Chapter 4 Arrangement of Electrons in Atoms

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Chapter 4Arrangement of Electrons in Atoms
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Section 4.1The Development of a New Atomic Model

• Objectives
• Discuss the significance of the photoelectric
  effect and the line-emission spectrum of hydrogen
  to the development of the atomic model.
• Describe the Bohr model of the hydrogen atom.

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Electromagnetic radiation energy that exhibits
wavelike behavior All electromagnetic radiation
travels at the same speed. Different types are
the result of different wavelengths and
frequencies.
This spectrum shows the different forms of
electromagnetic radiation.
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• Light is a small part of the electromagnetic
  spectrum.

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Absorption/Emission Spectra

• When light strikes a metal, the metal ejects e-
  from the surface and creates an electric current
  Photoelectric effect
• Remember the cathode-ray tube?
• When the glass tube was filled with a pure gas
  and an electric current passed through, the gas
  will gave off light.
• Different gases give off different colors of
  light.
• If this light is passed through a prism, a
  series of bright lines is seen (emission
  spectrum).
• Every element has a distinct emission spectrum.

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So, whats happening to create the emission
spectrum?
Carbon
Oxygen
Iron
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• Rutherfords model of the atom provided
  information about the structure of atoms, it did
  not explain where the electrons were located in
  the space surrounding the nucleus.

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Albert Einstein
Max Planck

• In 1900, Max Planck suggested that objects could
  give off energy in small, specific amounts he
  called quanta.
• A quantum (singular of quanta) is the minimum
  amount of energy that can be lost or gained by an
  atom.

• Einstein proposed that different elements require
  different frequencies of energy to eject electrons

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• So scientists agreed
• Light is a form of energy
• Different colors of light have different levels
  of energy on the electromagnetic spectrum
• Atoms of different elements had different values
  for a quantum (the minimum amount of energy they
  can gain and lose)
• When atoms of different elements absorb their
  quantum of energy they can temporarily eject
  electrons
• Different colors of light are created by the
  different levels of energy being absorbed, and
  then given off, by electrons.

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The question still remained Why would different
elements absorb different amounts of energy and
then give off different light colors when energy
was applied and they ejected electrons? What was
it about their structures that allowed this?
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Neils Bohr Model of the H Atom

• When an e- is hit by light energy, it absorbs
  the energy.
• If the energy is of the correct frequency
  (quantum), the e- will jump to another energy
  level (excited state vs ground state).
• The electron cannot stay in excited state so it
  falls back to the ground state. It cannot take
  the energy with it so it releases the energy in
  the form of LIGHT energy!
• We see this energy as different colors of light!

When the electron returns to ground state, it
emits energy in the form of different colors of
light depending on frequency of energy gained.
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Section 4.2The Quantum Model of the Atom

• Objectives
• Compare and contrast the Bohr model and the
  quantum model of the atom.
• List the 4 quantum numbers, and describe their
  significance.
• Relate the number of sublevels corresponding to
  each of an atoms main energy levels, the number
  of orbitals per sublevel, and the number of
  orbitals per main energy level.

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Todays Quantum Model of the Atom

• Orbital -- a 3-d region around the nucleus that
  indicates the probable location of an e- .
• Quantum theory -- describes mathematically the
  wave properties of e- and other very small
  particles
• Quantum numbers -- numbers or letters that
  specify the properties of atomic orbitals and the
  properties of e- in orbitals

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4 Quantum Numbers

• Principal Quantum Number
• Angular Momentum Quantum Number
• Magnetic Quantum Number
• Spin Quantum Number

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Principal quantum number (n) - indicates the
main energy level occupied by the e- n
1,2,3,4, etc.
Representation of Bohrs proposal of orbitals.
Lower numbered levels are closer to nucleus and
of lower energy.
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Angular Momentum Quantum Number indicates the
orbital and its shape s, p, d, f
The maximum number of electrons in an energy
level is 2n2, thus a shell with n 2 may hold a
maximum of 8 electrons.
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Magnetic Quantum Number orientation of orbital
around the nucleus (the axis, or axes, it is on
or between) uses x, y, z, xy, yz, etc.
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These orbitals can overlap, cause interference
with each other, and affect over all energy of
each other.
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Spin Quantum Number spin state of an electron
in an orbital can be ½, -½ - often
represented with arrows
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Review

• 4 quantum numbers describe each e- .
• Principal quantum number - indicates the main
  energy level occupied by the e- 1,2,3,4, etc.
• Angular Momentum Quantum Number indicates the
  orbital and its shape s, p, d, f
• Magnetic Quantum Number orientation of orbital
  around the nucleus (the axis, or axes, it is on
  or between) x, y, z, xy, yz, etc.
• Spin Quantum Number spin state of an electron
  in an orbital can be ½, -½ - often
  represented with
• Soto accurately describe the energy (location)
  of electron, all 4 quantum numbers must be used.

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Section 4.3Electron Configurations

• Objectives
• List the total number of electrons needed to
  fully occupy each main energy level
• State and explain the Aufbau principle, the Pauli
  exclusion principle, and Hunds rule.
• Write orbital notations, electron-configuration
  notation and noble-gas notation for atoms.

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Electron Configuration

• Electron configuration -- the arrangement of e-
  in an atom
• Electron configurations summarize the locations
  of each e- in atoms.
• Like all systems in nature, e- tend to assume
  arrangements that have the lowest possible
  energies
• Ground-state configuration -- lowest energy
  arrangement of the e- for each element

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3 basic rules help govern these ground-state
configurations.

1. Aufbau principle -- an e- occupies the
   lowest-energy orbital possible
2. Pauli exclusion principle -- no 2 e- in the
   same atom can have the same set of 4 quantum
   numbers
3. Hunds rule -- orbitals of equal energy are each
   occupied by one e- before any orbital is
   occupied by a second e- , and all e- in singly
   occupied orbitals must have the same spin

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Types of Electron Configurations

• Orbital Notation includes all 4 quantum
• numbers (spin is indicated with an arrow)

Fluorine 9 e-
_____ _____ _____ _____
_____ 1s 2s 2px
2py 2pz
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Complete Orbital Notations for Elements 1 10
on the Periodic Table

• Hydrogen
• Helium
• Lithium
• Beryllium
• Boron
• Carbon
• Nitrogen
• Oxygen
• Fluorine
• Neon

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Types of Electron Configurations
2. Electron Configuration Notation includes
principal and Angular Momentum (combines
magnetic and ignores spin)
Chromium 24 e- 1s2 2s2 2p6 3s2 3p6
4s2 3d4
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Complete Electron Configurations for Elements 11
20 on the Periodic Table

1. Sodium
2. Magnesium
3. Aluminum
4. Si
5. P
6. S
7. Cl
8. Ar
9. K
10. Ca

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Types of Electron Configurations

• Noble Gas Notation abbreviates part of the
• electron configuration by using the Noble
• Gas Symbol just prior to the element and
• adds the rest of the electron configuration

Magnesium 12e- Mg Ne 3s2
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Complete Noble-Gas Notations for the following
elements
S Ca Rb Li O Al
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Chapter 4 Study Guide

1. Know the contributions of Plank and Einstein to
   the Quantum Theory.
2. Be able to define ground state and excited
   state and know what happens to cause an atom to
   moved from ground state to excited state and what
   happens when it moves back to the ground state.
3. Know the 4 quantum numbers by name and
   descriptions.
4. Know the 3 rules Aufbau, Pauli Exclusion, and
   Hund and know how they are applied when doing
   configurations.
5. Know the order of energy levels and orbitals
   from lowest energy to highest energy.
6. Be able to write configurations orbital
   notation, electron configuration, and noble gas
   notations.
7. Be able to use configurations to identify
   elements.
8. Know the s, p, d, and f blocks on the Periodic
   Table so you can double check your work.

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Class work

• A quantum of electromagnetic energy is called
• a(n) .
• A bright-light spectrum of an atom is caused by
• the energy released when electrons .
• Explain the difference in an atoms ground state
• and excited state.
• Name and describe the information provided
• by all four quantum numbers.
• Describe the shapes of the s, p, and d orbitals.
• How many possible orientations can each
• of the following sublevels have?
• p sublevel?
• d sublevel?
• f sublevel?

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• How many electrons can be held in a
• p sublevel?
• d sublevel?
• f sublevel?
• How many total electrons can be held in each
• energy level?
• 1st
• 2nd
• 3rd
• 4th

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• Explain each of the following rules governing
• electron configurations.
• Aufbau principle
• Pauli exclusion principle
• Hunds rule
• Do the following for an atom of Cesium.
• Orbital Notation
• Electron Configuration
• Noble Gas Notation