Wave Nature of Light

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Wave Nature of Light

• Light travels through space as a wave.
• There are 2 primary characteristics of waves that
  interest us
• Wavelength (?) the distance between two
  consecutive crests or troughs.
• most often in nm
• Frequency (v or f) the number of wave cycles
  that pass a given point in a unit of time
  (usually per second). Once cycle per second 1
  Hz.

c ? ? c the speed of light, 3.00 x 108
m/s ? in meters ? reciprocal seconds
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Wave Characteristics
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Example 6.1

• The red light associated with the aurora borealis
  is emitted by excited (high energy) oxygen atoms
  at 630.0 nm. What is the frequency of the light.
• Ans 4.759 x 1014 Hz

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Particle Nature of Light

• Photons a stream of particles that give off
  energy in the form of light.
• E photon h? ? E atom (1 mole)
• ? E h? hc/?
• h plancks constant, 6.626 x 10-34 J s
• c speed of light, 3.00 x 108 m/s

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Electromagnetic Spectrum
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Example 6.2

• Referring back to example 1 calculate
• The energy in joules, of a photon emitted by an
  excited oxygen atom.
• Ans 3.153 x 1019 J/mole
• The energy, in kJ, in a mole of such photons.
• Ans 1.899 x 102 kJ/photons

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Atomic Spectra

• Atomic spectra give discrete lines given off at
  specific wavelengths.
• The fact that photons making up atomic spectra
  have only certain discrete wavelengths implies
  that they can have only certain discrete energies
  because these photons are produced when an
  electron moves from one energy level to the next.

These series appear in different regions of the
electromagnetic spectrum.
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Line Emission Spectra
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Bohr Model of the Hydrogen Atom

• There are three points to be made with the Bohr
  Model
• Bohr designated zero energy as the point at
  which the proton and electron are completely
  separated.
• Ordinarily the hydrogen electron is in its lowest
  energy state, referred to as the ground state
  (n1). When an electron absorbs enough energy, it
  moves to a higher, excited state. For hydrogen,
  the first excited state is n2, then n3.
• When an excited electron drops back to its lower
  energy state it gives off energy as a photon of
  light.

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The Rydberg Equation

• Bohr derived the following equation in applying
  his model to the hydrogen atom
• v RH 1 1__
• h (nlo)2 (nhi)2
• RH Rydbergs constant 2.18 x 10-18 J
• Balmer series when an electron jumps down to n2
  from n3,4,5..
• Lyman series when an electron jumps down to n1
  from n2,3,4,.

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Example 6.3

• Calculate the wavelength, in nm, fo the line in
  the Balmer series that results from the
  transition when the electron is in n4.
• Ans 486.0 nm

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Problem...

• Bohrs model was flawed. We cannot assume that
  electrons move about in specified orbitals.
• DeBroglie purposed that if light can behave as
  particles (photons) then electrons can act as
  waves (wave-particle duality).
• This led to wave mechanics and the quantum
  mechanical model of the atom. This differs from
  the Bohr model, mainly, in that
• The kinetic energy of an electron is inversely
  related to the volume of the region to which it
  is confined.
• It is impossible to specify the precise position
  of an electron in an atom at a given instant.

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DeBroglie Table of Wavelengths
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Electron Configuration

• Energy Levels (1-7) looking at the periodic
  table you can tell how many energy levels an atom
  has by looking at the period the electron is in.
• Sublevels (s,p,d,f) looking at the periodic
  table you can tell sublevels based on the 4
  blocks the periodic table is broken up into.
• Orbitals (3-D orientations)
• s_, p _ _ _, d _ _ _ _ _, f _ _ _ _ _ _ _

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Examples of Electron Configuration
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Orbital Occupancy for first 10 elements.
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Predicting Electron Configuration

• Some rules to follow when predicting electron
  configuration
• Start at lowest energy level possible (hydrogen).
• Follow atomic numbers when filling.
• Use arrows to represent electrons up arrows
  fill first.

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Rules to know by name.

• Pauli Exclusion Principle
• no 2 electrons in an atom may have the same set
  of quantum numbers.
• Hunds Rule
• When several orbitals of equal energy are
  available, as in a given sublevel, electrons
  enter singly with parallel spins (up arrows
  first).
• Aufbau Principle
• The principle postulates a hypothetical process
  in which an atom is "built up" by progressively
  adding electrons. As they are added, they assume
  their most stable conditions with respect to the
  nucleus and those electrons already there (lowest
  energy level first).

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Electron Arrangement in Ions

• Transition metal cations to the right of the
  scandium group do not form ions with noble-gas
  configurations (like most main group elements),
  they would have to lose four or more electrons to
  do so. In transition metals the outer s-electrons
  are usually lost first to form positive ions.
• For example Mn
• Mn2
• In ions like Fe, electrons will be lost from 4s
  first then the 3d. This is usually referred to as
  the first in, first out rule.

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Example of Fe3 Ion
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Electron energy levels in order of increasing
energy (pg 142).
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Example 6.6 6.9

• Find the electron configuration of iodine,
  sulfur, iron, copper, and chromium.
• Find the electron configuration for Fe2 and Br1-
• Show the configuration
• Write the abbreviated notation

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Magnetism

• Paramagnetic If there are unpaired electrons
  present the solid will be attracted into the
  field
• Diamagnetic If the atoms in the solid contain
  only paired electrons it is slightly repelled by
  the field.

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Quantum Numbers

• The principal quantum number
• Symbolized by n, basically the energy level the
  electron is in.
• The orbital quantum number
• Symbolized by l, basically represents the
  sublevel the electron is in s,p,d, or f.
• values for l s 0, p l 1, d l 2, f l
  3
• The letters s,p,d, and f come from the
  adjectives used to describe spectral lines
  sharp, principal, diffuse, fundamental.
• The magnetic quantum number
• Symbolized by m l, this determines the direction
  in space of the electron cloud surrounding the
  nucleus.
• All of the orbitals in a sublevel have the same
  energy
• s __ p __ __ __ d __ __ __ __ __ f
  __ __ __ __ __ __ __
• The spin quantum number
• Symbolized by ms, this represents the electron
  spin.
• ms can equal 1/2 (up arrow) or -1/2 (down arrow)

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Examples

• Give the quantum numbers for the outermost
  electron in neon, copper, and barium.

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Example 6.4 6.5

• Consider the following set of quantum numbers,
  which ones could not occur
• (a) 3,1,0,1/2
• (b) 1,1,0,-1/2
• (c) 2,0,0, ½
• (d) 4,3,2,1/2
• (e) 2,1,0,0

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Periodic Trends

• The Periodic Law The chemical and physical
  properties of elements are a periodic function of
  atomic number.
• Specific trends occur because of this
• Atomic Radius
• Ion Radius
• Ionization Energy
• Electronegativity

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Atomic Radius

• One half the distance of closest approach between
  atoms in an elemental substance. Basically from
  the center of the nucleus of an atom to its
  outermost electrons.
• Decreases across a period from left to right on
  the periodic table
• Increases down a group on the periodic table

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Ionic Radius

• Positive ions are smaller than the metal atoms
  from which they are formed
• Negative ions are larger than the nonmetal atoms
  from which they are formed

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Ionization Energy

• The energy required to remove an electron from
  its outermost shell.
• There can be a first, second, third, and so on
  ionization energy.
• For example, in Magnesium 2 there is a large
  jump in ionization energy from the 2nd to 3rd
  ionization energies. Why?
• Increases across the periodic table from left to
  right
• Decreases down the periodic table
• Noble gases generally have the highest
  ionization energies except when compared to
  fluorine. There are several exceptions to the
  trend (reference the textbook table of IEs).

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Ionization Energies for Be
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Electron Affinity Electronegativity

• Electron Affinity is the actual energy change
  associated with the gaining of an electron.
• Electronegativity is he ability to attract an
  electron.
• Increases across the periodic table from left to
  right fluorine is the most electronegative
  element.
• Decreases down the periodic table
• There are a few exceptions to this trend you
  must know them.
• Noble gases essentially (Kr and Xe are an
  exception) have no electronegativities.

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MC 1

• Use these answers for questions 1 - 3.
• (A) O(B) La(C) Rb(D) Mg(E) N
• 1. What is the most electronegative element of
  the above?
• 2. Which element exhibits the greatest number of
  different oxidation states?
• 3. Which of the elements above has the smallest
  ionic radius for its most commonly found ion?

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MC 2

• Use these answers for questions 1-4
• (A) Heisenberg uncertainty principle(B) Pauli
  exclusion principle(C) Hund's rule (principle of
  maximum multiplicity)(D) Shielding effect(E)
  Wave nature of matter
• 1. Can be used to predict that a gaseous carbon
  atom in its ground state is paramagnetic
• 2. Explains the experimental phenomenon of
  electron diffraction
• 3. Indicates that an atomic orbital can hold no
  more than two electrons
• 4. Predicts that it is impossible to determine
  simultaneously the exact position and the exact
  velocity of an electron

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MC 3

• 1s22s22p63s23p3
• Atoms of an element, X, have the electronic
  configuration shown above. The compound most
  likely formed with magnesium, Mg, is
• (A) MgX(B) Mg2X(C) MgX2(D) MgX3(E) Mg3X2

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MC 4

• Which of the following represents the ground
  state electron configuration for the Mn3 ion?
  (Atomic number Mn 25)
• (A) 1s2 2s2 2p6 3s2 3p6 3d4(B) 1s2 2s2 2p6 3s2
  3p6 3d5 4s2(C) 1s2 2s2 2p6 3s2 3p6 3d2 4s2(D)
  1s2 2s2 2p6 3s2 3p6 3d8 4s2(E) 1s2 2s2 2p6 3s2
  3p6 3d3 4s1

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MC 5

• One of the outermost electrons in a strontium
  atom in the ground state can be described by
  which of the following sets of four quantum
  numbers?
• (A) 5, 2, 0, 1/2(B) 5, 1, 1, 1/2(C) 5, 1, 0,
  1/2(D) 5, 0, 1, 1/2(E) 5, 0, 0, 1/2

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MC 6

• The elements in which of the following have most
  nearly the same atomic radius?
• (A) Be, B, C, N(B) Ne, Ar, Kr, Xe(C) Mg, Ca,
  Sr, Ba(D) C, P, Se, I(E) Cr, Mn, Fe, Co

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MC 7

• Ca, V, Co, Zn, As
• Gaseous atoms of which of the elements above are
  paramagnetic?
• (A) Ca and As only(B) Zn and As only(C) Ca, V,
  and Co only(D) V, Co, and As only(E) V, Co, and
  Zn only

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FRQ 1

• (a) A major line in the emission spectrum of neon
  corresponds to a frequency of 4.34?1014 s-1.
  Calculate the wavelength, in nanometers, of light
  that corresponds to this line.
• (b) In the upper atmosphere, ozone molecules
  decompose as they absorb ultraviolet (UV)
  radiation, as shown by the equation below. Ozone
  serves to block harmful ultraviolet radiation
  that comes from the Sun.
• O3 (g) ? O2 (g) O (g)
• A molecule of O3 (g) absorbs a photon with a
  frequency of 1.00?1015 s-1.
• (i) How much energy, in joules, does the O3(g)
  molecule absorb per photon?
• (ii) The minimum energy needed to break an
  oxygen-oxygen bond in ozone is 387 kJ mol-1. Does
  a photon with a frequency of 1.00?1015 s-1 have
  enough energy to break this bond? Support your
  answer with a calculation.

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FRQ 2

• Discuss some differences in physical and chemical
  properties of metals and nonmetals. What
  characteristic of the electronic configurations
  of atoms distinguishes metals from nonmetals? On
  the basis of this characteristic, explain why
  there are many more metals than nonmetals.

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FRQ 3

• Use the details of modern atomic theory to
  explain each of the following experimental
  observations.
• (a) Within a family such as the alkali metals,
  the ionic radius increases as the atomic number
  increases.
• (b) The radius of the chlorine atom is smaller
  than the radius of the chloride ion, Cl-. (Radii
  Cl atom 0.99Å Cl- ion 1.81 Å)
• (c) The first ionization energy of aluminum is
  lower than the first ionization energy of
  magnesium. (First ionization energies 12Mg 7.6
  ev 13Al 6.0 ev)
• (d) For magnesium, the difference between the
  second and third ionization energies is much
  larger than the difference between the first and
  second ionization energies. (Ionization energies
  for Mg 1st 7.6 ev 2nd 14 ev 3rd 80 ev)

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FRQ 4

• The elements are numbered randomly. Use the
  information in the table to answer the following
  questions.
• (a) Which element is most metallic in character?
  Explain your reasoning.
• (b) Identify element 3. Explain your reasoning.
• (c) Write the complete electron configuration for
  an atom of element 3.
• (d) What is the expected oxidation state for the
  most common ion of element 2?
• (e) What is the chemical symbol for element 2?
• (f) A neutral atom of which of the four elements
  has the smallest radius?

1st IE kJ mol-1 2nd IE kJ mol-1 3rd IE kJ mol-1
Element 1 1,251 2,300 3,820
Element 2 496 4,560 6,910
Element 3 738 1,450 7,730
Element 4 1,000 2,250 3,360
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FRQ 5

• (a) Write the ground state electron configuration
  for an arsenic atom, showing the number of
  electrons in each subshell.
• (b) Give one permissible set of four quantum
  numbers for each of the outermost electrons in a
  single As atom when it is in its ground state.
• (c) Is an isolated arsenic atom in the ground
  state paramagnetic or diamagnetic? Explain
  briefly.
• (d) Explain how the electron configuration of the
  arsenic atom in the ground state is consistent
  with the existence of the following known
  compounds Na3As, AsCl3, and AsF5.

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Equations 1 (reaction prediction)

• (a) Chlorine gas, an oxidizing agent, is bubbled
  into a solution of potassium bromide at 25C.
• (i) Balanced equation
• (ii) What state(s) of matter will be present at
  the end of the reaction.
• (b) Solid strontium hydroxide is added to a
  solution of nitric acid.
• (i) Balanced equation
• (ii) How many moles of strontium hydroxide would
  react completely with 500. mL of 0.40 M nitric
  acid?
• (c) A solution of barium chloride is added drop
  by drop to a solution of sodium carbonate,
  causing a precipitate to form.
• (i) Balanced equation
• (ii) What color if any will the precipitate have?

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Equations 2

• (a) A barium nitrate solution and a potassium
  fluoride solution are combined and a precipitate
  forms.
• (i) Balanced equation
• (ii) If equimolar amounts of barium nitrate and
  potassium fluoride are combined, which reactant,
  if any, is the limiting reactant? Explain.
• (b) A piece of cadmium metal is oxidized by
  adding it to a solution of copper(II) chloride.
• (i) Balanced equation
• (ii) List two visible changes that would occur
  in the reaction container as the reaction is
  proceeding.
• (c) A hydrolysis reaction occurs when solid
  sodium sulfide is added to distilled water.
• (i) Balanced equation
• (ii) Indicate whether the pH of the resulting
  solution is less than 7, equal to 7, or greater
  than 7. Explain.

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Graphics

• (Silberberg, Martin S.. Chemistry The Molecular
  Nature of Matter and Change, 5th Edition.
  McGraw-Hill)