Transcript chapter 5 electronx

Chapter 5
Electrons in Atoms
Bohr
In 1913 Bohr published a theory about
the structure of the atom based on an
earlier theory of Rutherford's.
Rutherford had shown that the atom
consisted of a positively charged
nucleus, with negatively charged
electrons in orbit around it
Bohr’s “Quantum Mechanical
Model” of the atom
• Bohr expanded upon this theory by proposing
that electrons travel only in certain
successively larger orbits
• He suggested that the outer orbits could hold
more electrons than the inner ones
• these outer orbits determine the atom's
chemical properties
Organizing atoms in the periodic table
• The Periodic Table: organizes elements
by atomic number and…
•
Groups/families: elements have the
same physical and chemical properties.
•
Rows/periods: elements have the
same number of electron shells.
Practice Question 1
1. Name another element that would have
similar chemical properties to chlorine.
2. Name an atom that is in the same
period as chlorine.
Electrons
• All atoms have an equal number of
protons and electrons
– Atoms are electrically neutral
• Atoms have no charge
• Symbol: Ne
An equal number of
positive protons and
negative electrons
results in zero charge
Practice Question 2
• How many electrons orbit:
– A magnesium atom?
– A sulfur atom?
– A hydrogen atom?
Valence electrons
are:
• responsible for chemical behavior of atom
• used for chemical bonding
• located in the outer electron shell
1 valence e-
4 valence e-
study question 5
Total number of
electrons
nitrogen
phosphorus
Number of valence
electrons
• There are 7 possible energy levels where an
electron can be found.
• Energy levels are represented by the periods
(horizontal rows) on the periodic table
• The total number of electrons that can fit in an
energy level can be found using the equation:
2n2 (where n = an energy level 1-7)
• Within each energy level, there are
sublevels
– s sublevel- consist of e- from groups 1 and 2
– p sublevel- consist of e- from groups 13-18
– d sublevel- consists of e- from groups 3-12
– f sublevel- consists of e- from the lanthanide
and actinide series
s
p
1
2
3
4
5
6
7
d (n-1)
f (n-2) 6
7
Orbitals
• Each sublevel can be broken down into
orbitals:
– s sublevel: has 1 orbital
– p sublevel: has 3 orbitals
– d sublevel: has 5 orbitals
– f sublevel: has 7 orbitals
• Each orbital can only hold a maximum of 2
electrons.
Orbitals
Energy
Level
Sublevels
Total Orbitals
Total
Electrons
Total Electrons
per Level
n=1
s
1 (1s orbital)
2
2
n=2
s
p
1 (2s orbital)
3 (2p orbitals)
2
6
8
n=3
•
•
s
1 (3s orbital)
2
18
Complete
the3chart
in your notes as6 we discuss this.
p
(3p orbitals)
(3d orbitals)
10 It has only 1.
Thedfirst level5 (n=1)
has an s orbital.
There are no other orbitals in the first energy level.
n = •4 We scall this orbital
1 (4s orbital)
32
the 1s orbital. 2
p
d
f
3 (4p orbitals)
5 (4d orbitals)
7 (4f orbitals)
6
10
14
Where are these Orbitals?
http://www.biosulf.org/1/images/periodictable.png
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
4f
5f
Electron configurations
• Electron configurations are similar to postal
“zipcodes”.
– They represent a general area where an electron can
be found.
• Examples
Hydrogen has 1 electron: 1s1
He has 2 electrons: 1s2
Li has 3 electrons: 1s2 2s1
Be has 4 electrons: 1s2 2s2
B has 5 electrons: 1s22s22p1
Electron Configurations
Rules for Electon Configurations
• 3 rules govern electron
configurations.
– Aufbau Principle
– Pauli Exclusion Principle
– Hund’s Rule
Each line represents
an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
High Energy
Low Energy
 The Pauli Exclusion Principle
states that an atomic orbital
may have up to 2 electrons and
then it is full.
Wolfgang Pauli, yet
another German
Nobel Prize winner
Don’t pair up the 2p electrons
until all 3 orbitals are half full.
Element
Configuration
Element
Configuration
H Z=1
1s1
He Z=2
1s2
Li Z=3
1s22s1
Be Z=4
1s22s2
B
Z=5
1s22s22p1
C
Z=6
1s22s22p2
N Z=7
1s22s22p3
O
Z=8
1s22s22p4
F Z=9
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K Z=19
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
Fe Z=26
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
Note that all the numbers in the electron configuration add up to the atomic
number for that element. Ex: for Ne (Z=10), 2+2+6 = 10
Electron Configurations
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1
This similar configuration causes them to behave the
same chemically.
It’s for that reason they are in the same family or group
on the periodic table.
Each group will have the same ending configuration, in
this case something that ends in s1.