Download Atomic Structure

Atomic Structure
Mrs. Daniels
September 2002
Chemistry .2
Revised August 2006
IN YOUR JOURNALS…

Describe what you think atoms are made of.

Can you see an atom with your naked eye or
under a microscope?

Draw a picture of a simple atom and its
components.
Dalton’s Atomic Theory
1.
2.
3.
4.
All elements are composed of tiny indivisible
particles called atoms
Atoms of the same element are identical to each
other & are different from those of other
elements
Atoms can physically mix together or
chemically combine in whole number ratios to
form compounds.
Chemical reactions occur when atoms are
separated, joined, or rearranged.
How
BIG or how small is an atom?
Is a penny big or small?
 Imagine grinding up the penny into copper
dust. Is each grain of copper big or small?
 A pure copper penny contains
approximately 2.4 x 1022 individual copper
atoms
 Can you imagine dividing the penny into
24,000,000,000,000,000,000,000 different
parts?

“Cutting it down to size activity”
Pair up with the person sitting next to you
 You and your partner will be given two items: a
pair of scissors and a piece of paper
 Your task: predict how many times you and
your partner can cut this piece of paper in half
with the scissors
 Each time, discard (set aside and we’ll discard at
the end) one half of the paper you’ve cut and
continue on with one piece

“Cutting it down to size activity”
How many times were you able to cut the paper
in half?
 Which pair was able to make the most cuts?
 How many times would you have to divide this
original 8 1/2 x 11” piece of paper in order to
get it to be the width of one atom?

31
Atomic Structure
First of all, an atom has no overall electric
charge.
 Secondly, we know that an equal number of
negative and positive particles combine to
form a neutral particle
 Keeping this in mind, let’s look at three
subatomic particles.

Atomic Structure
Proton (p+):
A positively charged
particle found in the central core of an atom
(called the nucleus)
 Neutron (n0):
A neutral particle found
in the nucleus of an atom
 Electron (e-):
A tiny negatively
charged particle found outside of the atomic
nucleus

Atomic Structure
The mass of a proton and a neutron is
relatively equal
 However, the electron has a mass equal to
1/1840 of a proton


Which subatomic particles do you think
take up the most space in the atom?
Atomic Structure

If we can’t see these subatomic particles,
how do we know they exist?
Before we answer that…
The atoms we’re discussing each have a
representative symbol as you should recall.
 See how many of these elements you can
“discover” in the symbol recognition
worksheet.

Homework tonight:
 Read pages 55-61

J.J. Thomson’s
Cathode Ray Tube

1897 was a big year for J.J…. He discovered the
electron.

Thomson passed electric current through gases
under low pressure in a sealed glass tube
At one end of the glass tube was an electrode with
a positive charge (anode) and at the other end was a
negative electrode (cathode)
When electricity was passed through, a glowing
beam formed between the cathode and the anode.
This beam was then called a cathode ray.


J.J. Thomson’s
Cathode Ray Tube
Thomson found that the cathode ray was
attracted to a positively charged metal plate
 Knowing that opposites attract, he
concluded that the beam (cathode ray) is
made up of tiny negatively charged particles
moving at high speed
 These particles were called electrons

Other Scientists
Millikan:
measured the charge of an
electron
 Moseley:
used an X-ray to determine the
number of protons in an atom
 Rutherford:
used a gold foil experiment to
determine that most of the atom is empty space
and the tiny center of the atom is positively
charged
 Chadwick:
demonstrated the existence of
neutrons

Organization of the Atom
The protons and neutrons are tightly packed
in a central core called the nucleus
 If an atom were the size of a football
stadium, the nucleus would be a tiny marble
sitting in the center of it
 The electrons are found in different layers
(energy levels) of a “cloud” around the
nucleus

Atomic #
& Mass #
Elements differ because of the number of
protons they have
 The Atomic Number is the number of
protons
 The number of electrons in an atom must
equal the number of protons in order for the
atom to be neutral
 The Mass Number is the whole number of
protons plus neutrons in an atom

Atomic # & Mass #
Mass
Number
16
O
Atomic
Number
8
Chemical
Symbol
Mass # vs. Atomic Mass
Atomic
mass
20.18
Ne
Atomic
Number
10
Chemical
Symbol
Isotopes

Isotopes of an atom occur when the number
of neutrons changes

Isotopes have the same chemical properties
as the original atom because the charged
particles remain the same
Atomic Mass
A weighted average mass of the atoms in a
naturally occurring sample of an element is
called the Atomic Mass
 This number represents the mass as well as
the relative abundance of each isotope
 Since atoms are so small, grams are not
typically used as units of mass
 Instead, an Atomic Mass Unit is used
(mathematically defined as 1/12th of the
mass of Carbon-12.)

Bell Work
Calculate this student’s grade if the class is weighted
as follows:
 Tests = 75%
Homework = 5%
 Lab = 10%
Final exam = 10%
Test scores: 89, 84, 72, 90
Lab :
99, 100, 98, 99, 94, 97
Homework : 92, 93, 96, 98, 105, 94
Final exam : 90

Bell Work
Calculate this student’s grade if the class is weighted
as follows:
 Tests = 75%
Homework = 5%
 Lab = 10%
Final exam = 10%
Test scores: 89, 84, 72, 90 = 335/4 = 83.75%
Lab :
99, 100, 98, 99, 94, 97 = 587/6= 97.8%
Homework : 92, 93, 96, 98, 105, 94 = 578/6= 96.3%
Final exam : 90%
83.75(.75) + 97.8(.10) + 96.3(.05) + 90(.10) =

86.4%
B
Bell Work

Now if this teacher did NOT weight grades,
what would this student’s grade be?
Test scores:
Lab :
Homework :
Final exam :
89, 84, 72, 90
99, 100, 98, 99, 94, 97
92, 93, 96, 98, 105, 94
90
1590/1700 = 93.5%
A…very different
Periodic Table of Elements
What do you think of when you hear the
word “periodic”?
 Periodic actually means: occurring on a
regular basis
 There are certain trends that exist on the
periodic table that are consistent

Periodic Table of Elements
A horizontal row across the periodic table
is called a period.
 When you read across the page, you
eventually come the end of a sentence. At
the end of a sentence is a period.


A vertical column on the periodic table is
called a family or group.
Alkali Metals

Group/Family I is called the Alkali Metals
Li
Hydrogen is not included in this group
 It is in a group of its own

Na
K
Rb
Cs
Fr

This family shares certain
characteristics: react vigorously with
water, are metals, and have 1 e- in
their outermost shell
Alkaline Earth Metals
Be
Mg
Ca
Sr
Ba
Ra
The alkaline earth metals are all
metals and all have 2 e- in their
outermost shell
 The second family from the left
of the periodic table

Transition Metals
The transition metals are located in the
center of the periodic table.
 They vary in their number of electrons,
however, they all share in the common
properties of metals.
 ~80% of all of the elements are metals
 The inner transition metals are referred to as
the rare earth elements. These are the two
rows found at the bottom of the periodic
table

Metalloids
Along the zigzag borders are
the metalloids
 These share some properties
of metals (some of the time)
 Aluminum is an exception: it
is a metal

Non-metals

In the upper right hand corner of the
periodic table are the non-metals
Typically non-lustrous and are poor
conductors of electricity
 Halogens (group 7): include chlorine and
bromine
 Noble Gases (group 8): Undergo few or no
chemical reactions

Puzzle Activity Instructions
As a group, you are resonsible for:
 Drawing each missing piece of your puzzle (be
sure to number it) on the white paper
 Guess what design or picture is on the piece and
then draw and color it.
 When you’re finished, give the puzzle to your
instructor. She will give you the puzzle pieces.
 Compare the real pieces to the ones you’ve drawn.
Write down ANY differences.
 In your Journal, what did this activity have to do
with Mendeleev and the first periodic table?

Valence Electrons




The shell or energy level (n) containing the
outermost electrons for an element is called the
valence shell
The electrons in that shell are called valence
electrons
These electrons are the farthest from the atom’s
nucleus and are therefore the easiest to remove
How many valence electrons do each of the alkali
metals have?
Valence Electrons

The similarity in the # of valence electrons
causes members of the same family to share
chemical behaviors

Hydrogen is so tiny, however that it reacts
very differently than other members of its
family
How many valence electrons ?
How many valence electrons do each of the
following have?
 Na
O
C
 Cl
B

Ionization Energy
Some energy is required to remove an
electron from that valence shell
 This energy is referred to as the ionization
energy
 This energy is measured in Volts
 Valence electrons are much easier to
remove than electrons closer to the nucleus
and are therefore usually the only ones
capable of being removed

Octet Rule




We will soon be talking about chemical bonding
One important rule to remember is that atoms tend
to want 8 electrons in their outermost shell
This could mean that they give electrons up, take
on electrons, or share electrons in order to achieve
this goal
Hydrogen & Helium are exceptions…they only
want 2.
Energy Levels or Orbits
Each orbit around the nucleus has a very
specific energy associated with it
 When an element was treated with heat
or an electric current, where did the
energy go?
 The electrons will absorb this energy
 If each energy level is assigned a specific
amount of energy, what does the electron
have to do in order to absorb the extra
outside energy?

Energy Levels or Orbits
It has to jump to the next energy level
located farther from the nucleus
 This is now an “excited electron”
 Excited electrons are very unstable and
cannot remain in the excited state
 They must return to their original orbit or
“ground state”

Energy Levels or Orbits
In order for it to return to its ground state, it
must give off the exact amount of energy it
picked up from the outside source
 When it returns to its ground state, it emits or
gives off the energy in the form of light and heat
 The light emitted by excited electrons in atoms
is not a continuous spectrum (all the colors) but
a line spectrum (only certain wavelengths)
 No two elements have the same line spectrum

Visible Light Spectrum


Reminder: Light is a form of energy
Review from gradeschool:
 ROY G BIV
 Violet light has higher energy than
red light
 There is an inverse relationship
between light wavelengths and
energy
 So as the wavelength of light gets
larger, the energy of light gets
smaller
Other atomic models
As you may recall…
Before Bohr, there was Thompson and
Rutherford
 Thompson proposed that an atom was a ball
of positive charges which contained several
electrons
 Rutherford, with his gold foil experiment,
showed that the bulk of the atom’s mass was
concentrated in a small, positively charged
region called the nucleus

Quantum Mechanical Model
Bohr’s model gave rise to the quantum
mechanical model
 When Bohr proposed that the energy
required to excite an electron (which was
then later emitted) was “quantized”
 There is a specific amount of energy
required in order for an electron to become
excited and move to the next energy level…
 …BUT, each orbit or energy level has its
own requirements. They are not all the
same.

Quantum Mechanical Model
Differing from Bohr’s model, the quantum
mechanical model suggests that the
electrons don’t just follow an exact path
around the nucleus like our planets do
around the sun
 Instead, the true location of the electron is
uncertain and only a probability of its
location is mapped
 This idea lends to the analogy of a cloud
(the more dense the cloud, the higher the
probability of finding the electron there)

Sublevels
Each energy level (n) is made up of one or
more subshells or energy sublevels
 The number of energy sublevels is the same
as the number of the energy level (n)
 So, the 3rd energy level has 3 sublevels; the
5th energy level has 5 sublevels and so on.
 The sublevels are designated s, p, d, and f.

Orbitals
As you proceed beyond the 3rd energy level,
overlapping of sublevels occurs and
becomes more complex as you increase the
energy level number.
 The s sublevels have only one orbital
 The p sublevels have 3 orbitals
 The d sublevels have 5 orbitals
 The f sublevels have 7 orbitals

Orbitals

s orbitals are spherical

p orbitals are dumbbell-shaped with three
different spatial orientations

d orbitals are interesting: 4 of the 5 kinds of d
orbitals are clover-leafed and the fifth has two
opposite nodes with a ring in between

f orbitals are too difficult to visualize
RULES FOR FILLING ATOMIC
ENERGY LEVELS
1.
2.
3.
4.
Electrons fill up the energy sublevels
The lowest energy sublevel must be completely
filled before the next higher sublevel can begin
to be filled. (Aufbau principle)
Each orbital can hold a maximum number of 2
electrons of opposite spin (Pauli exclusion
principle)
Due to their negative charge, electrons repel
one another. They will not pair up in an orbital
of any given sublevel until all orbitals in that
sublevel have been half-filled. (Hund’s rule)
Electron Configuration

There is a pattern that can be used to help you remember
which energy sublevel is next in line:
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
7d
6d
5d
4d
3d
7f
6f
5f
4f
7g
6g
5g
g is theoretical and is
not used in current electron
configurations
Electron Configuration
Remember, the maximum number of
electrons an s sublevel can hold is 2.
 The p = 6
The d = 10
The f = 14

There is an easier way to indicate which
sublevels are filled compared with drawing
out the line diagrams each time
 This is called Electron Configuration

Electron Configuration
Electron configuration is a shorthand way of
showing which orbitals of each sublevel are
filled
 When done correctly, the sum of the
superscripts of all orbitals equals the number of
electrons in the atom
 For example: the electron configuration for
phosphorus is
 P 1s22s22p63s23p3
 Add the superscripts 2+2+6+2+3 =15 e- in P

Exceptions to the rule






Chromium and Copper have unusual electron
configurations
They do not follow Aufbau’s energy diagram
Write down the electron configuration for Cr
What does it end with?
The true electron configuration for Cr is
1s22s22p63s23p64s13d5
Likewise, the electron configuration for Cu ends with 3d10
with only 1 electron in the 4s level