Arrangement of Electrons in Atoms

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Chapter 4

• Arrangement of Electrons in Atoms

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4-1 The Development of a New Atomic Model

• Rutherfords model did not explain where
  electrons were what prevented electrons from
  being drawn into nucleus?
• New model arose from experiments involving
  absorption and emission of light by matter

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4-1 Properties of Light

• Light can behave as a wave
• Visible light is a kind of electromagnetic
  radiation (energy that exhibits wavelike behavior
  as it travels through space)
• EM radiation includes X rays, UV and IR light,
  microwaves, radiowaves

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4-1 Properties of Light

• All EM radiation moves at the same speed in a
  vacuum 3.0 x 108 m/s
• Wave motion is repetitive
• Wavelength (?) distance between corresponding
  points on adjacent waves (m, cm, nm)
• Frequency (v) number of waves that pass a given
  point in a specific time, usually one second
  (1/s, Hz)

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4-1 Properties of Light
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4-1 Properties of Light

• Since speed is constant, frequency and wavelength
  are related to each other mathematically
• c ?v
• Wavelength and frequency are INVERSELY
  proportional because their product is a constant.

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4-1 Sample Problem

• Determine the frequency of light with wavelength
  550 nm.
• Convert nm to m
• Use formula c?? to determine v

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4-1 The Photoelectric Effect (Light as a Particle)

• 1900s an experiment that cannot be explained by
  the wave theory of light
• Photoelectric effect refers to the emission of
  electrons from a metal surface when light shines
  on the metal
• For a given metal, no electrons are emitted if
  the lights frequency is below a certain minimum,
  regardless of how intense the light or how long
  it is shone on the metal

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4-1 The Photoelectric Effect

• Wave theory predicts that ANY frequency of light
  could supply enough energy to eject an electron
  from the metal surface
• Wave theory cant explain why light must be of
  certain minimum frequency

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4-1 The Particle Description of Light

• 1900 Max Planck studying emission of light by
  hot objects
• Proposed matter does not emit energy continuously
  but in small, specific amounts called quanta
• Idea is called Quantum Theory. Planck wins Nobel
  prize in 1918 for his work.

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4-1 The Particle Description of Light

• Quantum minimum quantity of energy that can be
  lost or gained by an atom
• Energy of a quantum is related to frequency
• E hv

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4-1 The Particle Description of Light

• 1905 Einstein light has a dual nature
  sometimes it acts like a wave, sometimes it acts
  like a particle
• Light has wave properties
• Light is also like a stream of particles, each
  particle carries a quantum of energy
• Einstein called these particles photons

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4-1 Explanation for the Photoelectric Effect

• Electrons are bound to the atom with a certain
  amount of energy.
• Metal surface must be struck by a photon of light
  carrying at least this amount of energy to knock
  the electron loose.
• Energy and frequency are directly proportional.
  (Eh?)
• Only frequencies equal to or greater than the
  threshold frequency will knock an electron off an
  atom.

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4-1 Sample Problem

• Calculate the energy associated with a photon of
  light of frequency 4.1 x 1014 Hz.

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4-1 Hydrogen-Atom Line-Emission Spectrum

• Spectrum a pattern of energy observed when
  matter absorbs and emits energy
• Ground state lowest energy state of an atom or
  molecule
• Excited state state in which atom or molecule
  has higher PE than ground state

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4-1 Hydrogen-Atom Line-Emission Spectrum

• Current passed through vacuum tube with hydrogen
  gas inside
• Pink light passed through prism to separate into
  specific frequencies of light

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4-1 Hydrogen-Atom Line-Emission Spectrum

• Why does hydrogen give off only specific
  frequencies of light?
• 1913 Niels Bohr proposed a model for hydrogen
  atom that linked the atoms electron with photon
  emission
• Ties line emission spectrum to quantum theory.

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4-1 Bohr Model of the Hydrogen Atom

• Electron can circle nucleus only in allowed
  paths, or orbits
• Orbit closest to nucleus has lowest energy
  (ground state)
• Orbits farther from nucleus have higher energy
  (excited states)
• When electron absorbs energy, it jumps to higher
  orbit
• When electron emits energy, it drops to lower
  orbit

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4-1 Bohr Model of the Hydrogen Atom

• Electron can only exist in certain allowed
  orbits.
• Can only absorb and emit amounts of energy that
  correspond to energy differences between orbits.

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4-1 Bohr Model of the Hydrogen Atom

• Bohrs model did not explain the spectra of atoms
  with more than one electron
• Bohrs theory did not explain the chemical
  behavior of atoms

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4-2 Electrons as Waves

• It was already known that light can behave as a
  particle or a wave.
• 1924 Louis deBroglie asked if electrons could
  also have dual wave-particle nature

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4-2 deBroglies Hypothesis

• Electrons are particles but they can act like
  waves
• A wave confined to a space can only have certain
  frequencies seems to correspond to Bohrs
  quantized electron orbits

• The electron-wave is confined to a certain space
  the region around the nucleus so
  electron-waves can only have certain frequencies,
  which correspond to certain energies (E hv)

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4-2 Wave-Particle Duality of Nature

• Particles can have wave properties.
• Waves can have particle properties.

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4-2 Heisenberg Uncertainty Principle

• If the electron is both a particle and a wave,
  where is it?
• Werner Heisenberg, German physicist, 1927
• Electrons are detected by hitting them with
  photons, but hitting them changes their position
• It is impossible to determine simultaneously the
  position and velocity of an electron

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4-2 The Schrodinger Wave Equation

• 1926 Erwin Schrodinger uses assumption that
  electron behaves as a wave to describe
  mathematically the wave properties of electrons
  and other very small particles (Quantum theory)

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4-2 What does it mean?

• Solutions to the Schrodinger equation are called
  wave functions
• Wave functions can give probability of finding an
  electron at a particular position in the space
  around the nucleus
• An orbital is a 3D region around the nucleus that
  indicates the probable location of an electron

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4-2 Atomic Orbitals and Quantum Numbers

• Quantum numbers specify the properties of atomic
  orbitals and the properties of electrons in
  orbitals.
• Each electron in an atom can be assigned a set of
  four quantum numbers.

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4-2 The Principal Quantum Number

• Symbolized by n
• Indicates the main energy level occupied by an
  electron
• Values of n are positive integers (ex. n 1 is
  the first energy level)
• Principal quantum number also gives approximate
  distance from nucleus/size of energy level or
  shell
• Total number of electrons that can exist in a
  given energy level, n, is equal to 2n2.

Energy level, n Maximum number of electrons, 2n2
1 2
2 8
3 18
4 32
5 50
6 72
7 98
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4-2 Angular Momentum Quantum Number

• Symbolized by l
• Indicates the shape of the orbital
• sublevels
• For each energy level, n, the number of orbital
  shapes possible is equal to n
• The first four shapes are given letter symbols
  (s, p, d and f)

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4-2 Magnetic Quantum Number

• Symbolized by m
• Indicates the orientation of an orbital around
  the nucleus

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s orbital (1 orientation)
p orbital (3 orientations)
d orbitals (5 orientations)
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f orbitals (7 orientations)
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sublevel number of orbitals available number of electrons sublevel can hold
s 1 2
p 3 6
d 5 10
f 7 14
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4-2 Spin Quantum Numbers

• Electrons in orbitals spin on internal axes.
• When charged bodies spin, they induce a magnetic
  field.
• An electron can spin in one of two possible
  directions.
• The spin quantum number has two possible values,
  ½ and ½
• A single orbital can hold a total of two
  electrons, which MUST have opposite spins.

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4-3 Electron Configuration

• The arrangement of electrons in an atom
• Assigns an energy level and sublevel to each
  electron in an atom.

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4-3 Rules Governing Electron Configurations

• The Aufbau Principle an electron occupies the
  lowest-energy orbital available. (aufbau is
  German for building up
• Electrons fill low energy orbitals before filling
  higher energy orbitals.

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4-3 Electron Configuration

• 1s has the lowest energy.
• Energies of sublevels in different main energy
  levels begin to overlap in n3
• Use orbital filling diagram to determine order in
  which sublevels are filled.

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4-3 Rules Governing Electron Configurations

• Pauli Exclusion Principle no two electrons in
  the same atom can have the same set of four
  quantum numbers
• In other words, if two electrons are going to
  occupy the same orbital, they must have opposite
  spin.

-let horizontal line represent orbital -an up
arrow and a down arrow represent two electrons of
opposite spin
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4-3 Rules Governing Electron Configurations

• Hunds rule orbitals of equal energy
  (degenerate orbitals) 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
• Bus seat rule

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4-3 Ways to Represent Electron Configuration
Examples Na P Br Rb K Ar

• Electron Configuration Notation
• Assigns each electron to an energy level and a
  sublevel.

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4-3 Electron Configuration Sample Problems

• Name the elements indicated by the following
  electron configurations
• 1s22s22p63s23p5
• 1s22s22p63s23p64s23d5

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4-3 Electron Configuration Sample Problems

• Write the electron configuration for an element
  that has the following number of electrons
• 7
• 14
• 19
• 33

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4-3 Ways to Represent Electron Configuration

• Orbital Notation uses lines and arrows to
  represent orbitals and electrons
• Example Write the orbital notations for nitrogen
  and oxygen.

N O
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4-3 Ways to Represent Electron Configuration

• Noble Gas Notation to simplify an elements
  electron configuration, use the preceding noble
  gas as shorthand to indicate all the electrons
  possessed by that noble gas
• Example Ne and Na

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4-3 Valence Electrons
How many valence electrons does sodium
have? Bromine? Silicon?

• Valence electrons are electrons in the outermost
  energy level of an atom, farthest from the
  nucleus
• They are important because they are the electrons
  that are usually involved in chemical reactions.

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4-3 Electron Configurations with Special Stability

• Octet the outer energy level is considered
  filled when the s and p sublevels are completely
  filled with 8 electrons
• A filled outer energy level (8 electrons) is a
  very stable electron configuration.
• The noble gases have filled outer energy levels.
  This is why they are unreactive.

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4-3 Electron Configurations with Special Stability
Chromium Copper Molybdenum Silver

• Filled and half-filled sublevels have special
  stability (especially d).
• This fact sometimes results in electron
  configurations that deviate from the Aufbau
  principle.

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