Transcript Chapter 9 Electrons in Atoms and the Periodic Table Helium gas He has 2 protons and if neutral how many electrons? 2006, Prentice Hall.

Chapter 9
Electrons in Atoms and
the Periodic Table
Helium gas
He has 2 protons and
if neutral how many
electrons?
2006, Prentice Hall
CHAPTER OUTLINE


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




Waves
Electromagnetic Radiation
Dual Nature of Light
Bohr Model of Atom
Quantum Mechanical Model of Atom
Electron Configuration
Electron Configuration & Periodic Table
Abbreviated Electron Configuration
Periodic Properties
2
Blimps
• blimps float because they are
filled with a gas that is less dense
than the surrounding air
• early blimps used the gas
hydrogen, however hydrogen’s
flammability lead to the
Hindenburg disaster
• blimps now use helium gas, which
is not flammable
Why is hydrogen gas diatomic but helium gas not?
Because hydrogen is reactive and helium is inert.
What makes hydrogen reactive?
Recall that when elements are arranged in order of
increasing atomic number (# of protons), certain sets
of properties periodically recur
-All group I elements have similar reactivity
-All noble gases are inert
-Alkali metals- +1; metals
-Alkaline earth metals- +2; metals
-Halogens- -1; nonmetals
-Noble gases- 0; nonmetals
We need a theory/model that would explain why these occur:
Classical View of the Universe
• since the time of the ancient Greeks, the stuff of the
physical universe has been classified as either matter
or energy
• we define matter as the stuff of the universe that has
mass and volume
– therefore energy is the stuff of the universe that does not have
mass and volume
• we know that matter is ultimately composed of
particles, and the properties of particles determine the
properties we observe
• energy therefore should not be composed of particles,
in fact the thing that all energy has in common is that it
travels in waves
Light: Electromagnetic Radiation
• light is a form of energy
• light is one type of energy called
electromagnetic radiation with
electric and magnetic field
components and travels in waves
Electromagnetic Waves
1.
velocity = c = speed of light
– its constant! = 2.997925 x 108 m/s (m•sec-1) in vacuum
– all types of light energy travel at the same speed
2. amplitude = A = measure of the intensity of the wave,
“brightness”
– height of the wave
3. wavelength = l = distance between crests
– generally measured in nanometers (1 nm = 10-9 m)
– same distance for troughs or nodes
– determines color
4. frequency = n = how many peaks pass a point in a second
– generally measured in Hertz (Hz),
– 1 Hz = 1 wave/sec = 1 sec-1
Electromagnetic Waves
 Wavelength (λ) is the distance between
any 2 successive crests or troughs.
Amplitude
Nodes
WAVES
 Frequency (nu,n) is the number of waves
produced per unit time.
 Wavelength and frequency are inversely
proportional.
As wavelength of a wave increases
its frequency decreases
inversely
 proportional
Speed tells how fast waves travel through space.
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10.1
ELECTROMAGNETIC
RADIATION
 Energy travels through space as
electromagnetic radiation. This radiation
takes many forms, such as sunlight,
microwaves, radio waves, etc.
 In a vacuum, all electromagnetic waves
travel at the speed of light (3.00 x 108 m/s),
and differ from each other in their
frequency and wavelength.
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ELECTROMAGNETIC
RADIATION
 The
These
classification
waves rangeoffrom
electromagnetic
-rays (short λ,
Long
Short
waves
high
f) according
to radio waves
to wavelength
their
(long
frequency
λ, low f).is
wavelength
calledHigh
electromagnetic Low
spectrum.
frequency
frequency
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ELECTROMAGNETIC
RADIATION
Visible
light is waves
a
Infrared
have
X-rays have longer
small part of the EM
longer λ but lower 
λ but lower  than spectrum
than visible light
-rays
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10.2
Particles of Light
• Albert Einstein and other scientists in the early
20th century showed that wave properties do not
completely explain electromagnetic radiation (EM)
and showed that EM was composed of particlelike properties called photons
– photons are particles of light energy
• each wavelength of light has photons that have a
different amount of energy
– the longer the wavelength, the lower the energy of the
photons
DUAL NATURE
OF LIGHT

have
much evidence
that
light
RedScientists
light has also
longer
wavelength
and less
energy
act as a stream of tiny particles, called
thanbeams
blue light
photons.
A photon of
red light
A photon of
blue light
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DUAL NATURE
OF LIGHT
 Scientists, therefore, use both the wave and
particle models for explaining light. This is
referred to as the wave-particle nature of light.
 Scientists also discovered that when atoms are
energized at high temperatures or by high
voltage, they can radiate light. Neon lights are
an example of this property of atoms.
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Light’s Relationship to Matter
He
• Atoms can acquire extra energy,
but they must eventually release it
• When atoms emit energy, it
always is released in the form of
light
• However, atoms don’t emit all
colors, only very specific
wavelengths
– in fact, the spectrum of wavelengths
can be used to identify the element
Hg
When an atom absorbs energy it reemits it as light
White Light Source
Emits at every wavelength
(all colors) also called a
continuous spectrum
emission spectrum of Hydrogen
H produces its own unique and
distinctive emission spectrum
Spectra
The Bohr Model of the Atom
• The Nuclear Model of the atom does not
explain how the atom can gain or lose
energy
• Neils Bohr developed a model of the
atom to explain the how the structure of
the atom changes when it undergoes
energy transitions
• Bohr’s major idea was that the energy of
the atom was quantized, and that the
amount of energy in the atom was
related to the electron’s position in the
atom
– quantized means that the atom could only
have very specific amounts of energy
1885 to1962
Nobel Prize in
Physics in 1922
The Bohr Model of the Atom
Electron Orbits
• the Bohr Model, electrons travel in orbits
around the nucleus
– more like shells than planet orbits
• the farther the electron is from the nucleus
the more energy it has
Rank the electrons from
highest to lowest energy
eeeee-
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The Bohr Model of the Atom
Orbits and Energy
• each orbit has a specific
amount of energy
• the energy of each orbit is
characterized by an integer the larger the integer, the
more energy an electron in
that orbit has and the farther
it is from the nucleus
– the integer (whole #s), n, is
called a quantum number
Bohr orbits are like steps in a ladder. It is possible to be on one
step or another, but it is impossible to be between steps.
BOHR MODEL
OF ATOM
 Bohr’s
In thisBohr,
Neils
model,
model
a Danish
of
thethe physicist, studied the
hydrogen
atom
electrons
consisted
could
atom of
only
extensively, and developed a
model for
electrons
occupy
particular
orbiting
the atomthe
energy
that was able to explain
the lineand
nucleus
levels,
spectrum.
at different
could “jump”
distances
to
higher levels
from the
by
nucleus, called
energy
absorbing
energy.
levels.
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BOHR MODEL
OF ATOM
 The lowest energy level is
called ground state, and
the higher energy levels
are called excited states.
 When electrons absorb
energy through heating
or electricity, they move
to higher energy levels
and become excited.
energy
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BOHR MODEL
OF ATOM
 When excited electrons
return to Lower
the ground
state,
energy
energy is emitted
as a
transition
photon of light
is released.
give off
red
 The color (wavelength)
of
light
the lightHigher
emittedenergy
is
determined
by thegive
transition
difference
energy
offinblue
light
between the two states
(excited and ground).
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The Bohr Model of the Atom
Ground and Excited States
• The lowest amount of energy hydrogen’s one
electron can have corresponds to being in the
n = 1 orbit –this is the ground state
• when the atom gains energy, the electron
leaps to a higher energy orbit –this is the
excited state
• the atom is less stable in an excited state,
and so it will release the extra energy to
return to the ground state
– either all at once or in several steps
The Bohr Model of the Atom
Hydrogen Spectrum
• every hydrogen atom has identical orbits, so
every hydrogen atom can undergo the same
energy transitions
• however, since the distances between the
orbits in an atom are not all the same, no two
leaps in an atom will have the same energy
– the closer the orbits are in energy, the lower the
energy of the photon emitted
– lower energy photon = longer wavelength
• therefore we get an emission spectrum that
has a lot of lines that are unique to hydrogen
The Bohr Model of the Atom
Hydrogen Spectrum
Which e- has longer wavelength and lower
energy (red, violet or blue-green)?
The Bohr Model of the Atom
Success and Failure
• the mathematics of the Bohr Model very
accurately predicts the spectrum of
hydrogen
• however its mathematics fails when
applied to multi-electron atoms
– it cannot account for electron-electron
interactions
• a better theory was needed
QUANTUM MECHANICAL
MODEL OF ATOM
 In 1926 Erwin Shrödinger created a mathematical
model that showed electrons as both particles and
waves. This model was called the quantum
1887 to 1961
Nobel Prize in
mechanical model.
Physics in 1933
 This model predicted electrons to be located in a
probability region called orbitals.
 An orbital is defined as a region around the
nucleus where there is a high probability of
finding an electron.
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Orbits vs. Orbitals
Pathways vs. Probability
Orbit – acts as a particle and
follows a well-defined path
Orbital – acts as a wave and particle
thus could end up anywhere in a
probability map
QUANTUM MECHANICAL
MODEL OF ATOM
 Based on this model, there
are discrete principal
energy levels within the
atom.
 Principal
levelsthe
are
As energy
n increases,
designatedenergy
by n. of the
 The electrons
inincreases
an atom
electron
can exist in any principal
energy level.
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QUANTUM MECHANICAL
MODEL OF ATOM
 Each principal energy level
is subdivided into sublevels.
 The sublevels are
designated by the letters
s, p, d and f.
 As n increases, the number
of sublevels increases.
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10.7, 10.8
QUANTUM MECHANICAL
MODEL OF ATOM
 Within
The number
the sublevels,
of orbitals
thewithin
electrons
the sublevels
are located in
orbitals.
vary
withThe
theirorbitals
type. are also designated by the
letters s, p, d and f.
s sublevel = 1 orbital = 2 electrons
p sublevel = 3 orbitals = 6 electrons
d sublevel = 5 orbitals = 10 electrons
f sublevel = 7 orbitals = 14 electrons
An orbital can hold a maximum of 2 electrons
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How does the 1s Subshell
Differ from the 2s Subshell
Probability Maps & Orbital Shape
s Orbitals
Probability Maps & Orbital Shape
p Orbitals
Probability Maps & Orbital Shape
d Orbitals
ELECTRON
CONFIGURATION
 The distribution of electrons into the various energy
shells and subshells in an atom’s ground state is
called its electron configuration
 The electrons occupy the orbitals from the lowest
energy level to the highest level (Aufbau Principal).
 The energy of the orbitals on any level are in the
following order: s < p < d < f.
 Each orbital on a sublevel must be occupied by a
single electron before a second electron enters
(Hund’s Rule).
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ELECTRON
CONFIGURATION
 Electron configurations can be written as:
2
Principal
energy
level
6
p
Number of
electrons in
orbitals
Type of
orbital
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ELECTRON
CONFIGURATION
 Another notation, called the orbital notation is
shown below:
Electrons in
orbital with
opposing spins
Principal
energy level
Type of
orbital
1s
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Filling an Orbital with Electrons
• each orbital may have a maximum of 2
electrons with opposite spins
– Pauli Exclusion Principle
• electrons spin on an axis
– generating their own magnetic field
• when two electrons are in the same orbital,
they must have opposite spins
– so there magnetic fields will cancel
ELECTRON
CONFIGURATION
H
↑
1s1
1s
Hydrogen has 1 electron. It will occupy the orbital of
lowest energy which is the 1s.
He
↑↓
1s2
1s
Helium has two electrons. Both helium electrons occupy the
1s orbital with opposite spins.
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ELECTRON
CONFIGURATION
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 due to Hund’s Rule.
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ELECTRON
CONFIGURATION
Ne
↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
1s22s22p6
The last p electron for neon pairs up with the last lone
electron and completely fills the 2nd energy level.
Na ↑
1s22s22p6 3s1
3s
Sodium has 11 electron. The first 10 will occupy the
orbitals of energy levels 1 and 2.
core
electrons
valence
electron
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ELECTRON
CONFIGURATION
 As electrons occupy the 3rd energy level and
higher, some anomalies occur in the order of the
energy of the orbitals.
 Knowledge of these anomalies is important in
order to determine the correct electron
configuration for the atoms.
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ELECTRON CONFIG.
& PERIODIC TABLE
46
10.15
ELECTRON CONFIG.
& PERIODIC TABLE
 The horizontal rows in the periodic table are
called periods. The period number corresponds
to the number of energy levels that are occupied
in that atom.
 The vertical columns in the periodic table are
called groups or families. For the main-group
elements, the group number corresponds to the
number of electrons in the outermost filled
energy level (valence electrons).
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ELECTRON CONFIG.
& PERIODIC TABLE
energy
4One
energy
3 energy
level levels
levels
48
ELECTRON CONFIG.
& PERIODIC TABLE
3 valence
1 valence
5 valence
electrons
electron
electrons
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ELECTRON CONFIG.
& PERIODIC TABLE
 The
Notevalence
that elements
electrons
in the
configuration
same groupfor
have
the
elementselectron
similar
in periods
configurations.
1-3 are shown below.
50
10.15
ELECTRON CONFIG.
& PERIODIC TABLE
Arrangement of orbitals in the periodic table
51
10.16
ELECTRON CONFIG.
& PERIODIC TABLE
d orbital numbers are 1 less
than the period number
52
10.16
ELECTRON CONFIG.
& PERIODIC TABLE
f orbital numbers are 2 less
than the period number
53
10.16
ABBREVIATED
ELECTRON CONFIG.
 When writing electron configurations for larger
atoms, an abbreviated configuration is used.
 In writing this configuration, the non-valence
(core) electrons are summarized by writing the
symbol of the noble gas prior to the element in
brackets followed by configuration of the
valence electrons.
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ABBREVIATED
ELECTRON CONFIG.
K
Z = 19
1s22s22p63s23p6 4s1
core[Ar]
electrons
4s1
Previous
noble gas
valence
electron
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ABBREVIATED
ELECTRON CONFIG.
Br
Z = 35
1s22s22p63s23p6 4s2 3d10 4p5
core[Ar]
electrons
4s23d104p5
valence
electrons
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Electron Configuration of As
from the Periodic Table
8A
1A
1
2
3
4
5
6
7
3A 4A 5A 6A 7A
2A
3d10
4s2
Ar
As
4p3
As = [Ar]4s23d104p3
As has 5 valence electrons
TRENDS IN
PERIODIC PROPERTIES
 The electron configuration of atoms are an
important factor in the physical and chemical
properties of the elements.
 Some of these properties include: atomic size,
ionization energy and metallic character.
 These properties are commonly known as
periodic properties and increase or decrease
across a period or group, and are repeated in
each successive period or group.
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ATOMIC SIZE
 The size of the atom is determined by its atomic
radius, which is the distance of the valence
electron from the nucleus.
 For each group of the representative elements,
the atomic size increases going down the group,
because the valence electrons from each energy
level are further from the nucleus.
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ATOMIC SIZE
60
Group IIA
2e2e-
Be
(4p+
&
4e-)
4
Be
4 p+
2e-
Mg (12p+ & 12e-)
8e2e-
12
Mg
12 p+
2e8e-
Ca
(20p+
&
20e-)
8e-
20
Ca
2e20 p+
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ATOMIC SIZE
 The atomic radius of the representative elements
are affected by the number of protons in the
nucleus (nuclear charge).
 For elements going across a period, the atomic
size decreases because the increased nuclear
charge of each atom pulls the electrons closer to
the nucleus, making it smaller.
62
ATOMIC SIZE
63
Period 2
1e2e3 p+
Li (3p+ & 3e-)
4e2e-
2e2e4 p+
Be (4p+ & 4e-)
6e2e-
6 p+
8 p+
C (6p+ & 6e-)
O (8p+ & 8e-)
3e2e5 p+
B (5p+ & 5e-)
8e2e10 p+
Ne (10p+ & 10e-)
IONIZATION
ENERGY
 The ionization energy is the energy required to
remove a valence electron from the atom in a
gaseous state.
 When an electron is removed from an atom, a
cation (+ ion) with a 1+ charge is formed.
Na (g) + IE
Na+ + e-
65
IONIZATION
ENERGY
 The ionization energy
decreases going down a
group, because less energy
is required toLarger
removeatom
an
Less
IE
electron from the
outer
shell since it is further
from the nucleus.
66
IONIZATION
ENERGY
 Going across a period,
the ionization energy
increases because the
increased nuclear
charge of the atom
holds the valence
electrons more tightly
and therefore it is more
difficult to remove.
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IONIZATION
ENERGY
 In general, the ionization
energy is low for metals
and high for non-metals.
 Review of ionization
energies of elements in
periods 2-4 indicate
some anomalies to the
general increasing trend.
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IONIZATION
ENERGY
 These anomalies are caused by more stable
More
stable
electron configurations
ofstable
the atoms in groups 2
More
(1/2and
filled)
(complete “s” sublevel)
group 5 (half-filled “p”
Higher
IE
Higher
IE in their ionization
sublevels) that cause an
increase
energy compared to the next element.
Be
1s2 2s2
N
1s2 2s2 2p3
B
1s2 2s2 2p1
O
1s2 2s2 2p4
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METALLIC
CHARACTER
 Metallic character is the ability of an atom to lose
electrons easily.
 This character is more prevalent in the elements
on the left side of the periodic table (metals), and
decreases going across a period and increases for
elements going down a group.
70
METALLIC
CHARACTER
Most
metallic
elements
Least
metallic
elements
71
Example 1:
Select the element in each pair with the larger
atomic radius:
K
or
Br
Larger due to
less nuclear
charge
72
Example 2:
Indicate the element in each set that has the higher
ionization energy and explain your choice:
F
N
or
C
Highest IE due
to most
nuclear charge
73
THE END
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