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Electrons in Atoms and

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Title: Electrons in Atoms and


1
Chapter 9 Electrons in Atoms and the Periodic
Table
Helium gas
He has 2 protons and if neutral how
many electrons?
2006, Prentice Hall
2
CHAPTER OUTLINE
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
3
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
    hydrogens flammability lead to the Hindenburg
    disaster
  • blimps now use helium gas, which is not flammable

4
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
5
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

6
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

7
Electromagnetic Waves
  • velocity c speed of light
  • its constant! 2.997925 x 108 m/s (msec-1) in
    vacuum
  • all types of light energy travel at the same
    speed
  • amplitude A measure of the intensity of the
    wave, brightness
  • height of the wave
  • wavelength l distance between crests
  • generally measured in nanometers (1 nm 10-9 m)
  • same distance for troughs or nodes
  • determines color
  • frequency n how many peaks pass a point in a
    second
  • generally measured in Hertz (Hz),
  • 1 Hz 1 wave/sec 1 sec-1

8
Electromagnetic Waves
  • Wavelength (?) is the distance between any 2
    successive crests or troughs.

Amplitude
Nodes
9
WAVES
  • Frequency (nu,?) is the number of waves produced
    per unit time.
  • Wavelength and frequency are inversely
    proportional.
  • Speed tells how fast waves travel through space.

As wavelength of a wave increases
its frequency decreases
10.1
10
ELECTROMAGNETICRADIATION
  • 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.

11
ELECTROMAGNETICRADIATION
  • The classification of electromagnetic waves
    according to their frequency is called
    electromagnetic spectrum.
  • These waves range from ?-rays (short ?, high f)
    to radio waves (long ?, low f).

Long wavelength Low frequency
Short wavelength High frequency
12
ELECTROMAGNETICRADIATION
Visible light is a small part of the EM spectrum
Infrared waves have longer ? but lower ? than
visible light
X-rays have longer ? but lower ? than ?-rays
10.2
13
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 particle-like
    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

14
DUAL NATUREOF LIGHT
  • Scientists also have much evidence that light
    beams act as a stream of tiny particles, called
    photons.

Red light has longer wavelength and less energy
than blue light
A photon of red light
A photon of blue light
15
DUAL NATUREOF 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.

16
Lights 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 dont emit all colors, only very
    specific wavelengths
  • in fact, the spectrum of wavelengths can be used
    to identify the element

Hg
17
When an atom absorbs energy it reemits it as light
emission spectrum of Hydrogen
White Light Source
Emits at every wavelength (all colors) also
called a continuous spectrum
H produces its own unique and distinctive
emission spectrum
18
Spectra
19
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
  • Bohrs major idea was that the energy of the atom
    was quantized, and that the amount of energy in
    the atom was related to the electrons 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
20
The Bohr Model of the AtomElectron 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
e-
e-
e-
e-
e-
21
The Bohr Model of the AtomOrbits 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.
22
BOHR MODELOF ATOM
  • Neils Bohr, a Danish physicist, studied the
    hydrogen atom extensively, and developed a model
    for the atom that was able to explain the line
    spectrum.
  • Bohrs model of the atom consisted of electrons
    orbiting the nucleus at different distances from
    the nucleus, called energy levels.
  • In this model, the electrons could only occupy
    particular energy levels, and could jump to
    higher levels by absorbing energy.

23
BOHR MODELOF 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
24
BOHR MODELOF ATOM
  • When excited electrons return to the ground
    state, energy is emitted as a photon of light is
    released.
  • The color (wavelength) of the light emitted is
    determined by the difference in energy between
    the two states (excited and ground).

Lower energy transition give off red light
Higher energy transition give off blue light
25
The Bohr Model of the AtomGround and Excited
States
  • The lowest amount of energy hydrogens 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

26
The Bohr Model of the AtomHydrogen 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

27
The Bohr Model of the AtomHydrogen Spectrum
Which e- has longer wavelength and lower energy
(red, violet or blue-green)?
28
The Bohr Model of the AtomSuccess 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

29
QUANTUM MECHANICALMODEL OF ATOM
  • In 1926 Erwin Shrödinger created a mathematical
    model that showed electrons as both particles and
    waves. This model was called the quantum
    mechanical model.
  • 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.

1887 to 1961 Nobel Prize in Physics in 1933
30
Orbits vs. OrbitalsPathways vs. Probability
Orbital acts as a wave and particle thus could
end up anywhere in a probability map
Orbit acts as a particle and follows a
well-defined path
31
QUANTUM MECHANICALMODEL OF ATOM
  • Based on this model, there are discrete principal
    energy levels within the atom.
  • Principal energy levels are designated by n.
  • The electrons in an atom can exist in any
    principal energy level.

As n increases, the energy of the electron
increases
32
QUANTUM MECHANICALMODEL OF ATOM
  • Each principal energy level is subdivided into
    sublevels.
  • The sublevels are designated by the letterss, p,
    d and f.
  • As n increases, the number of sublevels
    increases.

10.7, 10.8
33
QUANTUM MECHANICALMODEL OF ATOM
  • Within the sublevels, the electrons are located
    in orbitals. The orbitals are also designated by
    the letters s, p, d and f.
  • The number of orbitals within the sublevels vary
    with their type.

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
34
How does the 1s Subshell Differ from the 2s
Subshell
35
Probability Maps Orbital Shapes Orbitals
36
Probability Maps Orbital Shapep Orbitals
37
Probability Maps Orbital Shaped Orbitals
38
ELECTRONCONFIGURATION
  • The distribution of electrons into the various
    energy shells and subshells in an atoms 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 lt p lt d lt f.
  • Each orbital on a sublevel must be occupied by a
    single electron before a second electron enters
    (Hunds Rule).

39
ELECTRONCONFIGURATION
  • Electron configurations can be written as

Number of electrons in orbitals
2 p6
Principal energy level
Type of orbital
40
ELECTRONCONFIGURATION
  • Another notation, called the orbital notation is
    shown below

Electrons in orbital with opposing spins
Principal energy level
1 s
Type of orbital
41
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

42
ELECTRONCONFIGURATION
?
H
1s1
Hydrogen has 1 electron. It will occupy the
orbital of lowest energy which is the 1s.
?
?
1s2
He
Helium has two electrons. Both helium electrons
occupy the 1s orbital with opposite spins.
43
ELECTRONCONFIGURATION
?
B
1s22s22p1
Boron has the first p electron. The three 2p
orbitals have the same energy. It does not
matter which orbital fills first.
?
?
?
C
1s22s22p2
The second p electron of carbon enters a
different p orbital than the first p due to
Hunds Rule.
44
ELECTRONCONFIGURATION
?
?
?
?
?
?
Ne
1s22s22p6
The last p electron for neon pairs up with the
last lone electron and completely fills the 2nd
energy level.
?
Na
1s22s22p6
3s1
Sodium has 11 electron. The first 10 will occupy
the orbitals of energy levels 1 and 2.
core electrons
valence electron
45
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.

46
ELECTRON CONFIG. PERIODIC TABLE
10.15
47
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).

48
ELECTRON CONFIG. PERIODIC TABLE
One energy level
4 energy levels
3 energy levels
49
ELECTRON CONFIG. PERIODIC TABLE
3 valence electrons
1 valence electron
5 valence electrons
50
ELECTRON CONFIG. PERIODIC TABLE
  • The valence electrons configuration for the
    elements in periods 1-3 are shown below.
  • Note that elements in the same group have similar
    electron configurations.

10.15
51
ELECTRON CONFIG. PERIODIC TABLE
Arrangement of orbitals in the periodic table
10.16
52
ELECTRON CONFIG. PERIODIC TABLE
d orbital numbers are 1 less than the period
number
10.16
53
ELECTRON CONFIG. PERIODIC TABLE
f orbital numbers are 2 less than the period
number
10.16
54
ABBREVIATEDELECTRON 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.

55
ABBREVIATEDELECTRON CONFIG.
Previous noble gas
K
Z 19
1s22s22p63s23p6
4s1
valence electron
core electrons
Ar
4s1
56
ABBREVIATEDELECTRON CONFIG.
Br
Z 35
1s22s22p63s23p6
4s2
3d10
4p5
valence electrons
core electrons
Ar
4s23d104p5
57
Electron Configuration of As from the Periodic
Table
8A
1A
1 2 3 4 5 6 7
3A
4A
5A
6A
7A
2A
Ar
3d10
As
4s2
4p3
As Ar4s23d104p3 As has 5 valence electrons
58
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.

59
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.

60
ATOMIC SIZE
61
Group IIA
4
Be (4p 4e-)
Mg (12p 12e-)
20
Ca (20p 20e-)
62
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.

63
ATOMIC SIZE
64
Period 2
Li (3p 3e-)
Be (4p 4e-)
B (5p 5e-)
C (6p 6e-)
O (8p 8e-)
Ne (10p 10e-)
65
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.

66
IONIZATION ENERGY
  • The ionization energy decreases going down a
    group, because less energy is required to remove
    an electron from the outer shell since it is
    further from the nucleus.

Larger atom Less IE
67
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.

68
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.

69
IONIZATION ENERGY
  • These anomalies are caused by more stable
    electron configurations of the atoms in groups 2
    (complete s sublevel) and group 5 (half-filled
    p sublevels) that cause an increase in their
    ionization energy compared to the next element.

More stable (1/2 filled) Higher IE
More stable Higher IE
Be
1s2 2s2
N
1s2 2s2 2p3
B
1s2 2s2 2p1
O
1s2 2s2 2p4
70
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.

71
METALLIC CHARACTER
Most metallic elements
Least metallic elements
72
Example 1
Select the element in each pair with the larger
atomic radius
or
K
Br
Larger due to less nuclear charge
73
Example 2
Indicate the element in each set that has the
higher ionization energy and explain your choice
or
F
C
N
Highest IE due to most nuclear charge
74
  • THE END
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