Transcript Hein and Arena - Solano Community College

Electrons in Atoms
and the Periodic Table
Chapter 9
Tro, 2nd ed.
A Brief History of Atomic Theory
Greeks were the first to suggest that matter
was made up of small particles called
atoms
Early chemists performed experiments to
find these particles
Their experiments led to Dalton’s Atomic
Theory
Dalton’s Atomic Theory was revised further
to the Thompson and Rutherford
“nuclear” models of the atom
Further work has led to the modern nuclear
atomic theory, which is still being revised
A Brief History of Atomic Theory
~1900 Max Planck: energy emitted in
bursts called quanta
~1905 Einstein used quanta to explain
properties of light
Called quantum of light a photon
~1911 Rutherford proposed “nuclear atom”
Nucleus is a tiny center in the atom:
Contains most of the mass
Contains all of the positive charge
(Thought electrons in orbit like mini-solar
system - not!)
WHAT YOU NEED TO KNOW ABOUT
ENERGY AND LIGHT (& e-s & photons)
Light is part of the electromagnetic
spectrum of radiation
Continuous spectrum is like a “rainbow”
Discrete spectrum is lines of specific colors
of light
Each line of color has its own energy
associated with it
Light in a vacuum moves at constant speed
c = 2.998 X 108 m/sec (memorize!)
WHAT YOU NEED TO KNOW ABOUT
ENERGY AND LIGHT (& e-s & photons)
Electromagnetic Radiation - Examples
light from the sun
x-rays
microwaves
radio waves
television waves
radiant heat
All show wavelike behavior.
Each travels at the same speed in a
vacuum:
2.998 x 108 m/s
X-rays are part visible
of the light
Infrared
is partlight
of is
electromagnetic
the electromagnetic
part of the
spectrum
spectrum
electromagnetic
spectrum
The Electromagnetic Spectrum
Know visible
light as
ROYGBIV.
Characteristics of a Wave
Light has the properties of a wave.
Wavelength,
l
Wavelength,
l
(measured
from
(measured
from
peak totrough
peak) to trough)
Characteristics of a Wave
Frequency, n, is the number of wavelengths that pass
a particular point per second.
Characteristics of a Wave
Speed is how fast a wave moves through space.
WHAT YOU NEED TO KNOW ABOUT
ENERGY AND LIGHT (& e-s & photons)
Each specific “color” of visible light has a specific wavelength
called l and a specific frequency called n and both relate
to speed of light
c = ln
Wavelength, l, is length unit: meters/wavelength
Frequency, n, is number of wavelengths/second
Speed of light: c = meters
*
wavelength
wavelengths = m/s
second
Energy (Joules) of one photon of light of specific wavelength
or frequency:
E = hn = Joules/photon
h = Planck’s constant (6.626 x 10-34 J.s)
The Bohr Atom
At high temperatures or voltages,
elements in the gaseous state emit
light of different colors.
When the light is passed through a
prism or diffraction grating a line
spectrum results.
Niels Bohr, a Danish physicist,
in 1912-1913 carried out research
on the hydrogen atom.
Each element has its own
unique set of spectral emission
lines that distinguish it from
other elements.
These colored lines
indicate that light is
being emitted only at
certain wavelengths.
Line spectrum of hydrogen. Each line corresponds
to the wavelength of the energy emitted when the
electron of a hydrogen atom, which has absorbed
energy falls back to a lower principal energy level.
Tro: Figure 9.8: each element produces its own unique emission spectrum.
Spectra
Energy is absorbed by e-, then
emitted as a photon of light.
Tro:
Figure
9.11
The Bohr
Atom
E1
E2
E3
An electron can have
one of several possible
energies depending on
its orbit.
The Bohr
Atom
Electrons
revolve
An
electron
has a
around the
nucleus
in it
discrete
energy
when
orbits thatan
areorbit.
located
occupies
at fixed distances from
the nucleus.
Different lines of the
hydrogen spectrum
correspond to different
electron energy level shifts.
The Bohr
Atom
When
an electron
falls
The color
of the light
from
a higher
energy level
emitted
corresponds
to
to
a lower
one
of theenergy
lines oflevel
the a
quantum
energy in the
hydrogenofspectrum.
form of light is emitted by
the atom.
The Bohr
Atom
Light is not emitted
continuously, but is
emitted in discrete
packets called photons
CALCULATIONS FOR l, n AND E OF A
PHOTON EMITTED BY AN ELECTRON
When e- in H atom moves from energy level
1 up to energy level 4 and then drops back
down to energy level 2, we see a photon of
light emitted that has a wavelength, l, of
4.86 x 10-7 m. Calculate the frequency of
the light and the energy of the photon
emitted.
c = ln Rearrange to get frequency, n
n = c/l = (2.998x108 m/s)/(4.86x10-7 m)
= 6.17 x 1014 /s
E = hn = 6.626x10-34 J.s * 6.17 x 1014 /s
= 4.09 x 10-19 J/photon of light
MODERN ATOMIC THEORY
Bohr’s calculations succeeded very
well for the hydrogen atom.
Bohr’s methods did not succeed for
heavier atoms.
So…more theoretical work on atomic
structure was needed.
MODERN ATOMIC THEORY
Thompson had shown that light, which is
photons of energy, had the properties of
matter as well.
In 1924, Louis De Broglie suggested that all
matter must also have wave properties.
De Broglie showed that the wavelength of
ordinary sized objects, such as a baseball,
are too small to be observed.
For objects the size of an electron the
wavelength can be detected.
MODERN ATOMIC THEORY
In 1926, Schröedinger created a
mathematical model that showed electrons
as waves.
Schröedinger’s work led to a new branch of
physics called wave or quantum mechanics.
Using Schröedinger’s wave mechanics, the
probability of finding an electron in a
certain region around the atom can be
determined.
The actual location of an electron within an
atom cannot be determined (Heisenberg
Uncertainty Principle).
MODERN ATOMIC THEORY
Based on wave mechanics it is clear
that electrons are not revolving
around the nucleus in orbits.
Instead of being located in orbits, the
electrons are located in orbitals.
An orbital is a region around the
nucleus where there is a high
probability of finding an electron.
MODERN ATOMIC THEORY
According to Bohr the energies of
electrons in an atom are
quantized.
The wave-mechanical model of
the atom also predicts discrete
principal energy levels within
the atom.
Energy Levels
of Electrons
Principal energy levels, n = 1,2,3...(also called
principal quantum number)
Each energy level contain(s) sublevel(s) called s,
p, d, and f (the angular momentum quantum
number, l)
Within sublevels are orbitals (designated by
orientation in 3-D space, called the magnetic
quantum number, ml)
Each orbital can hold 2 e-s max (each electron is
assigned a spin quantum number, ms)
As n increases, the
energy of the electron
increases
The first four principal
energy levels of the
hydrogen atom
Each level is
assigned a principal
quantum number, n
Energy levels in atoms
Energy levels in
atoms
The number of sublevels
equals the assigned energy
level. For n=1, there is
one sublevel, s. For n=2,
there are two sublevels, s
& p, etc. The sublevels
have the quantum
designation, l. The
maximum value is n -1.
Each principal energy level
is subdivided into sublevels.
Within sublevels the electrons are found
in orbitals.
An s orbital is round and
soft, like a nerf ball.
The shape represents the
highest probability where the
electron might be found.
An atomic orbital can hold a maximum of
two electrons.
An electron can spin in one
of two possible directions
represented by ↑ or ↓.
The two electrons that
occupy an atomic orbital
must have opposite spins.
This is known as the Pauli
Exclusion Principal.
A p sublevel is made up of three p-type
orbitals.
Each p orbital has two lobes and can hold a
maximum of two electrons.
Since there are three orbitals, a p sublevel can
hold a maximum of 6 electrons.
The three p orbitals all center at
the atom’s nucleus…
pz
py
px
…and occupy one of the
three axes of 3-D space.
A d sublevel is made up of five orbitals.
The five d orbitals lie in different planes and point in different
directions. Each d orbital can hold a maximum of two electrons.
A d sublevel can hold a maximum of 10 electrons.
Tro Figure 9.19: The number of subshells within a shell is equal to the
value of n, the principal quantum number.
Shells & Subshells
Etc, etc, etc: What subshells exist for n = 5?
Energy Levels
of Electrons
Pauli Exclusion Principle: Each
orbital can hold a max of 2 e-s, so
possibilities are 0, 1 or 2 e-s.
s has only 1 orbital  2 e-s MAX
p has 3 orbitals  6 e-s MAX
d has 5 orbitals  10 e-s MAX
f has 7 orbitals  14 e-s MAX
(g has ? Orbitals  ?? e-s MAX)
Energy Levels
of Electrons
Max is 2 e-s because of a property of
e-s called spin. (Pauli Exclusion
Principle)
Each e- is spinning on its axis like the
earth.
Any spinning charge creates a
magnetic field, with N and S poles.
The direction of spin determines
which is North.
The e-s will pair up in an orbital so
their N poles are opposite each other.
Atomic Structure of the First 18
Elements: Use these guidelines
The ground state of the electron is the lowest
energy orbital it can occupy.
Higher energy orbitals are excited states.
The distribution of electrons into the various
energy shells and subshells in an atom in its
ground state is called its electron
configuration.
Each energy shell and subshell has a maximum
number of electrons it can hold:
s = 2, p = 6, d = 10, f = 14
Place electrons in the energy shells and
subshells in order of energy, from low energy
up: the Aufbau Principal.
Atomic Structure of the First 18
Elements: Use these guidelines
No more than two electrons can occupy one orbital
Electrons occupy the lowest energy orbitals
available, the ground state. They enter a higher
energy orbital only after the lower orbitals are
filled. (Aufbau again.)
For the atoms beyond hydrogen, orbital energies
vary as “s<p<d<f” for a given value of n.
Each orbital in a sublevel is occupied by a single
electron before a second electron enters. For
example, all three p orbitals must contain one
electron before a second electron enters a p
orbital. (Hund’s Rule)
7s
6s
Energy
5s
4s
6p
5p
4f
4d
3d
3p
2p
1s
5f
4p
3s
2s
6
d
5d
After 3p is filled, the next
lowest energy is 4s, not 3d.
See all the overlap beyond 3p.
Order of Subshell Filling
in Ground State Electron Configurations
1s
Start by drawing a diagram
putting each energy shell on
a row and listing the subshells,
(s, p, d, f), for that shell in
order of energy, (left-to-right)
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
Next, draw arrows through
the diagonals, looping back
to the next diagonal
each time
6s
6p
6d
7s
Writing Electron Configurations
Number of
electrons in
sublevel orbitals
Arrangement of
electrons within their
respective sublevels.
Principal
Type of orbital
energy level
6
2p
Order of Subshell Filling
in Ground State Electron Configurations
Try using the drawing with the arrows
to do aluminum, phosphorus and
chlorine.
Al 1s22s22p63s23p1
P 1s22s22p63s23p3
Cl 1s22s22p63s23p5
Orbital Box Notation
In the following diagrams boxes
represent orbitals.
Electrons are indicated by arrows: ↑ or
↓ for spin.
Electron configurations can be spdf or
orbital box. See both as follows.
Filling the 1s sublevel
H
↑
1s1
Hydrogen has 1 electron. It will occupy the orbital of
lowest energy which is the 1s.
He
↑↓
1s2
Helium has two electrons. Both helium electrons occupy the
1s orbital with opposite spins.
Filling the 2s sublevel
Li
↑↓
↑
1s
2s
1s22s1
The 1s orbital is filled. Lithium’s third electron will
enter the 2s orbital.
Be
↑↓
↑↓
1s
2s
1s22s2
The 2s orbital fills upon the addition of beryllium’s third and
fourth electrons.
B
↑↓
↑↓
1s
2s
↑
1s22s22p1
2p
Boron has the first p electron. The three 2p orbitals have the
same energy. It does not matter which orbital fills first.
C ↑↓
↑↓
1s
2s
↑
↑
1s22s22p2
2p
The second p electron of carbon enters a different p orbital than the
first p electron so as to give carbon the lowest possible energy.
N ↑↓
↑↓
1s
2s
↑
↑
↑
1s22s22p3
2p
The third p electron of nitrogen enters a different p orbital than its
first two p electrons to give nitrogen the lowest possible energy.
O ↑↓
↑↓
1s
2s
↑↓
↑
↑ 1s22s22p4
2p
There are four electrons in the 2p sublevel of oxygen. One 2p
orbitals is has a second electron, with a spin opposite to the
electron already in the orbital.
F
↑↓
↑↓
↑↓ ↑↓ ↑
1s
2s
2p
1s22s22p5
There are five electrons in the 2p sublevel of fluorine. Two of the 2p
orbitals are now occupied by a second electron, which has a spin
opposite to that of the first electron already in the orbital.
Ne ↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
1s22s22p6
There are 6 electrons in the 2p sublevel of neon, which fills the
sublevel.
Na ↑↓
↑↓
↑↓ ↑↓ ↑↓
↑
1s
2s
2p
3s
1s22s22p63s1
The 2s and 2p sublevels are filled. The next electron enters the
3s sublevel of sodium.
Mg ↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
↑↓ 1s22s22p63s2
3s
The 3s orbital fills upon the addition of magnesium’s twelfth
electron.
After element 18, the 4s sublevel is filled before the 3d is filled.
Writing Electron Configurations
Try writing the electron configurations
for calcium, vanadium, iron, and
arsenic.
Ca 1s22s22p63s23p64s2
V 1s22s2sp63s23p64s23d3
Fe 1s22s2sp63s23p64s23d6
As 1s22s2sp63s23p64s23d104p3
Electron Structures and the Periodic
Table
In 1869, Dimitri Mendeleev of Russia
and Lothar Meyer of Germany
independently published periodic
arrangements of the elements based
on increasing atomic masses.
Mendeleev’s arrangement is the
precursor to the modern periodic
table.
Period numbers correspond
Horizontal rows are
to the highest occupied
called periods.
energy level.
Elements
with
similar
Elements
in
the
B groups
groups
in
A
Groups are
numbered
properties
arenumerals.
organized
are
designated
transition
are
designated
with
Roman
in groups or families.
elements.
representative
elements.
SHORTHAND FOR ELECTRON
CONFIGURATIONS: The electron
configuration of any of the noble gas elements
can be represented by the symbol of the
element enclosed in square brackets.
1s22s22p1
[He]2s22p1
Na
1s22s22p63s1
[Ne]3s1
Cl
1s22s22p63s23p5
[Ne]3s23p5
B
VALENCE ELECTRONS
Valence electrons determine how atoms of the
representative elements will combine with each other.
(By forming ions or by sharing e-s.)
Valence electrons are in highest principal energy level.
FIND valence e-s:
- by looking at where atom is on periodic table
- or by looking at its written e- config and underlining
the e-s in the highest # principal energy level
Notice that Group 1A elements all have ns1 in last place,
Group 2A is ns2, etc.
What is Group 7A “generic” valence arrangement?
ns2np5
VALENCE ELECTRONS
EXAMPLE: IODINE is element #53.
Identify the number of valence e’s
and their location.
It is in group 7A, so it has 7 valence
e-s. Or underline in the
configuration: [Kr]5s24d105p5.
The
chemical
properties
of
For “A”
family behavior
elementsand
the valence
electron
elements
in a family
associated
with the
configuration
is the are
same
in each column.
electron configuration of its elements.
With the exception of helium which has one
filled s orbital, the nobles gases have filled all p
orbitals.
VALENCE ELECTRONS
TRY: Write electron configuration and use
periodic table for arsenic, magnesium,
argon, silicon, gallium. Determine which
are valence electrons and how many
there are.
As
Mg
Ar
Si
[18Ar] 4s23d104p3
[10Ne] 3s2
[10Ne] 3s23p6
[10Ne] 3s23p2
5
2
8
4
d orbital numbers are 1 less
than dthe
period
number
orbital
filling
Arrangement of electrons
according to sublevel being filled.
10.16
f orbital numbers are 2 less
than the
period
number
f orbital
filling
Arrangement of electrons
according to sublevel being filled.
10.16
Period number corresponds with the
highest energy level occupied by
electrons in that period.
10.17
Memorize the Cr and Cu difference!
Electron Configuration Using the
Periodic Table
See page 284 in your textbook for complete
instructions.
Locate the element on the Periodic Table and locate its
row and column. The outermost electrons are
determined by this. Then fill in the “inner” electrons
by looking at the noble gas just before the element
and filling in up to the atom. Make sure the
electrons add up to the element’s atomic number.
Example: Bromine. Row 4, column 7A or 17.
Outermost electrons will end with 4p5. Argon is just
before Br. From Argon, go to row 4, s block, cross
the 3d’s, and you’re at the 4p block.
Bromine is [Ar]4s23d104p5
GROUP PRACTICE:
Try to write Ti, Fe, Pd, and U abbreviated
configuration using the rules & using the
Periodic Table
Ti: [Ar] 4s23d2
Fe: [Ar] 4s23d6
Pd: [Kr] 5s24d8
U: [Rn] 7s25f4
Do the complete e- config and abbrev econfig for germanium, and underline its
valence e-s.
The Noble Gas
Electron Configuration
Noble gases have 8 valence electrons
- except for He, which has only 2
electrons
Noble gases are especially
nonreactive
- He and Ne are practically inert
Noble gases are nonreactive because
their electron configurations are
especially stable
Alkali Metals
Alkali metals have one more electron
than the previous noble gas
Alkali metals tend to lose their extra
electron, resulting in the same
electron configuration as a noble
gas, forming a cation with a 1+
charge
Halogens
Electron configurations for halogens have
one fewer electron than the next noble
gas
In reactions with metals, the halogens tend
to gain an electron and attain the electron
configuration of the next noble gas,
forming an anion with charge 1In reactions with nonmetals, they tend to
share electrons with the other nonmetal,
so that each attains the electron
configuration of a noble gas
Stable Electron Configuration
and Ion Charge
Metals form cations by
losing enough electrons
to get the same electron
configuration as the
previous noble gas
Nonmetals form anions by
gaining enough electrons
to get the same electron
configuration as the next
noble gas
Isoelectronic series: N3- O2F- Ne Na+ Mg2+ Al3+
Na
Atom’s
Electron
Config
[Ne]3s1
Na+
Ion’s
Electron
Config
[Ne]
Mg
[Ne]3s2
Mg2+
[Ne]
Al
[Ne]3s23p1
Al3+
[Ne]
O
[He]2s2p4
O2-
[Ne]
F
[He]2s22p5
F-
[Ne]
Atom
Ion
Make a table as above for the
isoelectronic series around
argon.
Periodic Trends in Atomic Properties
Characteristic properties and trends
of the elements are the basis of the
periodic table’s design.
These trends allow us to use the
periodic table to accurately predict
properties and reactions of a wide
variety of substances.
Atomic radii
increase down a
group.
11.2
For each step down a group electrons enter
the next higher energy level.
Atomic Radius
Radii of atoms tend to decrease
from left to right across a period.
Eachincrease
For
This
time an in
representative
positive
electron
nuclear
is added, a
elements
charge
proton
is
pulls
also
within
all
the same
electrons
added
to the
period,
closer to
the energy
nucleus.
nucleus.level
remains constant
as electrons are
added.
11.2
Example 9.6 – Choose the
Larger Atom in Each Pair
C or O
Li or K
C or Al
Se or I
Example 9.6 – Choose the
Larger Atom in Each Pair
C or O
Li or K
C or Al
Se or I?
IONIZATION ENERGY
The first ionization energy of an atom is the
energy required to remove the first
electron from an atom.
Na + ionization energy → Na+ + eThe second ionization energy is the amount
of energy required to remove the second
electron from an atom.
Na++ second ionization energy → Na2+ + e-
Ionization energies gradually increase from left
to right across a period.
1
2
VIIA
VA
IA
IVA
IIA
11.3
VIA
3
4
IIIA
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
Noble
Gases
Noble
Gases
nonmetals have higher ionization
potentials than metals
VIIA
VA
IA
VIA
IVA
IIA
IIIA
Distance of Outer Shell Electrons From Nucleus
Ionization energies of Group A elements
decrease from top to bottom in a group.
nonmetals
metals
11.3
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
Example 9.7 – Choose the Atom with the
Highest Ionization Energy in Each Pair
Mg or P
As or Sb
N or Si
O or Cl
Example 9.7 – Choose the Atom with the
Highest Ionization Energy in Each Pair
Mg or P
As or Sb
N or Si
O or Cl?