Honors Chemistry, Chapter 4

1
Chapter 4 Arrangement of Electrons in Atoms
2
Properties of Light

• Electromagnetic radiation (light) is a form of
  energy that exhibits wavelike behavior as it
  travels through space.
• All forms of electromagnetic radiation form the
  electromagnetic spectrum including
• X-rays -Infrared light
• Ultraviolet light -Microwaves
• Visible light -Radio Waves

3
Wavelength and Frequency

• The wavelength (l) of an electromagnetic wave is
  the distance between corresponding points on
  adjacent waves.
• The frequency (n) is defined as the number of
  waves that pass a given point in one second.
• The product of wavelength and frequency is a
  constant, the speed of light
• lnc

4
Photoelectric Effect

• The photoelectric effect refers to the emission
  of electrons from a metal when light shines on
  the metal.
• Max Plank discovered that light is emitted from
  hot bodies in small, specific amounts called
  quanta.
• A quantum is the minimum quantity of energy that
  can be lost or gained by an atom. E hn, where
  h is Planks constant.

5
The Photon

• In 1905, Albert Einstein proposed that
  electromagnetic radiation has a dual
  wave-particle nature. (1921 Nobel Prize)
• Einstein called the particle of light the photon.
• A photon is a particle of electromagnetic
  radiation having zero mass and carrying a quantum
  of energy. Ephoton hn

6
Atomic Energy States

• The ground state of an atom is the lowest energy
  state that an atom can posses.
• An excited state of an atom is one in which the
  atom has higher potential energy than its ground
  state.

7
Hydrogen-Atom Emission Spectrum

• When electricity is passed through hydrogen, the
  atoms emit a characteristic purple colored light.
• Classical theory predicted that the emission from
  a hydrogen atom should be continuous (white
  light).
• Instead, hydrogen was shown to exhibit a line
  emission spectrum.

8
Bohr Theory of the Atom

• The Bohr theory of the atom proposed that the
  electrons were in specific energy levels and the
  colors of light emitted corresponded to
  transitions of electrons between those energy
  levels.
• Bohrs theory explained the hydrogen atom but
  none of the other atoms on the periodic table!

9
Chapter 4, Section 1 Review

• Explain the relationship among the speed,
  wavelength, and frequency of electromagnetic
  radiation.
• Discuss the dual wave-particle nature of light.
• 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 H atom.

10
Electrons as Waves

• De Broglie hypothesized that electrons be
  considered waves confined to the space around the
  nucleus.
• Electrons were shown to exhibit deffraction, the
  bending of a wave as it passes by the edge of an
  object and interference, a decrease and increase
  in energy when waves overlap.

11
Heisenberg Uncertainty Principle

• The Heisenberg uncertainty principle states that
  it is impossible to determine simultaneously both
  the position and velocity of an electron or any
  other particle.

12
Schrodinger Wave Equation

• Quantum theory used the Schrodinger wave equation
  to describe mathematically the wave properties of
  electrons and other very small particles.
• Now the energy levels in all atoms, not just the
  hydrogen atom could be understood.
• An orbital is a three-dimensional region around
  the nucleus that indicates the probable location
  of an electron.

13
Atomic Quantum Numbers
14
Orbital Letter Designations
15
Chapter 4, Section 2 Review

• Discuss De Broglies role in the development of
  the quantum model of the atom.
• Compare and contrast the Bohr model and the
  quantum model of the atom.
• Explain how the Heisenburg uncertainty principle
  and the Schrodinger wave equation led to the idea
  of atomic orbitals.

16
Review Continued

• List the four quantum numbers and describe their
  significance.
• Relate the number of sublevels corresponding to
  each of the atoms main energy levels, the number
  of orbitals per sub level and the number of
  orbitals per main energy level.

17
Electron Configuration

• The arrangement of electrons in an atom is known
  as the atoms electron configuration.
• The electronic energy levels defined by quantum
  theory are filled with electrons using three
  principles
• Aufbau principle
• Pauli exclusion principle
• Hunds rule

18
Aufbau (Building Up) Principle

• In the ground state, an electron occupies the
  lowest-energy orbital that can receive it.

19
Pauli Exclusion Principle

• No two electrons in the same atom can have the
  same set of four quantum numbers.

20
Hunds Rule

• Orbitals of equal energy are occupied by one
  electron before any orbital is occupied by a
  second electron, and all electrons in singly
  occupied orbitals must have the same spin.

21
Orbital Notation

• In orbital notation the electrons are represented
  as arrows and each energy level is shown by a
  horizontal line
• __ __ __ __ __
• N 1s 2s 2px 2py 2pz

22
Electron Configuration Notation

• In electron configuration notation, the electrons
  are indicated as superscripts for each level
• N 1s2 2s2 2p3

23
Sample Problem 4-1

• The electron configuration for boron is
  1s2 2s2 2p1. How many electrons does the atom
  have? What is the atomic number? Write the
  orbital notation for boron.
• Electrons 2 2 1 5
• Atomic number 5
• __ __ __
  __ __
• Orbital notation B 1s 2s 2px 2py 2pz

24
Electron Shells

• The highest occupied level is the
  electron-containing main energy level with the
  highest principle quantum number.
• The inner-shell electrons are those which are not
  in the highest occupied energy level.

25
The Noble Gas Configuration

• A noble-gas configuration is an outer main energy
  level fully occupied, in most cases, by eight
  electrons.
• It is designated as the noble gas symbol
  surrounded by brackets Ne
• So, for example, the electron configuration for
  sodium may be either
• Na 1s2 2s2 2p6 3s1 or Na Ne 3s1

26
Sample Problem 4-2

• Write both the complete electron-configuration
  notation and the noble-gas notation for iron, Fe.
• Write the orbital notation for iron. How many
  orbitals have electrons? How many of these are
  filled? How many levels are not filled? In
  which level are the unpaired electrons?

27
Sample Problem 4-2 continued

• Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6
• Or
• Fe Ar 4s2 3d6
• __ __ __ __ __ __ __ __ __
• Fe 1s 2s 2px 2py 2pz 3s 3px 3py 3pz
• __ __ __ __ __ __
• 4s 3d 3d 3d 3d 3d

28
Sample Problem 4-3

• Write both the complete electron-configuration
  notation and the noble-gas notation for a
  rubidium atom.
• Identify the elements in the second, third, and
  fourth periods that have the same number of
  highest-energy-electrons as rubidium.

29
Chapter 4, Section 3 Review

• List the total number of electrons needed to
  fully occupy each main energy level.
• State the Aufbau principle, the Pauli exclusion
  principle, and Hunds rule.
• Describe the electron configurations for the atom
  of any element using orbital notation, electron
  configuration notation, and when appropriate,
  noble-gas notation.