3'4 Where Are the Electrons in Atoms

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3.4 Where Are the Electrons in Atoms
Important Topics

• Continuous and Line Spectra
• Bohr Model of the Atom
• Quantum Mechanical Model of the Atom
• Atoms Building - Electron Energy Levels

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The Bohr Model of the Hydrogen Atom

• In 1911, Rutherfords ? - particle experiments
  were very controversial
• The idea that all the positive charge of an
  atom
• was crammed into the nucleus was hard for
• many to accept

• In 1913, Neils Bohr, a Danish physicist proposed
• That indeed all the positive charge was in the
  nucleus
• The electrons orbited the nucleus much like
• planets orbit the sun

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The Bohr Model of the Hydrogen Atom (cont)
Bohr based his model of several well known facts
at the time

• Visible light, x-rays, ultraviolet radiation,
  infrared
• radiation, microwaves and radio waves are
  all part
• of the electromagnetic spectrum

Gamma Rays X-rays
Ultraviolet

Radio and Infrared Microwaves
Television
waves
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The Bohr Model of the Hydrogen Atom (cont)

• Waves can be described by the wave equation which
  includes
• velocity (c speed of light), wavelength
  (?) and frequency (?)

Wavelength The distance between peaks of a
wave
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The Bohr Model of the Hydrogen Atom (cont)

• Light from the sun (white light) appears as a
  continuous
• spectrum of light.
• Continuous There are no discrete,
  individual
• Spectrum of Light wavelengths of light but
  rather all
• wavelengths appear, one after
• the other in
  a continuous fashion
• When an element (hydrogen) is placed in a
  container
• and heated or excited electrically the light
  emitted
• forms a line spectrum of light
• Line Spectrum of Light Only certain
  wavelengths appear
• 
  and not others

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The Bohr Model of the Hydrogen Atom (cont)
Bohr boldly proposed that the energy possessed by
an electron in a hydrogen atom and the radius of
the orbit are quantized Quantized Values
limited to only specific values and not a
continuous range of values
The ramp is an example of a continuous situation
in which any energy state is possible up the
ramp
Like a set of stairs, the energy states of an
electron is quantized
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The Bohr Model of the Hydrogen Atom (cont)

• The Energy Absorption Process
• Light or energy excites an electron from a
  lower energy level
• (orbits in the Bohr theory) to a higher energy
  level
• Since these energy levels are quantized and
  
• hence the electron can not ever be at any
• intermediate level, this means that the
  electron
• simply disappears from one orbit and
  reappears
• in another
• This absorption or excitation process is called
  
• a quantum leap or quantum jump
• Quantum Jump The quantized process by which
  an electron
• moves up or
  down from one energy level
• to another
• The electron in a hydrogen atom is usually
  found in

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The Bohr Model of the Hydrogen Atom (cont)
Ground State The condition of an atom in
which all electrons are in their normal,
lowest energy levels
What is a good analogy for this process?
A spring and two balls
The electron absorbs energy in the ground
state and is excited to a higher level
Both the atom and the electron now have higher
energy
This is an energy emission process and what we
observe in the hydrogen line spectrum
The Excited State
The Ground State
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The Bohr Model of the Hydrogen Atom (cont)
The Excited State The condition in which at
least one electron in an atom is at an energy
level above the ground state

• An unstable, higher energy state of an atom

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The Bohr Model of the Hydrogen Atom (cont)

• The atomic line spectral lines result when an
  electron in an
• excited state decays back to the ground
  state

The electron loses energy, light is emitted and
the electron returns to the ground state
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The Bohr Model of the Hydrogen Atom (cont)

• The Bohr model works well for the hydrogen atom
  only
• For elements larger than hydrogen the model
• does not work

• Never-the-less, Bohr made two huge contributions
  to the
• development of modern atom theory
• He suggested a reasonable explanation for the
• atomic line spectra in terms of electron
  energies
• He introduced the idea of quantized electron
  energy
• levels in the atom

The Bohr atom lasted for about 13 years and was
quickly replaced by the quantum mechanical model
of the atom The Bohr model is a good starting
point for understanding the quantum mechanical
model of the atom
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The Quantum Mechanical Model of the Atom
In 1924, Louis de Broglie, a French physicist
suggested that all matter has both wave like
properties and particle like properties Accepta
nce of the dual nature of matter and a principle
called the uncertainty principle led to the
development of Quantum Mechanics Uncertainty
Principle It is impossible to know
simultaneously the exact position and
momentum of an
electron (Werner Heisenberg) Between 1925
and 1928 Erwin Schrodinger applied the
principles of wave mechanics to atoms and
developed the quantum mechanical model of the
atom The model describes an atom that has
certain allowed quantities of energy due to the
wavelike motion of an electron whose exact
location is impossible to know
Rather than think of the electron as a particle
orbiting a nucleus, now the electron is treated
as a matter-wave in three dimensional space
around the nucleus
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The Quantum Mechanical Model of the Atom (cont)

• A series of wave functions were developed by
  Schrodinger
• to describe the motion of the electrons
  matter-wave in
• terms of time and position
• Each solution to what is now called the
  Schrodinger
• equation is associated with a given wave
  function, also
• called an atomic orbital
• It is not possible to know precisely where the
  electron
• is at any moment, but it is possible to describe
  where
• it is most probable to find the electron
• An equation called the square of the wave
  function
• describes the possibility of finding the
  electron anywhere
• in three dimensional space

For any given energy level the electron
probability can be expressed by an electron
density diagram that describes an orbital
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Atom Building

• The Principle Energy Level (cont)
• The principle energy level (shell) gets larger
• (their radius increases) as the principle
  quantum
• number n increases

• As the principle level increases in size it can
  hold
• more electrons that the level below it

The number of electrons a shell can hold is
limited by the equation Bohr developed
When an atom is being built the level that is
filled first is the level with the lowest energy
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Atom Building (cont)
The Principle Energy Level
Maximum Electron Capacities of the First Four
Principle Energy Levels (Shells)
n 4 2n2 2 x 42 32 electrons n 3
2n2 2 x 32 18 electrons n 2
2n2 2 x 22 8 electrons n 1 2n2
2 x 11 2 electrons
58 electrons
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The Seventh Level is the Highest
Occupied Ground-State Electrons in any Element
now Known
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Atom Building (cont)
Hydrogen (H) 1 1e Helium (He)
2 2e Lithium (Li) 3 2e 1e Beryllium (Be)
4 2e 2e Boron (B) 5 2e 3e Carbon
(C) 6 2e 4e Nitrogen (N)
7 2e 5e Oxygen (O) 8 2e 6e Fluorine (F)
9 2e 7e Neon (Ne) 10 2e 8e
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Valence Electrons Outermost electrons in an
atom
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Atom Building (cont)
Electron Arrangements of the First 20 Elements
(cont)
Number of Electrons in Each Energy Level
Atomic Number
1st 2nd 3rd 4th
Element
Sodium (Na) 11 2e 8e 1e Magnesium (Mg)
12 2e 8e 2e Aluminum (Al)
13 2e 8e 3e Silicon (Si)
14 2e 8e 4e Phosphorus (P)
15 2e 8e 5e Sulfur (S) 16 2e 8e 6e Chlorin
e (Cl) 17 2e 8e 7e Argon (Ar)
18 2e 8e 8e Potassium (K)
19 2e 8e 8e 1e Calcium (Ca)
20 2e 8e 8e 2e
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Valence Electrons Outermost electrons in an
atom
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3.5 Development of the Periodic Table
Chemists were very busy during the 19th century

• Thousands of chemical compounds were decomposed
  
• into their component elements
• These elements were characterized to determine
  the
• properties of the fundamental chemical
  building blocks
• By 1860, 70 of the now 115 elements were
  identified
• This represented a great simplification for the
  science
• of chemistry
• It became necessary to organize these elements
  in
• some way to make sense of the trends in
  behavior
• that were identified

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Any Object in the Universe Could Now be
Understood in Terms of the Relatively Few
Elements that were at the Base of Its Composition
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Chemical Periodicity - The Law of Mendeleev
(cont)
First Periodic Table developed in 1869 by
Mendeleev and Meyer
Chemical Periodicity or - The repeating
behavior in the Periodic Behavior
chemical properties of the elements
The Law of Mendeleev Properties of the
elements recur in
regular cycles (periodically) when
the elements are arranged in order of

increasing atomic mass

• The periodic table is organized with atomic
  number
• increasing from left-to-right and from
  top-to-bottom

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Chemical Periodicity - The Law of Mendeleev
(cont)
Periodic Behavior of Elements - Repeating
Chemical Properties
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The Periodic Table and the Elements (cont)





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Inner Transition Elements





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3.6 The Modern Periodic Table

• There are two general regions of the periodic
  table

Metals
Nonmetals
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Organization Across the Periodic Table
Periods Arranged horizontally across
the periodic table (rows 1-7)

• These elements have the same valence shell
  number

2nd Period
6th Period
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Organization Down the Periodic Table
Groups (Family) Arranged vertically down the
periodic table (columns or groups,
1- 18 or 1-8 A,B)

• These elements have the same number of
  electrons in
• the outer most shells, the valence
  shell.

Alkali Family 1 e - in the valence shell
Halogen Family 7 e - in the valence shell
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3.7 Periodic Trends - Lewis Dot Symbols

• Another way to show valence electrons is to use
  Lewis Symbols
• Lewis Symbols The element displayed with an
  electron dot
• symbol to show the valence electrons

Group No V. e- Lewis Symbol
8A 8 . . Ne . .
7A 7 . F . .
6A 6 . O .
5A 5 . . N .
4A 4 . . C . .
3A 3 . . B .
1A 1 Li .
2A 2 . Be .
No V. e- Number of Valence Electrons
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3.7 Periodic Trends - Atomic Properties

• Atomic Size
• Moving down a group size increases

• Moving across a period (row) size decreases

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3.8 Properties of Main-Group Elements Noble Gases

• Group 8 Noble Gases
• He, Ne, Ar, Kr, Xe, Rn
• All colorless gases at room temperature
• Very non-reactive, practically inert
• Found in nature as a collection of separate atoms
  
• uncombined with other atoms

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3.8 Properties of Main-Group Elements Noble
Metals

• Noble Metals
• Ag, Au, Pt
• All solids at room temperature
• Least reactive metals
• Found in nature uncombined with other atoms

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3.8 Properties of Main-Group Elements Halogens

• Group 7A Halogens
• Very reactive nonmetals
• React with metals to form ionic compounds
• HX all acids

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3.8 Properties of Main-Group Elements Halogens
(cont)

• Fluorine F2
• Pale yellow gas
• Chlorine Cl2
• Pale green gas
• Bromine Br2
• Brown liquid that has lots of brown vapor over it
• Only other liquid element at room conditions is
  the metal Hg
• Iodine I2
• Lustrous, purple solid

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3.8 Properties of Main-Group Elements Ions

• Ions that have a positive charge are called
  cations
• Formed when an atom loses electrons
• Ions that have a negative charge are called
  anions
• Formed when an atom gains electrons
• Ions with opposite charges attract
• Therefore cations and anions attract each other
• Moving ions conduct electricity
• Compounds must have no total charge,
• Therefore the numbers of cations and anions in a
  compound must balance to get 0 total charge

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3.8 Properties of Main-Group Elements Ions (cont)

• Metals form cations
• For each positive charge the ion has 1 less
  electron than the
• neutral atom
• Na 11 e- Na 10 e-
• Ca 20 e- Ca 2 18 e-
• Cations are named the same as the metal
• Sodium Na ? Na 1e- Sodium ion
• Calcium Ca ? Ca 2 2e- Calcium ion
• The charge on a cation can be determined from the
  Group
• number on the Periodic Table for
  Groups IA, IIA, IIIA
• Group 1A ? 1

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3.8 Properties of Main-Group Elements Ions (cont)

• Nonmetals form anions
• For each negative charge the ion has 1 more
  electron than the
• neutral atom
• Fluorine F 9 e- F - 10
  e-
• Phosphorus P 15 e- P 3- 18 e-
• Anions are named by changing the ending of the
  name to -ide
• Fluorine F 1e- ? F - Fluoride ion
• Oxygen O 2e- ? O 2- Oxide ion
• The charge on an anion can be determined from the
  Group
• number on the Periodic Table

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