Adv Chem Chapt 7

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Transcript Adv Chem Chapt 7

ATOMIC
STRUCTURE
AND
PERIODICITY
Advanced Chemistry 7
ELECTROMAGNE
TIC RADIATION
ELECTROMAGNETIC
RADIATION
One of the ways that energy travels through
space is by electromagnetic radiation.
 light from the sun
 X-rays
 microwaves
WAVE CHARACTERISTICS
 Wavelength
(λ) – is the distance between
two consecutive peaks or troughs in a
wave.
 Frequency (ν) – is the number of waves
(cycles) per second that pass a given point
in space.

units – hertz or waves/sec (s-1)
 Speed
(c) – all types of electromagnetic
radiation travel at the speed of light.
2.9979 x 108 m/s
 c = λν

ELECTROMAGNETIC
RADIATION
ELECTROMAGNETIC RADIATION
THE NATURE OF
MATTER
WAVE AND PARTICLE
DUALITY
Planck found that matter could only absorb
or emit energy in whole number multiples
of the quantity hν.
 h is Planck’s constant = 6.626 x10-34 Js
 ΔE = hν
 Transfer of energy is not continuous but is
quantized and can occur only in discrete
amounts called quantum. Thus energy
has particle properties as well as wave
properties.
EINSTEIN
WAVE AND PARTICLE
DUALITY
Einstein proposed that
electromagnetic radiation was also
quantized and could be viewed as a
stream of “particles” called photons.
 Ephoton = hv = hc/λ
THE PHOTOELECTRIC
EFFECT
The photoelectric effect refers to the
phenomenon in which electrons are
emitted from the surface of a metal when
light strikes it.
1. No electrons are emitted by a metal below
a specific threshold frequency (vo)
2. For light with frequency lower than the
threshold frequency, no electrons are
emitted regardless of intensity of the
light.
THE PHOTOELECTRIC
EFFECT
3.
4.
For light with frequency greater than
the threshold frequency, the number of
electrons emitted increases with the
intensity of the light.
For light with frequency greater than
the threshold frequency, the kinetic
energy of the emitted electrons increases
directly with frequency of the light.
THE PHOTOELECTRIC
EFFECT
These observations can be explained by
assuming that electromagnetic radiation
is quantized (consists of photons), and
that the threshold frequency represents
the minimum energy required to remove
the electron from the metal’s surface.
 Minimum energy required to remove an
electron = Eo = hvo

KEelectron = ½ mv2 = hv – hvo
PLANCK AND EINSTEIN
CONCLUSIONS
 Energy
is quantized. It can occur
only in discrete units called quanta.
 Electromagnetic radiation, which
was previously thought to exhibit
only wave properties, seems to show
certain characteristics of particulate
matter as well. This phenomenon is
sometimes referred to as the dual
nature of light.
WAVE PARTICLE DUALITY
The main significance of the equation
E = mc2 is that energy has mass.
 m = E/c2
LOUIS DE BROGLIE (18921987)
Since light which previously was
thought to be purely wavelike, was
found to have certain characteristics
of particulate matter. But is the
opposite also true? Does matter
have that is normally assumed to be
particulate exhibit wave properties?
LOUIS DE BROGLIE (18921987)
 de
Broglie’s equation allows us to
calculate the wavelength for a
particle:
h

mv
DE BROGLIE’S PROOF
LOUIS DE BROGLIE (18921987)
Conclusion: Energy is really a form of
matter, and all matter shows the
same types of properties. All matter
exhibits both particulate and wave
properties.
THE ATOMIC
SPECTRUM OF
HYDROGEN
SPECTRUM
A
continuous spectrum results
when white light passes through a
prism and all wavelengths (colors)
are shown.
 An emission spectrum produces
only a few lines of color that is
limited to discrete wavelengths
produced by an atom. This is called
a line spectrum and is specific to
each atom.
HYDROGEN LINE
SPECTRUM
 The
significance of the line spectrum
is that it indicates that only certain
energies are allowed for the electron
in the hydrogen atom. In other
words the energy of the electron in
the hydrogen atom is quantized
HYDROGEN LINE
SPECTRUM
THE BOHR
MODEL
NIELS BOHR
Bohr developed a
quantum model for
the hydrogen atom
that allowed for
only specific energy
levels around the
atom that
corresponded with
specific radii.
NIELS BOHR (1885-1962)
 The
most important equation to come
from Bohr’s model is the expression for
the energy levels available to the
electron in the hydrogen atom.
E  2.178x10

Z
18
Z 
J 2 
n 
2
is the nuclear charge, n is the energy
level.
NIELS BOHR (1885-1962)
 The
most important equation to come
from Bohr’s model is the expression for
the energy levels available to the
electron in the hydrogen atom.

E  2.178x10
 the
18
Z 
J 2 
n 
2
negative sign calculates a lower
energy closer to the atom, not the
radiation of negative energy.
EXAMPLE
What is the change in energy if an
electron in level 6 (excited state)
returns to level 1 (ground state)
in a hydrogen atom?
 ni=6; nf=1; Z=1 (hydrogen
nucleus contains a single proton)
EXAMPLE
What is the change in energy if an
electron in level 6 (excited state)
returns to level 1 (ground state)
in a hydrogen atom?
2


1
18
E6  2.178x10 J  2   6.050x10 20 J
6 
2


1
18
E1  2.178x10 J  2   2.178x10 18 J
1 
EXAMPLE
ΔE=Ef – Ei = E1 – E6=-2.117 x 10-18J
The negative sign for the change in energy
indicates that the atom has lost energy
and is now more stable. This loss of
energy produces a photon.
EXAMPLE
What is the corresponding wavelength
for the energy produced from the
electron jump?
E = -2.117 x 10-18J

 c
E  hv  h  
 
9.383x10-8 m
BOHR MODEL
CONCLUSIONS
 The
model correctly fits the quantized energy
levels of the hydrogen atom and postulates
only certain allowed circular orbits for the
electrons.
 As the electron becomes more tightly bound,
its energy becomes more negative relative to
the zero-energy reference state. As the
electron is brought closer to the nucleus,
energy is released from the system.
BOHR MODEL
CONCLUSIONS
E  E final  Einitial 


1
1
18
2.178x10 J  2  2 
 n final ninitial 
BOHR MODEL
CONCLUSIONS
 The
energy levels calculated by Bohr
closely agreed with the values
obtained from the hydrogen emission
spectrum but does not apply well to
other atoms. The Bohr’s model is
fundamentally incorrect but is very
important historically because it
paved the way for our current theory
of atomic structure.
THE QUANTUM
MECHANICAL
MODEL OF THE
ATOM
QUANTUM MECHANICS
Quantum Mechanics or Wave
Mechanics were developed by three
physicists: Heisenberg, de Broglie,
and Schrodinger.
 Emphasis was given to the wave
properties of the electron.
 The electron bound to the nucleus
behaves similar to a standing wave.
QUANTUM MECHANICS
 Like
a standing
wave, electrons can
travel in patterns
that allow for a
common node. In
other words, wave
patterns around the
nucleus must be in
whole number wave
patterns. But their
exact movement is
not known.
HEISENBERG
UNCERTAINTY PRINCIPLE
There
is a fundamental limitation
to just how precisely we can know
both the position and momentum of
a particle at a given time. This
limitation is small for large
particles but substantial for
electrons.
h
x   mv  
4
PROBABILITY
DISTRIBUTION
A probability
distribution is used
to indicate the
probability of finding
an electron in a
specific position.
 Electron density
map
 Radial probability
distribution
PROBABILITY
DISTRIBUTION
For the hydrogen 1s orbital, the maximum
radial probability occurs at a distance of
5.29x10-2nm or .529Å from the nucleus.
This is the exact radius of the innermost
orbit calculated in the Bohr Model.
 The definition most often used by
chemists to describe the size of the
hydrogen 1s orbital is the radius of the
sphere that encloses 90% of the total
electron probability
QUANTUM
NUMBERS
QUANTUM NUMBERS
Each orbital is characterized by a
series of numbers called quantum
numbers, which describe various
properties of an orbital:
1. Principal quantum number (n)has integral values : 1,2,3,4. It
describes the size and energy of the
orbital. Energy Level
QUANTUM NUMBERS
2.
Angular momentum quantum
number (l) – has integral values from
0 to n-1. This is related to shape of the
atomic orbitals. Sublevel
o
o
o
o
o
l =0 is s
l =1 is p
l =2 is d
l =3 is f
l =4 is g
QUANTUM NUMBERS
3.
Magnetic quantum number
(ml) has values between l and – l
, including 0. The value of ml is
related to the orientation of the
orbital in space. Axis designation
QUANTUM NUMBERS
4.
Electron spin quantum
number (ms)- can only have
one of two values, +½, -½.
Electrons can spin in one of two
opposite directions.
QUANTUM NUMBERS
In a given atom no two electrons can
have the same set of four quantum
numbers (n, l, ml , ms). This is called
the Pauli exclusion principle; an
orbital can only hold two electrons,
and they must have opposite spins.
QUANTUM NUMBERS
ORBITAL SHAPES
AND ENERGIES
S ORBITALS
 The
s orbitals
have a
characteristic
spherical shape
and contain areas
of high probability
separated by areas
of zero probability.
These areas are
called nodal
surfaces, or nodes.
S ORBITALS
 The
number of
nodes increases as
n increases. The
number of nodes =
n - 1.
P ORBITALS
P
orbitals each have two lobes separated by a
node at the nucleus. The p orbitals are
labeled according to the axis of the xyz
coordinate system along which the lobes lie.
P ORBITALS
Cross section of electron probability of a p
orbital
D ORBITALS
 The
five d orbitals first occur in energy
level 3. They have two fundamental
shapes. Four of the orbitals (dxz, dyz, dxy,
and dx2-y2) have four lobes centered in the
plane indicated in the orbital label. dx2-y2
lie along the x and y axes and dxy lie
between the axes. The fifth orbital dz2 has
a unique shape with two lobes along the z
axis and a belt centered in the xy plane.
D ORBITALS
F ORBITALS
 The
f orbitals first occur in level 4 and
have shapes more complex than those of
the d orbitals. These orbitals are not
involved in the bonding in any of the
compounds that we will consider.
F ORBITALS
ORBITAL ENERGIES
 For
the hydrogen atom, the energy of
a particular orbital is determined by
its value of n. Thus all orbitals with
the same value of n have the same
energy – they are said to be
degenerate.
POLYELECTRONI
C ATOMS
POLYELECTRONIC ATOMS
Polyelectronic atoms are atoms with more
than one electron. To look at these atoms,
three energy contributions must be
considered:
 Kinetic energy of the electrons as they
move around the nucleus.
 The potential energy of attraction between
the nucleus and the electrons.
 The potential energy of repulsion between
the two electrons.
POLYELECTRONIC ATOMS
Since electron pathways are unknown,
dealing with the repulsions between
electrons cannot be calculated
exactly.
 This is called the electron correlation
problem.
POLYELECTRONIC ATOMS
The electron correlation problem
occurs with all polyelectronic atoms.
To deal with this, we assume each
electron is moving in a field of
charge that is the net result of the
nuclear attraction and the average
repulsions of all the other electrons.
In other words,…..
POLYELECTRONIC ATOMS
A valence electron is attracted to the
highly charged nucleus and still
repelled by the other ‘inner’
electrons. The net effect is that the
electron is not bound nearly as
tightly to the nucleus as it would be
if it were alone.
 This is a screened or shielded affect.
POLYELECTRONIC ATOMS
Because of this shielded affect.
orbitals within a principal energy
level do not have the same energy
(degenerate). Sublevels vary in
energy within a principal quantum
level.
s<p<d<f
POLYELECTRONIC ATOMS
Hydrogen
vs.
Polyelectronic
HISTORY OF THE
PERIODIC TABLE
EARLY GREEKS
 Earth
 Air
 Fire
 Water
DOBEREINER (1780-1849)
Johann Dobereiner was the first
chemist to recognize patterns and
found several groups of three
elements that have similar
properties.
 chlorine, bromine and iodine
 called triads.
NEWLANDS
 John
Newlands suggested that
elements should be arranged in
octaves, based on the idea that
certain properties seemed to repeat
for every eighth element in a way
similar to the musical scale.
MEYER AND MENDELEEV
The present form of the periodic table
was conceived independently by two
chemists: Meyer and Mendeleev.
Usually Mendeleev is given most of
the credit, because it was he who
emphasized how useful the table
could be in predicting the existence
and properties of still unknown
elements.
MEYER AND MENDELEEV
 In
1872 when Mendeleev first
published his table, the elements
gallium, scandium, and germanium
were unknown. Mendeleev correctly
predicted the existence and
properties of these elements from the
gaps in his periodic table.
Mendeleev also corrected the atomic
masses of several elements.
MENDELEEV’S TABLE
THE AUFBAU
PRINCIPLE AND
THE PERIODIC
TABLE
THREE RULES FOR
ORBITAL CONFIGURATION
 Aufbau
principle – As protons are added,
so are electrons, and fill in orbitals in
order of energy levels.
 Pauli Exclusion – Two electrons with
opposite spins can occupy an orbital.
 Hund’s rule – The lowest energy
configuration for an atom is the one with
one unpaired electrons in each
degenerate orbital. (Electrons don’t like
roommates)
VALENCE ELECTRONS
 Valence
electrons are the electrons
in the outermost principal quantum
level of an atom. These are the most
important electrons because they are
involved in bonding.
 The inner electrons are known as
core electrons.
VALENCE ELECTRONS
 The
elements in the same group
have the same valence electron
configuration. Elements with the
same valence electron configuration
show similar chemical behavior.
TRANSITION METALS
Transition metals have electron
configurations that fill in the order of 4s
before 3d. Copper and Chromium have
a configuration that is observed
different than what is expected.
 Expected: Cr: 1s22s23s23p64s23d4

Observed: 1s22s23s23p64s13d5
 Expected:

Cu:1s22s23s23p64s23d9
Observed: 1s22s23s23p64s13d10
TRANSITION METALS
ADDITIONAL ORBITAL
RULES
 The
(n+1)s orbital always fills before the nd
orbitals. The s orbitals fill prior to the d
orbitals due to the vicinity of the nucleus.
 After lanthanum, which has the
configuration of [Xe] 6s25d1, a group of 14
elements called the lanthanide series, or the
lanthanides occurs. This seris of elements
corresponds to the filling of the seven 4f
orbitals.
ADDITIONAL ORBITAL
RULES
ADDITIONAL ORBITAL
RULES
 After
actinium, a group of 14 elements
called the actinide series or actinides
occurs.
 The groups 1A, 2A, 3A…, the group
numbers indicate the total number of
valence electrons for the atoms in these
groups.
ADDITIONAL ORBITAL
RULES
 After
actinium, a group of 14 elements
called the actinide series or actinides
occurs.
 The groups 1A, 2A, 3A…, the group
numbers indicate the total number of
valence electrons for the atoms in these
groups.
ADDITIONAL ORBITAL
RULES
PERIODIC TRENDS
IN ATOMIC
PROPERTIES
IONIZATION ENERGY
Ionization energy is the energy
required to remove an electron from
a gaseous atom or ion when the atom
or ion is assumed to be in its ground
state:

X(g)
X+(g) + e-
IONIZATION ENERGY
 It
is always the highest-energy
electron (the one bound least tightly)
that is removed first. The first
ionization energy (I1) is the energy
required to remove that first
electron. The second ionization
energy (I2) is considerably larger.
IONIZATION ENERGY
 The
first electron is removed from a
neutral atom, the second from a +1
cation. The increase in positive charge
binds the electrons more firmly and the
ionization energy increases. The trend
continues for consecutive electrons
removed.
 Core electrons are always held tighter
than valence.
IONIZATION ENERGY
 First
ionization energy increases from
left to right across a period.
 First ionization energy decreases in
going down a group.
IONIZATION ENERGY
ELECTRON AFFINITY
Electron Affinity is the change in
energy change associated with
the addition of an electron to a
gaseous atom:
 X(g)
+ e-
X-(g)
ELECTRON AFFINITY
 If
the addition of the electron is
exothermic the corresponding value for
electron affinity will carry a negative
sign.
The more negative the energy, the
greater the quantity of energy
released.
ELECTRON AFFINITY
 Electron
affinities generally become
more negative from left to right across
a period and becomes more positive
down a group.
 As with Ionization energy. Some
exceptions occur due to repulsions
and electron configuration.
ELECTRON AFFINITY
ELECTRON AFFINITY
ATOMIC RADIUS
Atomic radii are measured by the distances
between atoms in chemical compounds.
 Covalent atomic radii are assumed to be
half the distance between atoms in
covalent bonds.
 For metallic atoms, the metallic radii are
obtained from half the distance between
metal atoms in a solid metal crystal
ATOMIC RADIUS
 Atomic
radii decrease in going from left to
right across a period because of
increasing nuclear charge and decreasing
shielding.
 Atomic radius increases down a group,
because of the increases in the orbitals
sizes associated with principal quantum
numbers.
ATOMIC RADIUS
ATOMIC RADIUS
THE PROPERTIES
OF A GROUP: THE
ALKALI METALS
INFORMATION AND THE
PERIODIC TABLE
 It
is the number and type of valence
electrons that primarily determine
an atom’s chemistry
 The organization of the period table
allows the prediction of electron
configuration without memorization.
INFORMATION AND THE
PERIODIC TABLE
 Groups
on the periodic table have
specialized names: Alkali metals,
Alkaline earth metals, Halogens, …etc.
 The most basic division of elements in
the periodic table is into metals and nonmetals. This division affects chemical
properties.

Metals tend to give up electrons and have
low ionization energies. The opposite is
true for non-metals.
INFORMATION AND THE
PERIODIC TABLE
INFORMATION AND THE
PERIODIC TABLE
 Metalloids
are elements along the
division line and exhibit both metallic
and nonmetallic properties under certain
circumstances. These elements are
sometimes called semimetals.
THE ALKALI METALS
 Lithium,
sodium, potassium, rubidium,
cesium, and francium are the most
chemically reactive of the metals.
Hydrogen is found in group 1 but
behaves as a nonmetal because its very
small and the electron is bound tightly
to the nucleus.
THE ALKALI METALS
 Going
down the group the first
ionization energy decreases and the
atomic radius increases. The overall
density increases due to the increase of
atomic mass relative to atomic size
(therefore more mass per unit volume).
THE ALKALI METALS
 There
is a smooth decrease in melting
point and boiling points in Group 1 that
is not typical for other groups.
 The most important chemical property
of Group 1 is its ability to lose its
valence electrons. Group 1 are very
reactive.
THE ALKALI METALS
 Hydration
energy of an ion represents the
change in energy that occurs when water
molecules attach to the metal cation.
 The hydration energy is greatest with Li+
because it has the most charge density (charge
per unit volume). This means that polar
water molecules are more strongly attracted to
the small Li+ ions
 The order of reducing abilities in an aqueous
reaction is Li > K > Na
THE ALKALI METALS
THE ALKALI METALS
THE END