
Element Na Mg Al
Atomic Radius (Å)
1.86 1.60 1.43
Ion Na+ Mg2+ Al3+
Ionic Radius (Å)
1.16 0.85 0.68
30
Ionic Radii
Anions (negative ions) are always larger than their neutral atoms.
Element N O F
AtomicRadius(Å)
0.75 0.73 0.72
Ion N3- O2- F1-
IonicRadius(Å)
1.71 1.26 1.19
31
Ionic Radii
32
Ionic Radii
Cation (positive ions) radii decrease from left to right across a period. Increasing nuclear charge attracts the electrons and
decreases the radius.
Ion Rb+ Sr2+ In3+
IonicRadii(Å)
1.66 1.32 0.94
33
Ionic Radii
Anion (negative ions) radii decrease from left to right across a period. Increasing electron numbers in highly charged ions
cause the electrons to repel and increase the ionic radius.
Ion N3- O2- F1-
IonicRadii(Å)
1.71 1.26 1.19
34
Ionic Radii
Example: Arrange these elements based on their ionic radii. Ga, K, Ca
35
Ionic Radii
Example: Arrange these elements based on their ionic radii. Cl, Se, Br, S
36
Electronegativity
Electronegativity is a measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another element. Electronegativity is measured on the Pauling
scale. Fluorine is the most electronegative element. Cesium and francium are the least
electronegative elements.For the representative elements, electronegativities usually increase from left to right across periods and decrease from top to bottom within groups.
37
Electronegativity
38
Electronegativity
Example: Arrange these elements based on their electronegativity. Se, Ge, Br, As
39
Electronegativity
Example: Arrange these elements based on their electronegativity. Be, Mg, Ca, Ba
See animation for Summary of Trends in the Periodic Table
http://www.learnerstv.com/animation/animation.php?ani=56&cat=chemistry
40
Recommended