Download Atomic structure

Topic 1 - Atomic Structure, Basic Atomic Vocabulary
Write a definition for the following terms in your own words:
Proton –
Neutron –
Electron –
Some vocabulary…
Atomic number – the number of protons in the nucleus of an atom (also the number of
electrons in an electrically neutral atom)
Mass Number – the total number of protons and neutrons in the nucleus of an atom
Atomic mass – mass of an atom of a chemical element expressed in atomic mass units
Atomic mass unit – used to express the masses of individual atoms (one atomic mass
unit (u) is 1/12th the mass of a carbon-12 atom)
Average atomic mass – average of the isotopic masses weighted according to the
naturally occurring abundances of the isotopes of the element (and relative to the value
of exactly 12u for an atom of C-12).
Isotopes – atoms of the same element that contain different numbers of neutrons (i.e.
carbon-14 vs. carbon-12, uranium-235 vs. uranium-238)
Ion – an atom or molecule in which the number of electrons is not equal to the number
of protons, giving the atom or molecule a net charge
 If the ion consists of two or more atoms, it is referred to as a polyatomic ion
 Ions form in order to fill their valence shell. An atom is in its most stable state when
the outermost shell (valence shell) is full. A stable valence shell usually contains 8
electrons, but not always (i.e. hydrogen, helium, lithium, beryllium)
- Cation:
- Anion:
Calculate the average atomic masses
1. Analysis of naturally occurring chlorine shows that it is composed primarily of
two isotopes, Cl-35 and Cl-37. The percentages of each are 75.5% Cl-35 and 24.5%
Cl-37. From this information, calculate the average atomic mass for a chlorine
atom.
2. Naturally occurring copper consists of 69% Cu-63 and 31% Cu-65. From this
information calculate the average atomic mass of a copper atom.
3. Rubidium has two common isotopes, Rb-85 and Rb-87. If the abundance of Rb-85
is 72.2% and the abundance of Rb-87 is 27.8%, what is the average atomic mass of
rubidium?
4. Uranium has three common isotopes. If the abundance of U-234 is 0.01%, the
abundance of U-235 is 0.71%, and the abundance of U-238 is 99.28%, what is the
average atomic mass of uranium?
5. Titanium has five common isotopes, and they are listed here with their
abundances: Ti-46 (8.0%), Ti-47 (7.8%), Ti-48 (73.4%), Ti-49 (5.5%), and Ti-50
(5.3%). What is the average atomic mass of titanium?
Use the atomic numbers and mass numbers to complete the following:
Name
Symbol
Lithium
Li
_______
N
_______
_______
Lead
_______
_______
Ca+2
# of Protons
# of electrons
# of neutrons
3
4
____________
____________
____________
____________
___________
____________
____________
__________
____________
____________
3
___________
24
_______
_______
___________
Sulfur
_______
___________
_______
_______
_______
Ag+
18
____________
____________
____________
____________
____________
___________
____________
____________
85
Silicon
_______
___________
____________
____________
_______
_______
___________
4
____________
_______
Rn
___________
____________
____________
Antimony
_______
___________
____________
____________
_______
_______
___________
____________
0
___________
____________
____________
_______
Am
_______
_______
___________
_______
_______
___________
____________
Curium
_______
___________
____________
____________
_______
_______
____________
____________
_______
Fe
____________
____________
23
___________
84
____________
155
Topic 2 – Trends in the Periodic Table, Lewis Dot, and Valence Electrons
 Based on the periodic law: when the elements are arranged in order of increasing
atomic number, there is a periodic repetition of their physical and chemical properties.
1) Groups: Elements located in the same group share similar chemical and physical
properties. i.e. Na and K are both found in the first group. They would share
similar chemical and physical properties along with Cs, Rb, and the rest of the
group. (Brainiac video…)
2) Atomic Radius:
 In a group, as atomic #, _______________________
 In a period, as atomic #, _______________________
- As you move down a group (therefore, atomic number is ____________), the
number of energy levels ___________. Recall that energy levels refer to the
distance away from the nucleus. Although the energy level increases, the number
of valence electrons _________________. As a result, as you get farther away from
the nucleus, the positive charge of the nucleus is less able to “pull” the valence
electrons closer. Result: the radius _______________ as you move down a group.
- As you move across a period left to right, the atomic number ____________ as
well. However, this time, electrons that are added remain in the same
_____________________ and therefore the same distance away from the nucleus.
Since the atomic number is _____________ while energy level remains constant,
the nucleus exerts a stronger “pull” on the valence electrons. Result: radius
__________________ as you move across a period left to right.
3) Ionization Energy:
 In a group, as atomic #, _________________________
 In a period, as atomic #, _________________________
- As you move down a group, valence electrons are _________________ from the
nucleus and therefore not as tightly bound to it. Therefore, less energy is
required to remove them from the atom to form an ion. Result: the ionization
energy value _________________ as you move down a group.
- As you move across a period left to right, electrons are more tightly bound to
the nucleus due to the increased “pull” of the protons, which are ______________
in number. Therefore there is more energy required to remove them from the
atom to form an ion. Result: the ionization energy value ________________ as
you move across a period left to right.
4) Lewis Dot Diagrams
 G.N. Lewis developed a way to symbolize the valence electrons of an atom. By
using an element’s symbol and a series of dots, we are able to predict bonding
tendencies of atoms. These representations are known as Electron Dot (or Lewis)
Diagrams.
In a Lewis diagram, the element’s symbol represents the nucleus and all nonvalence electrons. The dots surrounding the symbol represent the available
valence electrons.
Reminder/Summary:
Lewis Diagram – a symbolic representation of an atom and its valence electrons
Valence Electrons – electrons in the outermost orbitals of an atom, which are
most often involved in chemical bonding.
i.e.
Carbon
symbol:
valence e-:
Fluorine
symbol:
valence e-:
Rules for drawing Lewis diagrams:
1) Record the name of the element
2) Determine the number of valence electrons
** Use the periodic table! **
3) Write the symbol of the element
4) Place “electron dots” at the directional positions according to the number of
valence electrons available
** ORDER: Up, Down, Left, Right. **
Practice (Human Dot Diagrams!)
Lewis Diagram
1) Beryllium
# valence electrons: ______
2) Oxygen
# valence electrons: ______
3) Argon
# valence electrons: ______
4) Choose the example that shows a correct representation of a Lewis diagram.
a)
b)
c)
5) Complete Lewis Diagrams for the following elements:
a) Nitrogen
b) Fluorine
c) Calcium
d) Sodium
e) Gallium
f) Zinc
g) Xenon
h) Francium
Topic 3: Bohr Diagrams
Niels Bohr was the first scientist to propose the idea of electrons orbiting the nucleus in
shells or energy levels. His model (Bohr Diagrams) expands on the original simple
model of an atom. They show the nucleus with the correct number of protons and
neutrons, and more importantly, a more detailed view of how the electrons are
arranged.
Instructions for drawing a Bohr diagram:
1) Draw the nucleus as a solid circle
2) Put the number of protons (atomic number) in the nucleus with the number of
neutrons (atomic mass – atomic number = # neutrons) under it.
3) Place the number of electrons (same as protons) in orbits around the nucleus by
drawing circles around the nucleus.
When filling shells in energy levels, you have to fill the innermost shell before
beginning to fill the next energy level.
 Maximum number of electrons in 1st energy level: 2
 Maximum number of electrons in 2nd energy level: 8
 Maximum number of electrons in 3rd energy level: 18
 Maximum number of electrons in 4th energy level: 32
Practice drawing the following Bohr diagrams:
a) Na
b) P
c) Cl
d) In
Topic 4: Electron Configuration and Orbital Notation
Electron configuration – “a mathematical way to represent where the electrons are
found in an atom”
 Electrons may be found in 7 principle energy levels
 Each principle level is divided into sublevels of different sizes
First Level: 1 sublevel
Second Level: 2 sublevels
Third Level: 3 sublevels
Fourth Level: 4 sublevels
Fifth Level: 4 sublevels
Sublevels are identified as follows:
Sublevel
Maximum number of electrons
s
2
p
6
d
10
f
14
 When electrons fill up sublevels, they fill the lowest energy levels first. These
electrons are closest to the nucleus. This can be written using the following special
notation…
Ex) Hydrogen (H): 1s1
- 1 means 1st energy level
- s means s-sublevel
- 1 means 1 electron in the sublevel
Ex) Helium (He): 1s2
- 1 means 1st energy level
- s means s-sublevel
- 2 means 2 electrons in the sublevel
Ex) Oxygen (O): 1s2 2s2 2p4
*** Note ***
- You must be careful of the order in which
electrons enter the sublevels. The chart on
the right may help…
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
7d
4f
5f
6f
7f
Noble Gas Configuration
- An even more simplified and shorthand method for representing electron
configuration
- Emphasizes the outermost energy level only
- Instead of listing every energy level and amount of electrons individually, it utilizes
the nearest noble gas element of the energy level below as a representation of the inner
energy levels
For example:
- Sulfur (S)
 Electron configuration would be
1s2 2s2 2p6 3s2 3p4
 Noble gas configuration would be
[Ne] 3s2 3p4
…this is because we know that the electron configuration of Ne is: 1s2 2s2 2p6, therefore
there is no need to write it all out.
- Noble gas configurations are especially useful for elements with a large atomic
number, as their complete electron configurations become tiresome and redundant to
write out each time.
Electron Configuration Practice
In the space below, write the unabbreviated electron configurations of the following elements:
1)
sodium
________________________________________________________
2)
iron
________________________________________________________
3)
bromine
________________________________________________________
4)
barium
________________________________________________________
5)
neptunium
________________________________________________________
In the space below, write the abbreviated electron configurations of the following elements:
6)
cobalt
________________________________________________________
7)
silver
________________________________________________________
8)
tellurium
________________________________________________________
9)
radium
________________________________________________________
10)
lawrencium ________________________________________________________
Determine what elements are denoted by the following electron configurations:
11)
1s2 2s2 2p6 3s2 3p4
________________________________
12)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1
________________________________
13)
[Kr] 5s2 4d10 5p3
________________________________
14)
[Xe] 6s2 4f14 5d6
________________________________
15)
[Rn] 7s2 5f11
________________________________
Determine which of the following electron configurations are not valid:
16)
1s2 2s2 2p6 3s2 3p6 4s2 4d10 4p5
________________________________
17)
1s2 2s2 2p6 3s3 3d5
________________________________
18)
[Ra] 7s2 5f8
________________________________
19)
[Kr] 5s2 4d10 5p5
________________________________
20)
[Xe]
________________________________
Orbital Notation
 Another way used to represent the electrons in an atom
 A series of boxes is used to represent each sublevel
 Each box holds 2 electrons and represents 1 orbital
- an orbital is a space for 2 electrons (which tend to exist in pairs)
s
p
d
f
The way that electrons are placed into these boxes are governed by 3 main rules or
principles:
1) Aufbau Principle
- electrons are put into boxes left to right
2) Pauli Exclusion Principle
- 2 arrows (electrons) per box, 1 up and 1 down
3) Hund’s rule
- when filling boxes, one arrow per box until the set is filled, then second arrow
Ex) Silicon (14 electrons)
1s2
2s2
2p6
3s2
3p2


  

 
Ex) Titanium (22 electrons)
1s2
2s2
2p6
3s2
3p6
4s2


  

  

4d2
 