Chem 115: Chapter 9
Dr. Babb
Periodic Properties of the Atom
Properties that depend on position of element in the periodic table.
Factors that affect the periodic properties:
1.
Principal quantum number of valence shell (nvalence)
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2.
within a column/group.....as proceed down column nvalence increases and shell
size increases.
within a row/period…..as proceed across row nvalence remains constant.
Effective nuclear charge (Zeff): reduced nuclear charge felt by valence ein an atom due to shielding by core e-.
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within column/group…..as proceed down column Zeff remains relatively
constant.
within a row/period…..as proceed across a row Zeff increases.
Zeff = Z - # core e-
Compare the effective nuclear charge felt by the valence electrons in sodium vs. chlorine.
Periodic Properties
A.
Atomic Radius: distance from nucleus to outer limits of atom.
Note: Electron distribution on average is considered to be spherically
symmetric about nucleus.
* Why do atomic radii/sizes increase as
proceed down periodic table?
• Why do atomic radii/sizes decrease as
proceed across periodic table?
Compare the radii of F, C, N, Ne, and O.
B. Ionic Radii
1.
Cation < Neutral Atom
Which is larger, Na+ or Na? Explain.
2.
Anion > Neutral Atom
Which is larger, F- or F? Explain
3.
Isoelectronic species
Relative radii depend on Zeff.
Which is larger, F-, Ne, or Na+?
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Chem 115: Chapter 9
Dr. Babb
C.
Ionization Energy (IE): Energy absorbed when electron is removed from the
valence shell of gaseous atom/ion in ground state.
Note: Harder is to remove e-, the larger the IE.
In general…..
nonmetals, have higher IE relative to metals, which have lower IE.
Example: Consider oxygen.
O(g) O+(g) + 1 eO+(g) O+2(g) + 1 e-
IE1 (First IE)
IE2 (Second IE)
Examples:
Compare the IE of Li, F, and N.
Compare the IE of Ar, He, and Ne.
D. Electron Affinity (EA): Energy released when an electron is added to the valence shell
of a gaseous atom/ion in its ground state.
Note: Easier is to add e-, the more energy released (i.e. more exothermic).
In general…..
nonmetals, have most negative EA relative to metals, which have less negative EA.
Example: Consider oxygen.
1 e- + O(g) O-(g)
1 e- + O- (g) O-2 (g)
EA1 =EA2=+
Examples:
Compare the EA of Mg, Cl, and Si.
Compare the EA of K, Na, and Li.
Of special note:
Ionic Compound = Metal + Nonmetal
(Low IE) (Most neg. EA)
“easy to remove eto form cation”
“easy to add eto form anion”
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Chem 115: Chapter 9
Dr. Babb
Electronegativity and Ionic vs. Molecular Compounds
Ionic Compounds:
• formed by combination of metal cation with nonmetal anion.
• Ions held together by ionic bonds (due to electrostatic attraction between ions of
opposite charge).
Molecular Compounds:
• Formed by combination of two or more nonmetals.
• Atoms held together by covalent bonds (due to sharing of pair of e- between atoms).
Problem…..actual dividing line between ionic and covalent compounds is not so
clear cut!! Many compounds have both ionic and molecular properties.
Linus Pauling (American scientist, Nobel prize in Chemistry 1954)
• Studied the nature of chemical bond.
• Defined concept of electronegativity (en): ability of bonded atom to attract e- in
chemical bond toward itself.
Note: The higher the en, the more the bonded atom pulls e- toward itself. In
general…..metals have lower en, nonmetals have higher en. Exception, Group
VIIIA, en=0.
en=0 for
Group VIIIA
Electronegativity increases
Electronegativity increases
Use of Electronegativity to Determine Bond Type
1.
Look at difference in electronegativity, ∆en = |en2 – en1|.
2.
If ∆en ≥ 2, then ionic bond
If ∆en = 0, then covalent bond.
If 0 < ∆en < 2, then polar covalent bond (bond intermediate between
ionic and covalent and with partial charge separation).
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Chem 115: Chapter 9
Dr. Babb
Examples:
CaO
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What type of bonds are present in each of the compounds
shown below?
(en: Ca=1.1, O=3.5, C=2.5, H=2.1, Cl=2.9, S=2.4, Zn=1.7)
*
If polar covalent bonds are present, indicate which atom
has the partial negative and which the partial positive
charge.
CH4
Cl2
SO4-2
ZnS
Lewis Dot Symbols
Used to keep track of valence electrons.
To draw Lewis Dot Symbols: 1.
2.
Write chemical symbol for element.
Surround by dots, each dot corresponding to a
valence e-.
Remember…..for A-Group elements, # valence e- = Group #.
Examples:
A.
Draw Lewis Dot Symbols for Mg, S, Cl, and Ar.
B.
Use Lewis Dot Symbols to diagram the reaction between Na and F to form the
ionic compound NaF.
C.
In NaF, do the elements obey the octet rule? State the octet rule. Which
elements tend to obey the octet rule? Which elements are exceptions to the octet
rule? Do atoms in molecular compounds also tend to obey the octet rule?
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Chem 115: Chapter 9
Dr. Babb
Lewis Structures
Use Lewis Dot Symbols to show the arrangement of valence e- such
that octet rule is obeyed.
Examples:
A.
Draw Lewis structure for F2, O2, and N2.
B.
Define the following terms: single bond, double bond, triple bond,
bonding e- pair, nonbonding (or lone) e- pair, and bond order.
Lewis Structures: Steps for Drawing
1.
2.
Count up total # valence e-.
Determine skeletal structure (what’s bonded to what).
a.
b.
Central atom is usually written first in chemical formula.
H is never the central atom.
3.
4.
Place 2 e- in each bond.
Surround non-central atoms by octet of e-.
5.
Count up how many e- remain. Place any remaining e- on central atom
as e- pair.
If central atom does not have octet of e-, form multiple bond(s) to it to
give it an octet of e-.
a.
6.
a.
7.
H is exception and is only surrounded by 2 e-.
H, B, Be, F and metals do not generally form multiple bonds.
If central atom in Step #6 does not form multiple bonds, leave it short of
e-.
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Chem 115: Chapter 9
Dr. Babb
Examples:
A.
Draw Lewis structures for the substances given below.
CF4
CH4
H3O+
OCl2
ClO2-
ClO2+
PCl5
B.
CO
CO2
NO3-
HCN
SF4
NCl3
CO3-2
BeCl2
HClO4
BCl3
HNO2
Why can S and P have more than 8 e- in the valence shell? Can elements such
as C, N or O (Period 2 elements) expand their octet? Explain.
Resonance
Must be invoked when one single Lewis structure does not adequately
describe the bonding in a molecule or ion.
Resonance Structures:
1.
2.
Same skeletal arrangement.
Different arrangement of e- pairs.
Example:
A.
Draw a Lewis structure for carbonate. Does this one Lewis structure accurately
describe the bonding in carbonate? The actual carbonate structure is a
resonance hybrid of all reasonable Lewis structures. What is a resonance hybrid?
B.
Draw two reasonable resonance structures for sulfuric acid. Use formal charge
considerations to predict which form is more stable (or preferred).
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Chem 115: Chapter 9
Dr. Babb
Formal Charge
Gives information on how many e- are still associated with an atom
relative to the number of valence e- the atom came in with.
Formal Charge = # Valence e- - # bonds - total # nonbonding eRules for determining stability of resonance structures:
1.
Structure with least amount of formal charge separation is more stable.
2.
Structure with negative formal charge on more en element is more stable.
3.
Structure with two positive or two negative formal charges on adjacent atoms is
NOT stable.
Examples:
A.
Draw resonance structures for N2O. Which structure is most stable? least stable?
B.
Draw resonance structures for HClO3. Which structure is most stable? least
stable?
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