Transcript Lecture 8

Chemistry 103
Lecture 8
Outline

I. Electronic Structure (CH5)
 Brief Review
 Orbital Diagrams
 Valence Electrons
 Atomic & Ionic Radii

II. Chemical Bonds - Ionic Compounds (CH6)
 Electron Dot Symbols
 Lewis Dot Diagrams
 Predicting Formulas
Electron Configuration

Write the full electron configuration and the
nobel gas shorthand electron configuration
for tin (Sn, element 50)
Basic Questions about the Model

Which of the following orbitals will not be found in
our quantum mechanical model?
A. 1s
B. 3p
C. 2f
D. 4p
Justify your answer based on quantum number rules.
When n=2, can only have l=0 (s) and l=1(p) orbitals.
Basic Questions about the Model

What shell will be the first in which a “4” orbital will
be found?
n = 4 , l = 0, 1 ,2, 3
n = 4 (first time l can equal “3” (f))

How many individual orbitals are in the n=2 shell?
1 (s), l = 0, ml = 0
(1 value, 1 orbital)
3 (p), l = 1, ml = -1, 0, 1
(3 orbitals)
TOTAL = 4 orbitals in the n=2 shell
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Orbital (Box) Diagrams

Visual Representations of Electron
Configurations
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Each line (or circle) represents an orbital
Arrows represent electrons
The rules of electron configurations are followed
1s
2s
2p
3s
3p
4s
Orbital Occupancy

The Maximum Number of Electrons any single orbital can
hold is two.
They are distinguished from each other an arrow convention one must be up, one must be down if electrons are in the
same orbital.
He ____
1s
Writing Orbital Diagrams
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Electron configurations tells us which energy levels the
electrons for each element are located.
THREE rules:
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1. Electrons fill orbitals starting with lowest
energy first
2. There can be no more than 2 electrons in any
orbital, and those electrons must have different
“spins”
Writing Orbital Diagrams
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Electron configurations tells us which energy levels the
electrons for each element are located.
THREE rules:

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1. Electrons fill orbitals starting with lowest
energy first
2. There can be no more than 2 electrons in any
orbital, and those electrons must have different
“spins”
3. For orbitals in the same subshell, electrons fill
each orbital singly before any orbital gets a
second electron
Writing Orbital Diagrams
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Nitrogen
___
___
___ ___ ___
1s
2s
2p
Writing Orbital Diagrams
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Unpaired electrons
Nitrogen
___
___
1s
2s
___ ___ ___
2p
Paired electrons
Learning Check
Write the orbital diagrams for:
A. oxygen
B. phosphorus
C. calcium
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Question!

Which electrons are most likely to be involved
in chemical reactions?
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A) those nearest to the nucleus
B) those farthest from the nucleus
C) all are equally likely to be involved
Atomic Families

What differentiates one family from another?
 Let’s look at the electron configurations of the alkali metals to find
out.

Li
Na
K
Rb
Cs
Fr
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The Properties of Atoms Are Dependent on Their
Valence Electrons
• Elements that similar properties have the same number of
valence electrons!

Valence Electrons:
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For elements in the “s” and “p” block - valence electrons are the
electrons in the outer-most electron shell (the shell with the
highest n value)
Example: Mg: 1s22s22p63s2
Example: Cl: 1s22s22p63s23p5
How Many Valence Electrons?

How many valence electrons do the following
elements have?
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Na:
Al:
S:
Cl:
1
3
6
7
Atomic Radius
Atomic radius is the distance from the nucleus to the
valence electrons.
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17
Atomic Radius Within a Group
•
Atomic radius
increases going down
each group of
representative
elements.
Li Atom
Na Atom
K Atom
INCREASE
Radius = distance between nucleus and valence electrons
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Atomic Radius Across a Period
•
Atomic radius decreases from left to right across a period
because more protons increase nuclear attraction for
valence electrons.
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19
Learning Check
Select the element in each pair with the larger atomic
radius.
A.
B.
C.
Li or K
K or Br
P or Cl
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EXAM I - Summer 09
Material
General Course structure
Atoms ---> Compounds ---> Chemical Reactions
Atomic Radius
Atomic radius is the distance from the nucleus to the
valence electrons.
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23
Ionic Radii

An atom (or group of
atoms) that is
electrically charged as
a result of the loss or
gain of electrons.
Sizes of Metal Atoms and Ions
A positive ion
•
Lost its valence
electrons
•
Is smaller (about half
the size) than its
corresponding metal
atom
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Size of Sodium Ion
The sodium ion Na+ is smaller than Na atom because
the valence electrons have been lost from the 3rd
energy level.
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Sizes of Nonmetal Atoms and Ions
A negative ion
•
Increased its number of
valence electrons.
•
Is larger (about twice the
size) than its corresponding
metal atom.
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Size of Fluoride Ion
The fluoride ion F- is larger than F atom because
the added valence electron increase repulsions
between the outer electrons.
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Chapter 6 - Names & Formulas of Compounds

Compounds - Compounds result from the
formation of chemical bonds between two or
more different elements.
Bonds

Chemical bond: attractive force holding
two or more atoms together.
Bonds

Ionic Bonds - electron transfer process.
Typically between a metal and a nonmetals

Covalent Bonds - electrons shared.
Typically involving nonmetals.
Periodic Table of Elements
Electron Dot Symbols
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Focus on valence electrons - highest “n”
quantum number for representative elements
Electron Dot Symbols consists of symbol for
the element with one dot for each valence e-.
Na .
Rules for Electron Dot Symbols
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1. Representative elements in the same
group of the periodic table have the same
number of valence electrons
2. The number of valence electrons is
represented by “dots” around the symbol of
the representative element.
3. The maximum number of valence
electrons for any representative element is 8
- Octet Rule.
The Representative Elements
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Octet Rule - What is special about “8”?
An octet
• Is 8 valence electrons
• Is associated with the stability of the noble gases
• He is stable with two valence electrons (duet).
He 1s2
valence electrons
2
Ne 1s2 2s2 2p6
8
Ar
1s2 2s2 2p6 3s2 3p6
8
Kr
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
8
Ionic and Covalent Bonds
Atoms acquire octets
•
To become more stable
•
By losing, gaining, or
sharing valence electrons
•
By forming ionic bonds or
covalent bonds
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Focus on Ionic Compounds
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Ionic Compound
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Formed between a
positively charged
metal ion and a
negatively charged
nonmetal ion
Metals Form Positive Ions
Metals form
• Octets by losing all of their
valence electrons.
• Positive ions have the electron
configuration of the nearest noble
gas.
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Metal Ions – Why Are They Positive?

Let’s look at Mg. Its electron configuration is
1s22s22p63s2.

If it loses two electrons to become Mg2+, it will have the
electron configuration 1s22s22p6


If it gains 6 electrons to become Mg6-, it will have the
electron configuration 1s22s22p63s23p6
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
8 Valence Electrons!
8 Valence Electrons!
The question is which is easier? To lose TWO electrons or
to gain SIX?

Losing two: Therefore, Mg tends to form Mg2+ ions
Nonmetals Form Negative Ions
Nonmetals form
• Octets by gaining valence
electrons to acquire eight.
• Negative ions have the electron
configuration of the nearest noble
gas
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Publishing as Benjamin Cummings
Nonmetal Ions – Why Are They
Negative?

Let’s look at Cl. Its electron configuration is
1s22s22p63s23p5.

If it loses seven electrons to become Cl7+, it will have the
electron configuration 1s22s22p6


If it gains 1 electron to become Cl-, it will have the electron
configuration 1s22s22p63s23p6
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8 Valence Electrons!
8 Valence Electrons!
The question is which is easier? To lose SEVEN electrons
or to gain ONE?

Gaining ONE: Therefore, Cl tends to form Cl- ions