to review rutherford’s model of the atom to explore the ...€¢to explore the nature of...
TRANSCRIPT
Chapter 11 SMARTBOARD Class Notes.notebook
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• To review Rutherford’s model of the atom • To explore the nature of electromagnetic
radiation • To see how atoms emit light
Objectives
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A. Rutherford’s Atom
…….but there is a problem here!!
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Using Rutherford’s Model Atoms Should Collapse….
…..then how are they so stable?
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When were you last exposed to electromagnetic radiation?
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• What is the nature of electromagnetic radiation?• How are the types of electromagnetic radiation different?
B. Energy and Light
Light can be modeled as a wave, like a wave in water
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• How are the types of light different?
B. Energy and Light
Wavelength, λ (Greek letter “lambda”) units mFrequency, ν (Greek letter “nu”) units s1 or 1/s or Hertz (Hz)Amplitude, peak height – relates to energy in waveSpeed, c : velocity of light in a vacuum is 3x108 m/s
c = ν λ ν (s1)
(m)
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• Electromagnetic radiation
B. Energy and Light
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Wavelength and Frequency of Visible Light
Frequency in Terahertz (1THz = 1012 Hz) and Wavelength in Nanometers (1nm = 109 m)
Velocity of light = wavelength x frequency
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Light Sometimes Behaves in Unwavelike Ways
The Photoelectric Effect• Light shining on a metal surface can cause electrons to be separated from their atoms
• Below a threshold frequency no electrons are emitted however high the intensity
• At the threshold frequency electrons start to be emitted • At higher frequencies electrons have additional kinetic energy
Photoelectric Effect Demo
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• Dual nature of light – Two coexisting models
B. Energy and Light
Wave Photon – packet of energy
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• Different photons (from light of different wavelengths) carry different amounts of energy.
B. Energy and Light
Energy of photon = hν (h is Planck’s Constant = 6.626 x 1034 Js )
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C. Emission of Energy by Atoms
• Atoms can give off light • They first must receive energy and become excited. • The energy is released in the form of a photon.
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Typical Colors From Flame Tests
Na+
Li+ Cu2+
K+
Which flame represents a higher energy transition?
Periodic Table of Fireworks
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• To understand how the emission spectrum of hydrogen demonstrates the quantized nature of energy
• To learn about Bohr’s model of the hydrogen atom
• To understand how the electron’s position is represented in the wave mechanical model
Objectives
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A. The Energy Levels of Hydrogen
• Only certain types of photons are produced when excited Hydrogen atoms release energy. Why?
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A. The Energy Levels of Hydrogen
Atomic states • Excited state – atom with excess energy
• Ground state – atom in the lowest possible state
• When an H atom absorbs energy from an outside source it enters an excited state.
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A. The Energy Levels of Hydrogen
• Energy level diagram
Energy in the photon corresponds to the energy used by the atom to get to the excited state.
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A. The Energy Levels of Hydrogen
• Only certain types of photons are produced when H atoms release energy. Why?
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A. The Energy Levels of Hydrogen
Quantized Energy LevelsSince only certain energy changes occur the H atom must contain discrete energy levels.
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B. The Niels Bohr Model of the Atom (19131915)
Bohr’s model of the atom • Quantized energy levels • Electron moves in a circular orbit • Electrons jump between levels by absorbing or emitting photons of a particular wavelength
• Able to mathematically explain the emission spectrum of Hydrogen (compare with Rutherford)
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B. The Bohr Model of the Atom
Bohr’s model of the atom was not totally correct. • Had difficulties with spectra of larger atoms• Electrons do not move in a circular orbit and don’t seem to behave like discrete particles all the time
Maybe small particles, such as electrons, can also behave like waves having a dual nature (just like photons)……?
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C. The Wave Mechanical Model of the Atom (de Broglie and Schroedinger – mid1920’s) Orbitals • Nothing like orbits • Probability of finding the electron within a certain space
BohrSchrodinger
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• To learn about the shapes of the s, p and d orbitals
• To review the energy levels and orbitals of the wave mechanical model of the atom
• To learn about electron spin
Objectives
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A. The Hydrogen Orbitals
• Orbitals do not have sharp boundaries.
90% boundary
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A. The Hydrogen Orbitals
• Hydrogen has discrete energy levels.
• Called principal energy levels (electron shells)
• Labeled with whole numbers• Energy is related to 1/n2• En = E1/n2 • Energy levels are closer together the further they are from the nucleus
Hydrogen Energy Levels
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A. The Hydrogen Orbitals
• Each principal energy level is divided into sublevels.
Hydrogen Energy Levels
Labeled with numbers and letters Indicate the shape of the orbital
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Orbital Designations
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Orbitals Define the Periodic Table
“Orbitals” are “Energy Levels”
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A. The Hydrogen Orbitals
• The s and p types of sublevel
Hydrogen Energy Levels
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Representation of s, p, d atomic orbitals
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A. The Hydrogen Orbitals
• Why does an H atom have so many orbitals and only 1 electron?
• An orbital is a potential space for an electron. • Atoms can have many potential orbitals. • s, p, d, f orbitals named for sharp, principal, diffuse and fundamental lines on spectra. Further orbitals designated alphabetically
Hydrogen Orbitals
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g g gg
f ff
s p d d
Orbitals
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• Close examination of spectra revealed doublets• Need one more property to determine how electrons are arranged
• Spin – electron modeled as a spinning like a top• Spin is the basis of magnetism
B. The Wave Mechanical Model: Further Development
Electron Spin
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• Pauli Exclusion Principle (Wolfgang Pauli 1925) an atomic orbital can hold a maximum of 2 electrons and those 2 electrons must have opposite spins
• When an orbital contains two electrons (of opposite spin) it is said to be full
B. The Wave Mechanical Model: Further Development
Pauli Exclusion Principle
What are the four descriptors that define an energy level / electron’s position in an atom?
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• To understand how the principal energy levels fill with electrons in atoms beyond hydrogen
• To learn about valence electrons and core electrons
• To learn about the electron configurations of atoms
• To understand the general trends in properties in the periodic table
Objectives
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A. Electron Arrangements in the First 18 Atoms on the Periodic Table
H atom • Electron configuration – electron arrangement – 1s1
• Orbital diagram – orbital is represented as a box with a designation according to its sublevel. Contains arrow(s) to represent electrons (spin)
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A. Electron Arrangements in the First 18 Atoms on the Periodic Table
He atom
Electron configuration – 1s2 Orbital diagram
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A. Electron Arrangements in the First 18 Atoms on the Periodic Table
Li atom
Electron configuration– 1s2 2s1
Orbital diagram
Write the electron configuration and orbital diagrams for Boron, Nitrogen, Fluorine and Argon
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A. Electron Arrangements in the First 18 Atoms on the Periodic Table
Write the full electron configuration of Neon and SulfurDraw an orbital diagram for Magnesium and Chlorine
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Order of Filling of Orbitals
Atoms fill their orbitals in the order of their energies:
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• Orbital filling and the periodic table
B. Electron Configurations and the Periodic Table Dynamic Periodic Table
PT 2a
PT 2b
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• Electron configurations for K through Kr
B. Electron Configurations and the Periodic Table
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If there were more elements……….B. Electron Configurations and the Periodic Table
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A. Electron Arrangements in the First 18 Atoms on the Periodic Table
Classifying Electrons
• Valence electrons – electrons in the outermost (highest) principal energy level of an atom
• Core electrons – inner electrons • Elements with the same valence electron arrangement (same group) show very similar chemical behavior.
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B. Electron Configurations and the Periodic Table
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Using a Noble Gas Shorthand
• We can abbreviate electron configurations by using the configuration of the previous noble gas to cover the first part of the list of orbitals
• Mg is 1s2 2s2 2p6 3s2 or [Ne] 3s2• The noble gas portion is the equivalent to the group of core electrons
• Use the Noble Gas shorthand to show the electron configurations of Carbon, Chlorine and Zirconium
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C. Atomic Properties and the Periodic Table
Metals and Nonmetals Metals tend to lose electrons to form positive ions. Nonmetals tend to gain electrons to form negative ions.
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C. Atomic Properties and the Periodic Table
Atomic Size • Size tends to increase down a column. • Size tends to decrease across a row.
(close to scale) Larger
Effects of Shielding of outer electrons by inner orbitals
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C. Atomic Properties and the Periodic Table
Ionization EnergiesIonization Energy – energy (ΔH) required to remove an electron from an individual atom (gas)
Tends to decrease down a column Tends to increase across a row Changes in an opposite direction to atomic size
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Ionization Energies
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Electron Affinity
Electron Affinity is defined as ΔH for the process:
X(g) + e = X(g)
ΔH = Electron Affinity
Larger
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Electronegativity
• Ionization Energy and Electron Affinity are combined to give Electronegativity – a measure of how well atoms compete for electrons in a bond
Larger