Transcript Wave Nature of Light

Wave Nature of Light
• Light travels through space as a wave.
• There are 2 primary characteristics of
waves that interest us:
– Wavelength (λ): 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=ν×λ
c = the speed of light, 3.00 x 108 m/s
λ = in meters
ν = reciprocal seconds
Wave Characteristics
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
Particle Nature of Light
• Photons – a stream of particles that give
off energy in the form of light.
• E photon = hν = Δ E atom (1 mole)
• Δ E = hν = hc/λ
– h = planck’s constant, 6.626 x 10-34 J s
– c = speed of light, 3.00 x 108 m/s
Electromagnetic Spectrum
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
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.
Line Emission Spectra
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 (n=1).
When an electron absorbs enough energy, it moves to
a higher, excited state. For hydrogen, the first excited
state is n=2, then n=3.
– When an excited electron drops back to its lower
energy state it gives off energy as a photon of light.
The Rydberg Equation
• Bohr derived the following equation in applying
his model to the hydrogen atom:
v = RH
h
[1
–
(nlo)2
1__
]
(nhi)2
• RH = Rydberg’s constant = 2.18 x 10-18 J
• Balmer series: when an electron jumps down to
n=2 from n=3,4,5…..
• Lyman series: when an electron jumps down to
n=1 from n=2,3,4,….
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 n=4.
– Ans: 486.0 nm
Problem……...
• Bohr’s 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.
DeBroglie Table of Wavelengths
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 _ _ _ _ _ _ _
Examples of Electron Configuration
Orbital Occupancy for first 10 elements.
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.
Rules to know by name….
• Pauli Exclusion Principle:
– no 2 electrons in an atom may have the same set of
quantum numbers.
• Hund’s 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).
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:
Mn+2:
• 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.
Example of Fe3+ Ion
Electron energy levels in order of
increasing energy (pg 142).
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
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.
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)
Examples
• Give the quantum numbers for the
outermost electron in neon, copper, and
barium.
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
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
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
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
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
IE’s).
Ionization Energies for Be
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.
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?
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
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
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
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
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
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
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.
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.
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)
FRQ #4
1st
IE
2nd
IE
3rd
IE
kJ mol-1 kJ mol-1 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
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?
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.
Equations #1
(reaction prediction)
(a)
Chlorine gas, an oxidizing agent, is bubbled into a
solution of potassium bromide at 25°C.
(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?
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.
Graphics:
• (Silberberg, Martin S.. Chemistry: The
Molecular Nature of Matter and Change,
5th Edition. McGraw-Hill)