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Transcript
Atoms, Ions, and the Periodic Table
What is an atom?
It is smallest particle of an element that
retains the elements properties.
But how did we come to know all the
information we have about these tiny
particle?
Democritus (460-370 BC)




Matter is made of tiny, solid, indivisible particles which
he called atoms (from atomos, the Greek word for
indivisible).
Different kinds of atoms have different sizes and shapes.
Different properties of matter are due to the differences
in size, shape, and movement of atoms.
Democritus’ ideas, though correct, were widely rejected
by his peers, most notably Aristotle (384-322 BC).
Aristotle was a very influential Greek philosopher who
had a different view of matter. He believed that
everything was composed of the four elements earth, air,
fire, and water. Because at that time in history,
Democritus’ ideas about the atom could not be tested
experimentally, the opinions of well-known Aristotle won
out. Democritus’ ideas were not revived until John
Dalton developed his atomic theory in the 19th century!
John Dalton (1766-1844)





All matter is composed of extremely small
particles called atoms.
All atoms of one element are identical.
Atoms of a given element are different from
those of any other element.
Atoms of one element combine with atoms
of another element to form compounds.
Atoms are indivisible. In addition, they
cannot be created or destroyed, just
rearranged.


Dalton’s theory was of critical importance.
He was able to support his ideas through
experimentation, and his work revolutionized
scientists’ concept of matter and its smallest
building block, the atom.
Dalton’s theory has two flaws:


In point #2, this is not completely true. Isotopes
of a given element are not totally identical;
they differ in the number of neutrons. Scientists
did not at this time know about isotopes.
In point #5, atoms are not indivisible. Atoms
are made of even smaller particles (protons,
neutrons, electrons). Atoms can be broken
down, but only in a nuclear reaction, which
Dalton was unfamiliar with.
Discovery of the Electron
JJ Thomson (1856-1940)



Discovered the electron, and determined that it had
a negative charge, by experimentation with cathode
ray tubes. A cathode ray tube is a glass tube in which
electrons flow due to opposing charges at each end.
Televisions and computer monitors contain cathode
ray tubes.
Thomson developed a model of the atom called the
plum pudding model. It showed evenly distributed
negative electrons in a uniform
positive cage.
Diagram of plum pudding model:
Discovery of the Nucleus
Ernest Rutherford (1871-1937)


Discovered the nucleus of the atom in his
famous Gold Foil Experiment.
Alpha particles (helium nuclei) produced
from the radioactive decay of polonium
streamed toward a sheet of gold foil. To
Rutherford’s great surprise, some of the
alpha particles bounced off of the gold
foil. This meant that they were hitting a
dense, relatively large object, which
Rutherford called the nucleus.
Rutherford then discovered the proton, and next, working with a colleague,
James Chadwick (1891-1974), he discovered the neutron as well.
Models of the Atom - Niehls Bohr


Developed the Bohr model of the atom
(1913) in which electrons are restricted
to specific energies and follow paths called
orbits a fixed distance from the nucleus.
This is similar to the way the planets orbit
the sun. However, electrons do not have
neat orbits like the planets.
Diagram of Bohr model:
Quantum Mechanical Model

This is the current model of the atom.
We now know that electrons exist in
regions of space around the nucleus,
but their paths cannot be predicted.
The electron’s motion is random and
we can only talk about the probability
of an electron being in a certain region.
Sub-Atomic Particles
Each atom contains different numbers of
each of the three SUBatomic particles
Particle Symbol Charge
Molar
Mass
Where
found
Proton
p+
+1
1.007 825
Nucleus
Neutron
n0
0
1.008 665
Nucleus
Electron
e-
-1
0.000 549
Electron
Cloud
“A neutron walked into a bar and asked how much for a drink. The bartender
replied, “For you, no charge.”
Atomic Number
The periodic table is organized in order of increasing
atomic number.
The atomic number is the whole number that is unique
for each element on the periodic table. The atomic
number defines the element. For example, if the
atomic number is 6, the element is carbon. If the
atomic number is not 6, the element is not carbon.
The atomic number represents:


the number of protons in one atom of that element
the number of electrons in one atom of that element (with an ion,
the electrons will be different)
**Therefore, protons = electrons in a neutral
atom**
Atomic Mass




mass of an element measured in amu
(atomic mass units)
all compared to C-12 (the mass of carbon
12, which has a mass of exactly 12 amu
listed on the periodic table
Mass number= #of protons + # of
neutrons
Isotopes




Isotopes are atoms of an element with the same
number of protons but different numbers of
neutrons.
C
Most elements on the periodic table have more
than
one naturally occurring isotope.
There are a couple of ways to represent the different
isotopes. One way is to put the mass after the name or
symbol: Carbon-12 or C-12
Another way is to write the symbol with both the mass
number and atomic number represented in front of the
symbol:
12
6
mass # 12
atomic # 6 C
Determining Average Atomic
Mass


The atomic mass on the periodic table is
determined using a weighted average of
all the isotopes of that atom.
In order to determine the average atomic
mass, you convert the percent abundance
to a decimal and multiply it by the mass of
that isotope. The values for all the
isotopes are added to together to get the
average atomic mass.
Example of Average atomic mass
calculation
Given:
12C = 98.89% at 12 amu
13C = 1.11% at 13.0034 amu
Calculation:
(98.89%)(12 amu) + (1.11%)(13.0034 amu) =
(0.9889)(12 amu) + (0.011)(13.0034 amu) =
12.01 amu
Now you try one:

Neon has 3 isotopes: Neon-20 has a mass of 19.992 amu
and an abundance of 90.51%. Neon-21 has a mass of
20.994 amu and an abundance of 0.27%. Neon-22 has a
mass of 21.991 amu and an abundance of 9.22%. What
is the average atomic mass of neon?
The answer is:
(0.9051)(19.992 amu) + (0.0027)(20.994 amu) +
(0.0922)(21.991 amu) =
20.179 amu
Now compare this mass for Neon to the mass on the
periodic table!

Electromagnetic Radiation



Electromagnetic radiation is a form of energy
that travels through space in a wave-like
pattern. eg. Visible light
It travels in photons, which are tiny particles of
energy that travel in a wave like pattern.
Although we call them particles, they have no
mass. Each photon carries one quantum of
energy.
These photons of energy travel at the speed of
light (c) = 3.00 x 108 m/s in a vacuum
What is a wave and how do we
measure it?


Frequency (ν) – number of waves that passes
a given point per second (measured in Hz)
Wavelength (λ) – shortest distance between
two equivalent points on a wave (measured in
m, cm, nm)
Electromagnetic spectrum (EM)


The electromagnetic spectrum shows all wavelengths of electromagnetic
radiation – the differences in wavelength, energy and frequency
differentiates the different types of radiation.
Note that as the wavelength increases, the energy and the frequency
decrease.
Ground state vs. Excited state



Electrons generally exist in the lowest energy state
they can. We call this the ground state.
However, if energy is applied to the electrons, they
can be “excited” to a higher energy and we call this
an excited state.
The excited state electron doesn’t
stay “excited”. It will fall back to
the ground state quickly. When
the electron returns to the ground
state, energy is released in the
form of light. One example of this
is lasers.
Electrons in Atoms





We are most concerned with electrons because
electrons are the part of the atom involved in
chemical reactions.
Electrons are found outside the nucleus, in a
region of space called the electron cloud.
Electrons are organized in energy levels of
positive integer value (n = 1, 2, 3,...).
Within each energy level are energy sublevels,
designated by a letter: s, p, d, or f.
Each sublevel corresponds to a certain electron
cloud shape, called an atomic orbital.
The electron cloud is like an apartment building.
The
energy levels are like floors in the apartment building.
The sublevels are like apartments
on a floor of the building. Just like
there are different sizes of
sublevels, there are different sizes
of apartments: 1 bedroom, 2
bedroom, etc.
The orbitals are like rooms within an apartment.
The
electrons are like people living in the rooms.
What do these orbitals look like?

The s, p, d and f orbitals look different and
increase in complexity (f-orbitals not shown…
they are very complex)
Number of electrons in each sublevel
depends on number of orbitals!









Each orbital can hold a maximum of 2 electrons.
An “s” sublevel contains 1 s orbital. How many total
electrons can fit in an s sublevel?
2
A “p” sublevel contains 3 p orbitals. How many total
electrons can fit in a p sublevel?
6
A “d” sublevel contains 5 d orbitals. How many total
electrons can fit in a d sublevel?
10
An “f” sublevel contains 7 f orbitals. How many total
electrons can fit in an f sublevel?
14
The Aufbau Principle
Three rules govern the filling of atomic orbitals.
The first is:
 The Aufbau Principle: Electrons enter orbitals
of lowest energy first. The Aufbau order lists
the orbitals from lowest to highest energy:
(“Aufbau” is from the German verb aufbauen:
to build up)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10

The Pauli Exclusion Principle


An atomic orbital may hold at most 2
electrons, and they must have opposite
spins (called paired spins).
When we draw electrons to show these
opposite spin pairs, we represent them
with arrows drawn in opposite directions.
Write this
down in your
notes if you
haven’t!
Hund’s Rule

When electrons occupy orbitals of equal
energy (such as three p orbitals), one
electron enters each orbital until all the
orbitals contain one electron with spins
parallel (arrows pointing in the same
direction). Second electrons then add to
each orbital so that their spins are paired
(opposite) with the first electron in the
orbital.
An electron configuration uses the
Aufbau order to show how electrons are
distributed within the atomic orbitals.
 How to read a segment of an electron
configuration:
Example
3p6
3 = energy level
p = sublevel
6 = # of electrons
Now, let’s look at how to put these together
for a specific element!

Electron Configurations

This is one way to represent the electrons of an
atom. We will try a few together:
Element
Total # of
electrons
Electron Configuration
carbon
fluorine
magnesium
argon
6
1s2 2s2 2p2
9
12
1s2 2s2 2p5
1s2 2s2 2p6 3s2
18
1s2 2s2 2p6 3s2 3p6
Orbital Diagrams
Orbital diagrams show with arrow notation how
the electrons are arranged in atomic orbitals for a
given element.
Element
Total # of
Orbital Diagram
electrons

carbon
.
6
↑↓
1s
↑↓
2s
↑
9
↑↓
1s
↑↓
2s
↑↓
↑↓
2p
↑.
magnesium
12
↑↓
1s
↑↓
2s
↑↓
↑↓
2p
↑↓
argon
18
fluorine
↑↓
1s
↑↓
2s
↑↓
↑
2p
↑↓
2p
↑↓
↑↓
3s
↑↓.
3s
↑↓
↑↓
3p
↑↓
Valence electrons




Electrons in the outer energy level of an atom. They
are like the front lines of an army, because they are
the ones involved in chemical reactions (valence
electrons get shared or transferred during reactions).
The number of valence electrons that an atom has is
directly responsible for the atom’s chemical behavior
and reactivity.
We can represent the number of valence electrons
pictorially by drawing the electrons around the
symbol in a “dot diagram”. The electrons are drawn
in on each side of the symbol and are not paired up
until they need to be.
. Be .
Eg.
Element
Electron Configuration
# Valence
Electrons
Electron Dot
Structure
Li.
Li
1s2 2s1
1
Be
1s2 2s2
2
B
1s2 2s2 2p1
3
C
1s2 2s2 2p2
4
N
1s2 2s2 2p3
5
O
1s2 2s2 2p4
6
F
1s2 2s2 2p5
7
Ne
1s2 2s2 2p2
8
Be
.
.
B.
.̇
. C .
.̇
. N :
.̇
: O :
.̇.
: F :
..̇
.
:
Ne
̇
:
The Periodic Table


The rows on the periodic table are called
periods
The columns on the periodic table are called
groups or families


Elements within a group or a family have similar
reactivity. What do you know about all elements
in a period that could explain this?
They have the same number of valence electrons

Since many of the families on the periodic
table have such similar properties, they
some have specific names that you need
to know. Get out your periodic table and
label each section as we look at them
together.
Halogens
areare
in
group
17
and
they
are
the3most
most
The
Alkali
Alkaline
transition
Metals
Earth
are
Metals
metals
are
include
1
in
and
group
groups
are
2.
the
Noble
Gases
in group
group
18.
They
dothe
not
form
reactive
Theametals
form
-1
ions
byof
gaining 1
reactive
form
through
2+nonmetals.
metals.
12
ions
and
by these
losing
They
form
both
+1
ofall
the
ions
lose
by
ions
because
they
have
full
outer
shell
electron
to
fill
the
highest
energy
p
orbital.
They
1
electrons
and
do
not
need
any
more
electrons.
losing their
electrons
in highest
to
the
formhighest
compounds
energy
energy
s electron.
s orbital.1 2
have do
7 valence
electrons.
They
not
form
valence electrons.
electron.compounds.8 valence electrons
Electromagnetic Radiation



Electromagnetic radiation is a form of energy
that travels through space in a wave-like
pattern. eg. Visible light
It travels in photons, which are tiny particles of
energy that travel in a wave like pattern.
Although we call them particles, they have no
mass. Each photon carries one quantum of
energy.
These photons of energy travel at the speed of
light (c) = 3.00 x 108 m/s in a vacuum
What is a wave and how do we
measure it?


Frequency (ν) – number of waves that passes
a given point per second (measured in Hz)
Wavelength (λ) – shortest distance between
two equivalent points on a wave (measured in
m, cm, nm)
Electromagnetic spectrum (EM)


The electromagnetic spectrum shows all wavelengths of electromagnetic
radiation – the differences in wavelength, energy and frequency
differentiates the different types of radiation.
Note that as the wavelength increases, the energy and the frequency
decrease.
Ground state vs. Excited state



Electrons generally exist in the lowest energy state
they can. We call this the ground state.
However, if energy is applied to the electrons, they
can be “excited” to a higher energy and we call this
an excited state.
The excited state electron doesn’t
stay “excited”. It will fall back to
the ground state quickly. When
the electron returns to the ground
state, energy is released in the
form of light. One example of this
is lasers.
Nuclear Forces

The force that holds the protons together
within the nucleus even though there are
repulsive forces that would otherwise push
the positive protons away from each other.
(also known as strong force)
Radiation


Radiation-it’s the transfer of energy
Radioactivity-The spontaneous emission of
radiation by an unstable nucleus.
Good vs. Bad

Ionizing



Has enough energy to kick off an ion.
Very high energy
Non ionizing


Does not have enough energy to kick off an
ion
Low energy
A. Types of Radiation

Alpha particle ()


electron
Positron (+)


He
2+
paper
Beta particle (-)


helium nucleus
4
2
+’ly charged
e-
0
-1
1-
e
0
1
e
Gamma ()

high-energy photon
0
cardboard
1+
thick concrete
lead
B. Nuclear Decay

Alpha Emission
238
92
parent
nuclide
U
Th  He
234
90
daughter
nuclide
4
2
alpha
particle
Numbers must balance!!
B. Nuclear Decay

Beta Emission
131
53
I
131
54
Xe  e
0
-1
electron
B. Nuclear Decay

Gamma Emission


Usually follows other types of decay.
Transmutation

One element becomes another.
B. Nuclear Decay

Why nuclides decay…

need stable ratio of neutrons to protons
238
92
U
I
131
54
K
38
18
131
53
38
19
106
47
Th  He
234
90
4
2
Xe  e
Ar 
Ag  e 
0
-1
0
-1
0
1
106
46
e
Pd
DECAY SERIES TRANSPARENCY
C. Half-life

Half-life (t½)


Time required for half the atoms of a
radioactive nuclide to decay.
Shorter half-life = less stable.
Fission


splitting a nucleus into two or more
smaller nuclei
1 g of 235U =
3 tons of coal
235
92
U
Fission


chain reaction - self-propagating
reaction
critical mass the minimum
amount of
fissionable
material needed
to sustain a
chain reaction
Fission

Uranium-235 is the only naturally
occurring element that undergoes fission.
Uranium - 235
Fission


Why does fission produce so much energy?
Small quantities of mass are converted into
appreciable quantities of energy.
Fission
Energ
y
1
gram
matter
700,000
Gallons
of high
octane
gasoline
Fusion




combining of two nuclei to form one
nucleus of larger mass
thermonuclear reaction – requires temp of
40,000,000 K to sustain
1 g of fusion fuel =
20 tons of coal
occurs naturally in
stars
2
1
H H
3
1
Fission vs. Fusion
F
I
S
S
I
O
N
 235U




is limited
danger of
meltdown
toxic waste
F
U
S
I
O
N

fuel is abundant
no danger of
meltdown
no toxic waste
Nuclear Power

Fission Reactors
Cooling
Tower
Nuclear Power

Fission Reactors
Nuclear Power

Fusion Reactors (not yet sustainable)
Nuclear Power

Fusion Reactors (not yet sustainable)
National Spherical
Torus Experiment
Tokamak Fusion Test Reactor
Princeton University
Synthetic Elements

Transuranium Elements



elements with atomic #s above 92
synthetically produced in nuclear reactors and
accelerators
most decay very rapidly
238
92
U  He 
4
2
242
94
Pu