Transcript Chapter 3-part 3

Chapter 3 Atoms and Elements
3.7
Electron Energy Levels
1
Energy Levels
Energy levels
• are assigned numbers n =
1, 2, 3, 4, and so on.
• increase in energy as the
value of n increases.
• are like the rungs of a
ladder with the lower
energy levels nearer the
nucleus.
2
Energy Levels
Energy levels have a maximum number of electrons
equal to 2n2.
Energy level
electrons
n=1
n=2
n=3
Maximum number of
2(1)2 = 2(1) = 2
2(2)2 = 2(4) = 8
2(3)2 = 2(9) = 18
3
Orbitals
An orbital
• is a three-dimensional space around a nucleus,
where an electron is most likely to be found.
• has a shape that represents electron density (not
a path the electron follows).
• can hold up to 2 electrons.
4
Orbitals
An s orbital
•has a spherical shape around the nucleus.
•is found in each energy level.
A p orbital
• has a two-lobed shape.
• is one of three p orbitals in each energy level from n = 2.
5
Orbitals
6
Electron Level Arrangement
In the electron level arrangement for the first 18
elements
• electrons are placed in energy levels (1, 2, 3, etc.),
beginning with the lowest energy level
• there is a maximum number in each energy level.
Energy level
Number of electrons
1
2 (up to He)
2
8 (up to Ne)
3
8 (up to Ar)
4
2 (up to Ca)
7
Examples
Write the electron level arrangement for each:
1. N
2. Cl
3. K
8
Examples
Identify the element with each electron level
arrangement:
1.
2, 2
2.
2, 8, 3
3.
2, 7
9
Electron Configurations of
Multielectron Atoms
Electron Configuration: A description of which orbitals are occupied by
electrons.
Degenerate Orbitals: Orbitals that have the same energy level. For
example, the three p orbitals in a given subshell.
Ground-State Electron Configuration: The lowest-energy
configuration.
Aufbau Principle (“building up”): A guide for determining the filling
order of orbitals.
Electron Configurations of
Multielectron Atoms
Rules of the aufbau principle:
1. Lower-energy orbitals fill before higher-energy orbitals.
2. An orbital can only hold two electrons, which must have
opposite spins (Pauli exclusion principle).
3. If two or more degenerate orbitals are available, follow
Hund’s rule.
Hund’s Rule: If two or more orbitals with the same energy
are available, one electron goes into each until all are half-full.
The electrons in the half-filled orbitals all have the same spin.
Electron Configurations of
Multielectron Atoms
Electron
Configuration
H:
1s1
1 electron
s orbital (l = 0)
n=1
Electron Configurations of
Multielectron Atoms
Electron
Configuration
H:
1s1
He:
1s2
2 electrons
s orbital (l = 0)
n=1
Electron Configurations of
Multielectron Atoms
Electron
Configuration
H:
1s1
He:
1s2
Lowest energy to highest energy
Li:
1s2 2s1
1 electrons
s orbital (l = 0)
n=2
Electron Configurations and the
Periodic Table

Give expected ground-state electron configurations
for the following atoms, draw – orbital filling
diagrams and determine the valence shell
◦ O (Z = 8)
◦ Ca (Z = 20)
◦ Sr (Z = 38)
◦ Sn (Z = 50)
Valence Electrons
The valence electrons
• determine the chemical properties of the elements.
• are the electrons in the highest energy level.
• are related to the group number of the element.
Example: Phosphorus has 5 valence electrons.
5 valence electrons
P in Group 5A(15)
2, 8, 5
17
Groups and Valence Electrons
All the elements in a group have the same number of
valence electrons.
Example: Elements in group 2A (2) have two (2)
valence electrons.
Be
2, 2
Mg
2, 8, 2
Ca
2, 8, 8, 2
Sr
2, 8, 18, 8, 2
18
Periodic Table and Valence Electrons
Representative Elements Group Numbers
1
2
3
4
5
6
7
8
H
He
1
2
Li
2,1
Be
B
2,2 2,3
C
2,4
N
O
F
Ne
2,5 2,6
2,7 2,8
Na
Mg Al Si
P
S
Cl
2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6
Ar
2,8,7 2,8,8
19
Examples
State the number of valence electrons for each.
A. O
1) 4
2) 6
3) 8
B.
C.
Al
1) 13
2) 3
3) 1
Cl
1) 2
2) 5
3) 7
20
Examples
State the number of valence electrons for each.
A. calcium
1) 1
2) 2
3) 3
B. group 6A (16)
1) 2
2) 4
3) 6
C. tin
1) 2
2) 4
3) 14
21
Examples


How many electrons are in energy level 3 of the
following
A. Sulfur
b.
Phosphorous
Examples
Identify the elements that have the following
electron level
Energy level 1
2
3
a.
2
1
0
b.
2
8
6
c.
2
8
7
23
Electron-Dot Symbols
An electron-dot symbol
• shows the valence electrons
around the symbol of the
element.
• for Mg has 2 valence electrons
as single dots on the sides of
the symbol Mg.
.
.
·Mg · or Mg · or ·Mg or ·Mg
·
24
Writing Electron-Dot Symbols
Electron-dot symbols for
• groups 1A (1) to 4A (14) use single dots.
·
·
Na ·
· Mg ·
· Al ·
·C·
·
• groups 5A (15) to 7A (17) use pairs and single dots.
··
··
·P·
·
: O·
·
25
Groups and Electron-Dot Symbols
In a group, all the electron-dot symbols have the
same number of valence electrons (dots).
Example: Atoms of elements in Group 2A (2) each
have 2 valence electrons.
· Be ·
· Mg ·
· Ca ·
· Sr ·
· Ba ·
26
Examples
A.
.
X is the electron-dot symbol for
1) Na
B.
..
.X.
.
1) B
2) K
3) Al
is the electron-dot symbol of
2) N
3) P
27
Examples

What is the dot Lewis structure of
◦
◦
◦
◦
Carbon
Magnesium
Aluminum
Argon
Atomic Size
Atomic size is described
using the atomic radius;
the distance from the
nucleus to the valence
electrons.
29
Atomic Radius Within A Group
Atomic radius increases
going down each group of
representative elements.
30
Atomic Radius Across a Period
Going across a period from left to right,
• an increase in the number of protons increases
attraction for valence electrons.
• atomic radius decreases.
31
Examples
Select the element in each pair with the larger atomic
radius.
A. Li or K
B. K or Br
C. P or Cl
32
Ionization Energy
Ionization energy is the energy it takes to remove a
valence electron.
Na(g) + Energy (ionization) -> Na+(g) + e-
33
Ionization Energy In a Group
Going up a group of
representative elements,
• the distance decreases
between nucleus and
valence electrons.
• the ionization energy
increases.
34
Ionization Energy
• Metals have
lower
ionization
energies.
• Nonmetals
have higher
ionization
energies.
35
Examples
Select the element in each pair with the higher
ionization
energy.
A. Li or K
B. K or Br
C. P or Cl
36