1 modern atomic theory and the periodic table chapter 10
TRANSCRIPT
1
Modern Atomic Theoryand the Periodic Table Chapter 10
Modern Atomic Theoryand the Periodic Table Chapter 10
2
Chapter 10 - Modern Atomic Theory and the Periodic Table
10.1 A Brief History
10.2 Electromagnetic Radiation
10.3 The Bohr Atom
10.5 Atomic Structures of the First 18 Elements
10.6 Electron Structures and the Periodic Table
10.4 Energy Levels of Electrons
3
A B r ie f H is to ry o f A to m ic T h eo ry
W hile th ese m od e ls w o rk re a so na b ly w e llth e ir lim ita ton s h av e le d to m o re m o de rn the or ies
a s to th e n a tu re o f the a to m .
L im ita tio ns o f D a lton 's m od e lle d to th e T h om pson an d R u th e rfo rd
m o de ls o f the a to m .
T h e ir e xpe rim en ts led toD a lto n 's A to m ic T h eo ry
E a rly ch e m is tsp e rfo rm ed expe r im en ts
G re eks w e re th e f i r st to su gg e stth a t m a tte r is m a d e u p o f a to m s
A Brief HistoryA Brief History
4
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.
3.00 x 108 m/s
Electromagnetic RadiationElectromagnetic Radiation
Examples
5
Wavelength (λ)
Characteristics of a WaveCharacteristics of a WaveCharacteristics of a WaveCharacteristics of a Wave
wavelength(measured from peak to peak)
wavelength(measured from trough to trough)
10.1
Light has the properties of a wave.
6
Frequency () is the number of wavelengths that pass a particular point per second.
10.1
7
Speed (v) is how fast a wave moves through space.
10.1
8
• Light also exhibits the properties of a particle. Light particles are called photons.
• Both the wave model and the particle model are used to explain the properties of light.
9
10.2
visible light is part of the electromagnetic
spectrum
X-rays are part of the electromagnetic
spectrum
Infrared light is part of the
electromagnetic spectrum
The Electromagnetic Spectrum
10
• 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.
The Bohr AtomThe Bohr Atom
11
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.
These colored lines indicate that light is being emitted only at certain wavelengths.
Each element has its own unique set of spectral emission lines that distinguish it from other elements.
10.3
12
Niels Bohr, a Danish physicist, in 1912-1913 carried out researchon the hydrogen atom.
Niels BohrNiels BohrNiels BohrNiels Bohr
13
Electrons revolve around the nucleus in orbits that are located at fixed distances from the nucleus.
10.4
An electron has a discrete energy when it occupies an orbit.
The Bohr Atom
14
When an electron falls from a higher energy level to a lower energy level a quantum of energy in the form of light is emitted by the atom.
10.4
The color of the light emitted corresponds to one of the lines of the hydrogen spectrum.
The Bohr Atom
15
Different lines of the hydrogen spectrum correspond to different electron energy level shifts.
10.4
The Bohr Atom
16
Light is not emitted continuously. It is emitted in discrete packets called quanta.
10.4
The Bohr Atom
17
An electron can have one of several possible energies depending on its orbit.
E2 E3E1
10.4
The Bohr Atom
18
Bohr’s calculations succeeded very well in correlating the experimentally observed spectral lines with electron energy levels for the hydrogen atom.
• Bohr’s methods did not succeed for heavier atoms.
• More theoretical work on atomic structure was needed.
The Bohr Atom
19
In 1924 Louis De Broglie suggested that all objects 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.
20
In 1926 Erwin 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.
21
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.
23
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 ElectronsEnergy Levels of Electrons
24
The first four principal energy levels of the hydrogen atom.
As n increases, the energy of the electron increases.
10.7
Each level is assigned a principal quantum number n.
25
Each principal energy level is subdivided into sublevels.
10.7, 10.8
26
Within sublevels the electrons are found in orbitals.
An s orbital is spherical in shape.
The spherical surface encloses a space where there is a 90% probability that the electron may be found.
10.10
27
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.
10.10
An atomic orbital can hold a maximum of two electrons.
28
Each p orbital has two lobes.
Each p orbital can hold a maximum of two electrons.
A p sublevel can hold a maximum of 6 electrons.
10.10
A p sublevel is made up of three orbitals.
29
The three p orbitals share a common center.
10.10
pxpy
pz
The three p orbitals point in different directions.
The P orbitals
30
The five d orbitals all point in different directions.
Each d orbital can hold a maximum of two electrons.
A d sublevel can hold a maximum of 10 electrons.
10.11
A d sublevel is made up of five orbitals.
32
n=1 1s
n=2 2s 2p 2p 2p
n = 3 3s 3p 3p 3p 3d 3d 3d 3d 3d
n = 4 4s 4p 4p 4p 4d 4d 4d 4d 4d 4f 4f 4f 4f 4f 4f 4f
Distribution of Subshells by Principal Energy Level
33
The Hydrogen Atom
In the ground state hydrogen’s single electron lies in the 1s orbital.
Hydrogen can absorb energy and the electron will move to excited states.
10.12
• The diameter of hydrogen’s nucleus is about 10-13 cm.
• The diameter of hydrogen’s electron cloud is about 10-8 cm.
• The diameter of hydrogen’s electron cloud is about 100,000 times greater than the diameter of its nucleus.
34
To determine the electronic structures of atoms, the following guidelines are used.
Atomic Structure of the Atomic Structure of the First 18 ElementsFirst 18 Elements
35
1. No more than two electrons can occupy one orbital
10.10
Pauli exclusion principle
36
1 s orbital
2. Electrons occupy the lowest energy orbitals available. They enter a higher energy orbital only after the lower orbitals are filled.
3. For the atoms beyond hydrogen, orbital energies vary as s<p<d<f for a given value of n.
2 s orbital
10.10
37
4. 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.
10.10
38
Nuclear makeup and electronic structure of each principal energy level of an atom.
number of electronsin each sublevel
number of protons and neutrons in the nucleus
10.13
39
Electron Configuration
Arrangement of electrons within their respective sublevels. 2p6Principal
energy levelType of orbital
Number of electrons in
sublevel orbitals
40
In the following diagrams boxes represent orbitals.
• Electrons are indicated by arrows: ↑ or ↓.– Each arrow direction represents one of
the two possible electron spin states.
Orbital FillingOrbital FillingOrbital FillingOrbital Filling
41
↑ 1s2↓
H ↑ 1s1
Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s.
He
Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins.
Filling the 1s Sublevel
42
Li
1s22s2
The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital.
↑↓ ↑
↓
1s 2s
↑
1s22s1
Be ↑↓
The 2s orbital fills upon the addition of beryllium’s third and fourth electrons.
1s 2s
Filling the 2s Sublevel
43
B 1s22s22p1
1s 2s 2p
↑↓ ↑↓ ↑
Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first.
C1s 2s 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.
1s22s22p2↑ ↑
Filling the 2p Sublevel
44
↑ 1s22s22p4↑ ↑
1s 2s 2p
↑↓ ↑↓O
There are four electrons in the 2p sublevel of oxygen. One of the 2p orbitals is now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital.
↓
↑ ↑ ↑N1s 2s 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.
1s22s22p3
45
↑↓ ↑↓ ↑↓
2p
Ne1s 2s
↑↓ ↑↓
There are 6 electrons in the 2p sublevel of neon, which fills the sublevel.
1s22s22p6
↑↓ ↑↓ ↑
2p
F1s 2s
↑↓ ↑↓
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.
1s22s22p5
46
Na 1s22s22p63s1
1s 2s 2p 3s
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium.
Mg1s 2s 2p 3s
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
The 3s orbital fills upon the addition of magnesium’s twelfth electron.
1s22s22p63s2↓
Filling the 3s Sublevel
47
48
49
Electron Filling Order
Sublevel energy level order:
1s < 2s < 2p < 3s < 3p <
4s < 3d < 4p < 5s < 4d <
5p < 6s < 4f < 5d < 6p <
7s < 5f < 6dYou can memorize this sequence or....
52
In 1869 Dimitri Mendeleev of Russia and Lothar Meyer of Germany independently published periodic arrangements of theelements based on increasing atomic masses.
Mendeleev’s arrangement is the precursor to the modern periodic table.
Electron Structures and the Electron Structures and the Periodic TablePeriodic Table
53
Horizontal rows are called periods.
Period numbers correspond to the highest occupied energy level.
10.14
54
Elements in the A groups are designated
representative elements.
10.14
Elements in the B groups are designated transition
elements.
Elements with similar properties are organized in groups or families.
Groups are numbered with Roman numerals.
57
10.15
For A family elements the valence electron configuration is the same in each column.
The chemical behavior and properties of elements in a family are associated with the electron configuration of its elements.
58
10.15
With the exception of helium which has a filled s orbital, the nobles gases have filled p orbitals.
To write an electron configurationusing a noble-gas core:
2. Write the elemental symbol of the noble gas in square brackets, followed
by the remaining configuration
1. Find the highest atomic-numberednoble gas (Group 8A element)less than the atomic numberof the element for which theconfiguration is being written
60
The electron configuration of any of the noble gas elements can be represented by the symbol of the element enclosed in square brackets.
B 1s22s22p1 [He]2s22p1
Cl 1s22s22p63s23p5 [Ne]3s23p5
Na 1s22s22p63s1 [Ne]3s1
61
The electron configuration of argon is
Ar 1s22s22p63s23p6
Ca 1s22s22p63s23p6 4s2 [Ar]4s2
K 1s22s22p63s23p64s1 [Ar]4s1
The elements after argon are potassium and calcium Instead of entering a 3d orbital, the valence electrons of these elements enter the 4s orbital.
Exceptions to the conventional filling order:
1. d4 configurations generally do not exist
Chromium (Z = 24):Systematic prediction:
Cr: [Ar]4s23d4
But d4 is not likely,so promote an electron from the 4s sublevel:
Cr: [Ar]4s13d5
2. d9 configurations generally do not exist
Copper (Z = 29):Systematic prediction:
Cu: [Ar]4s23d9
But d9 is not likely,so promote an electron from the 4s sublevel:
Cu: [Ar]4s13d10
64
10.16
d orbital filling d orbital numbers are 1 less
than the period number
Arrangement of electronsaccording to sublevel being filled.
65
10.16
f orbital filling f orbital numbers are 2 less
than the period number
Arrangement of electronsaccording to sublevel being filled.
66
10.17
Period number corresponds with the highest energy level occupied by
electrons in that period.
67
10.17
The group numbers for the representative elements are equal to the total number of
outermost electrons in the atoms of the group.
The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels.
68
Chapter 10 - Modern Atomic Theory and the Periodic Table 10.1 A Brief History
10.2 Electromagnetic Radiation
10.3 The Bohr Atom – Niels Bohr description of the atom (electron orbitals).
10.4 Energy Levels of Electrons – Electron configuration (from the periodic table), s, p, d, and f orbitals.
10.5 Atomic structures of the First 18 Elements – Valence electrons, Representatives and Transition elements, Families names.
10.6 Electron Structures and the Periodic table – Relationship between group number and valence electrons.