ELECTRONS IN ATOMS

1
ELECTRONS IN ATOMS

• Chapter 5 pages 126-153
• Slides 2-15 are review from Chapter 4

2
The Structure of the Atom

• About 2500 years ago, Greek philosophers thought
  about matter and its composition
• 4th Century B.C. Democritus first suggested the
  idea of atoms (indivisible particles)
• Aristotle did not believe in atoms

3
Antoine Lavoisier (1782)

• Father of Modern Chemistry
• French chemist
• The first to use truly quantitative research
• Observations led to the Law of Conservation
  of Mass
• Identified components of water as hydrogen and
  oxygen

4
Joseph Proust (1799)

• French chemist
• Observed composition of water is always 11
  hydrogen and 89 oxygen by mass
• Studied many other compounds and always found a
  constant composition by mass for a given compound
• This is the Law of Definite Proportions

5
John Dalton (1808)

• English schoolteacher
• Studied the results of Lavoisier and Proust and
  many other scientists
• He wanted and atomic theory to explain the
  experimental evidence
• His theory led to the solid ball model of the
  atom

6
Daltons Atomic Theory

• All matter is composed of indivisible particles
  called atoms.
• Atoms of the same element are identical atoms of
  different elements are different.
• Atoms of different elements chemically combine in
  small whole number ratios to form compounds.
• Chemical reactions occur when atoms are
  separated, combined, or rearranged.

7
Crookes Experiment (1870s)

• Gas Tubes w/2 electrodes (conductors)
• Anode positive
• Cathode negative
• Cathode ray tube (CRT) when voltage was applied
  a beam of light composed of particles was
  deflected by a magnet determined they were
  charged particles

8
J.J. Thomson (1897)

• Was investigating the relationship between matter
  electricity
• Cathode ray tube experiment with a fluorescent
  screen allowed him to measure deflection when a
  magnet was used
• Measured ratio of charge to mass and determined
  particles were identical regardless of the gas
  used or the material of the cathode
• These particles were later named electrons
• This led to the plum pudding model

9
Thomsons Plum Pudding Model

• Nobel Prize 1907
• Pudding the charge and most of the mass of the
  atom
• Plums - charged electrons spread throughout to
  make the atom neutral
• Ions / - charged atoms result from the loss
  or gain of electrons
• Cations positive charge / lost electrons
• Anions negative charge / gain electrons

10
Robert A. Millikan (1909)

• Oil drop experiment suspended fine mist of oil
  droplets between charged plates
• Approximated the mass of an electron to be 1/2000
  the mass of an H atom
• Currently known to be 1/1840 of a H atom
• 9.11 x 10-28 g

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Protons

• Since atoms are neutral, a positive charge must
  also exist in the atom
• Thomson showed that positively charged rays
  existed in the CRT
• Protons finally identified by 1920
• Proton mass is 1836/1837 of H atom
• Mass of Proton 1.67 x 10-24 g

12
Radioactivity Discovered 1896 (Ch 25.1)

• Radioactivity discovered in Uranium by Becquerel
• Radiation energy that is emitted from a source
  and travels through space
• Radioactivity spontaneous radiation from the
  nucleus of an atom
• Marie/Pierre Curie radium polonium

13
Radioactivity (Ch 25.1)

• By 1900 3 types of radiation identified
• Alpha (a) He ions w no elctrons 1/10th the speed
  of light stopped by paper or clothing
• Beta (ß) electrons at high speeds / stopped by a
  few mm of Al
• Gamma (?) form of electromagnetic radiation more
  energetic than X-rays stopped by several cm of
  Pb or more concrete/ no mass or charge

14
Rutherford (1911)

• Gold foil experiment with alpha particles
• Led to the nuclear model of the atom
• Atoms contain a small dense nucleus
• Electrons move around like bees in a hive
• Diameter of nucleus 1/100,000 the size of the
  atom most of the atom is empty space
• 1920 Rutherford proposed neutral particles with
  the same mass as protons

15
Chadwick (1930s)

• Credited with the discovery of neutrons
• Nobel Prize 1935
• Neutron Mass 1.67 x 10-24 g

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Bohrs Model of the Atom (Chapter 5)

• 1913 Niels Bohr
• Linked electron with photon emission
• Electrons circle the nucleus in exact paths
• Paths farther from the nucleus are higher in
  energy
• When electrons drop to lower energy levels, the
  photon emitted is equal to the energy difference
  between the levels

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Problems with Bohrs Model

• It only worked for hydrogen, atoms with more
  electrons did not fit the model
• It did not fully explain the chemical behavior of
  atoms

18
Quantum Model

• Electrons as Waves If light has a dual nature,
  could electrons also?
• DeBroglie electrons considered waves confined
  to the space around the nucleus electrons could
  exist only at certain frequencies quantized
  energy levels
• Electrons like waves can be bent or diffracted
  show interference

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Heisenberg Uncertainty

• Electrons are detected by their interaction with
  photons
• Uncertainty Principle It is impossible to
  determine simultaneously both the position and
  velocity of an electron or any other particles

20
Schrodinger Wave Equation

• 1926 Used dual nature to develop equation that
  treated electrons as waves
• Like Bohr electrons exist in quantized energy
  levels
• Unlike Bohr Electrons do not travel in exact
  pathways, but are located in orbitals of electron
  density probability

21
Atomic Orbitals Quantum Numbers

• Electron Cloud surface drawn where electrons
  are likely to be found (orbital)
• Quantum Numbers Mathematical description of
  electrons in atoms
• 90 probability
• No two electrons have the same set of four
  quantum numbers

22
Principal Quantum Number

• n
• Main energy level n integers
• As n increases, energy and distance from the
  nucleus increases
• Always equals the number of sublevels within the
  principle energy level
• The number of orbitals in any energy level is n2
• Maximum electrons that can occupy a given
  energy level is 2n2.

23
Angular Momentum Quantum Number

• l
• Indicates the shape of an orbital the number of
  shapes possible n
• Also referred to as sublevles
• Its value ranges from 0 to n-1
• Shapes spherical (s, l0), dumbbell (p, l1),
  complex (d, l2), and more complex (f, l3)

24
Magnetic Quantum Number

• Orientation around the nucleus corresponds to
  3-D graph axes (x, y z)
• Its value ranges from l to l
• 1 for s orbitals, 3 for p orbitals, 5 for d
  orbitals, 7 for f orbitals

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Spin Quantum Number

• Which direction the electron is traveling
• 1/2 or -1/2 ( or )
• This satisfies Paulis exclusion principle

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Electron Configurations

• Defined Ways in which electrons are arranged
  around the nuclei of atoms, from lowest to
  highest energy.

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Three Rules for Filling Orbitals

• 1. Aufbau Principle
• Electrons enter orbitals of lowest energy first
• Orbitals within a sublevel are equal energy
• s,p,d,f E order of sublevels
• 2. Pauli Exclusion Principle
• An atomic orbital may describe at most two
  electrons
• Paired electrons have opposite spin
• No 2 electrons can have the same 4 quantum
  numbers
• 3. Hunds Rule
• When electrons occupy orbitals of equal E, one
  electron enters each orbital until all contain
  es with parallel spin. Second electrons are
  then added so they are paired.

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Types of Sublevels/Orbitals

• s 1 orbital 2 electrons
• p 3 orbital 6 electrons
• d 5 orbitals 10 electrons
• f 7 orbitals 14 electrons
• Remember each orbital holds up to 2 electrons.
• http//lrc-srvr.mps.ohio-state.edu/shell-cgi/world
  /genquiz.pl

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Summary of Electron Positions
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PROPERTIES OF LIGHT

• http//www.nelsonthornes.com/secondary/science/sci
  net/scinet/light/waves/wavediag.htm
• Properties of Light
• Electromagnetic radiation
• Electromagnetic spectrum
• Wavelength
• Frequency
• Amplitude
• Speed of light 3.00 x 108 m/s
• C?v
• Separated into continuous spectrum
• ROY G BIV Frequency increases from red to violet

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Hydrogen Line-Emission

• Ground State
• Excited State
• When energy is passed through hydrogen gas, an
  electron is excited to higher energy levels.
  When the electron falls back to its ground state,
  the energy is released as electromagnetic
  radiation. Visible light can be separated into
  separated bands of color known as a line emission
  spectrum. These bands are associated with
  specific frequencies (energy)
• DEMONSTRATIONS

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Photoelectric Effect

• http//www.colorado.edu/physics/2000/quantumzone/p
  hotoelectric.html
• Wave theory predicted light of any frequency
  would produce enough energy to emit an electron,
  however light has to be a minimum frequency
  support of particle theory
• Planck suggested quanta
• Quantum
• Ehv
• Albert Einstein Electromagnetic Radiation has a
  dual wave/particle nature
• Photons

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Links

• http//wine1.sb.fsu.edu/chm1045/notes/Struct/EConf
  ig/Struct08.htm
• http//wine1.sb.fsu.edu/chm1045/notes/Struct/EPeri
  od/Struct09.htm
• http//lrc-srvr.mps.ohio-state.edu/shell-cgi/world
  /genquiz.pl