Transcript The Structure of Atoms

The Structure of Atoms
Chapter 5
Why Study This Chapter First?
• A basic understanding of the atom is required to
comprehend many fundamental concepts that will be
discussed in this chemistry course.
• The behavior of atoms is a result of its fundamental
particles and how they are arranged.
• You will be able to answer
– Why atoms form compounds?
– Why atoms/elements have similar chemical and physical
properties?
– Why atoms combine in certain ratios to form compounds?
The Fundamental Particles in the
Atom
Particle
Mass (amu*)
Electron (e-)
0.00054858 amu
Charge (relative
scale)
1-
Proton (p or p+)
1.0073 amu
1+
Neutron (n or n0) 1.0087 amu
1 amu = 1.660  10-24 g
none
Discovery and Properties of
Electrons
• Humphry Davy and Michael Faraday
performed experiments in the early 1800s
illustrating the existence of charged
particles.
– Termed as ‘electrons’
– These charged particles were suggested to be
responsible for holding substances together.
Discovery and Properties of
Electrons
• J.J. Thomson (1897) used cathode-ray tubes to study these
‘electrons’ in more detail. He was able to determine the
ratio of charge to mass for the electron.
– e/m = 1.75882  108 coulomb/gram
Original instrument used by J.J. Thomson
Discovery and Properties of the
Electron
• Robert Millikan determined the
mass and charge of an electron
using the ‘Millikan oil-drop
experiment.
– Measured charges were all
integrals of the same number
(homework question).
– Charge (e-) = 1.60218  10-19 C
– The mass of the electron was
determined using the charge-tomass ratio determined by
Thomson.
Mass (e-) = 9.10940  10-28 grams
First Attempt at the Structure of
the Atom (Incorrect)
• In 1886, Eugen Goldstein
discovered positively
charged ions by using
modified cathode-ray tubes.
– Termed as ‘protons’
• The scientific community
proposed that these positive
and negative charges were
mixed together.
– Plum pudding model
The Structure of the Atom
Revisited (More correct)
• In 1909, Ernest Rutherford illustrated that the
plum pudding mode for the atom was incorrect.
Rutherford used alpha particles (He2+) to
establish the arrangements of protons and
electrons in the atom.
– Scattering of alpha particles through a thin piece of
gold foil was observed on a scintillation screen. If
the plum pudding model were correct, all the alpha
particles would be deflected by very small angles.
What did he observe?
The Structure of the Atom
Revisited (More correct)
Important observations from the Rutherford
experiment.
• Most of the particles passed through the foil
with no or little deflection.
• Surprisingly, a few were deflected at very
high angles, and a few came directly back to
the source (i.e. starting point).
The Structure of the Atom
• According to the Rutherford model each atom
consists of a very dense region of positive
charge surrounded by a diffuse region of
negative charge.
– The positive region contains most of the mass.
– The diameter of the positive region is 1/10,000 of
the atom.
• The density of positive region is ~1015 g/mL
– The atom is primarily empty space.
More Atom Information
• H.G.J. Moseley used X-rays to propose that each
element differs from the preceding element by one
positive charge (i.e. proton) in its nucleus.
– This led to the arrangement of the periodic table (according
to proton number). The number of protons in the nucleus is
equal to the atomic number.
• Understand the periodic table and the atomic number.
• James Chadwick discovered the presence of neutrons.
– All elements except hydrogen contain neutrons.
The Atom in Summary
• Atoms contain small, dense nuclei surrounded by
clouds of electrons.
– The nuclei contains protons and neutrons.
• Electrons are at relatively great distances from the
nuclei.
– Nuclear diameters are typically 10-5 nanometers while
atomic diameters are 10-1 nanometers.
• Basketball and six miles
• All atoms contain protons. All atoms except hydrogen
contain neutrons. Look at periodic table.
Mass Number and Isotopes
• Mass number is the sum of the ____ and ____ in
an atom’s nucleus. The number of ____
determined the atomic number. This number also
determines the identity of the atom.
• Isotopes are atoms of the same element with
different number of neutrons.
– Isotopes of he same element contain the same number
of ______.
– Nearly all elements have more than one isotope.
• hydrogen-1, hydrogen-2 (deuterium), and hydrogen-3 (tritium)
• chlorine-35 and chlorine-37
Using Nuclide Symbols to
Indicate Isotopes
• The isotopes is represented in the format, AZ E .
– A is the mass number, Z is the atomic number, and E is
the element.
• Write the chlorine-35 and chlorine-37 with nuclide
symbols. How many protons, neutrons and
electrons are in each isotope?
• Write plutonium-244 with nuclide symbols. What
is the number of protons, neutrons, and electrons in
this isotope?
Atomic Weight
• Atomic weight scale is based on the carbon-12
isotope.
– One amu (atomic mass unit) is exactly equal to 1/12 of
a carbon-12 atom.
• Atomic weight is a weighted average of all _____
occurring isotopes. This is the reason that most
numbers in the periodic table are significantly less
or more than a whole number.
– To calculate the atomic weight in amu multiply each
isotope by the natural abundance and add the terms.
Atomic Weight Calculations
• The two naturally occurring isotopes for chlorine are
chlorine-35 and chlorine-37. Isotopic masses for the two
isotopes were measured at 34.9689 and 36.9659 amu. The
natural abundances are 75.77% and 24.23%, respectively.
Calculate the atomic weight.
• The naturally occurring isotopes for strontium are Sr-84, Sr86, Sr-87, and Sr-88. The natural abundances of these
isotopes are 0.56%, 9.86%, 7.00%, and 82.58%. What is the
atomic weight?
• The atomic weight of gallium is 69.72 amu. The two
naturally occurring isotopes are Ga-69 and Ga-71. What is
the natural abundance of each isotope?
Look at webelements.com
Mass Spectrometers
• Mass spectrometers
are instruments that
measure the mass-tocharge ratio of charged
particles. The extent
of deflection of ions in
a magnetic field
depends on the massto-charge ratio.
Mass Spectrometers
• Mass spectrometers are able to separate and determine the
isotopic abundances of a particular element.
– Ions are detected in order of mass-to-charge ratio (m/q). The ions
with the smallest m/q are detected first, and the ions with greatest
m/q are detected last. The intensity or heights of the peaks in a
spectrum reflect the natural abundance.
• The naturally occurring isotopes of B are boron-10 and
boron-11. The natural abundances of the isotopes are 20%
and 80%. How would the spectrum for boron appear? (mass
spectra plotter)
• How would the mass spectrum for Ge appear?
http://www.webelements.com/webelements/elements/text/Ge
/isot.html
The Nature of Light
• Much of the information concerning the
arrangement of electrons in the atom has been
based on the light emitted (given off) or
absorbed by atom.
– Activate a light stick; Why are only certain color
emitted? This pertains to most objects that exhibit
a specific color (neon signs)
• Light has both ______ properties and ______
properties.
Wave-like Properties of Light
• Light is electromagnetic radiation that exhibits
oscillating wave-like behavior. All waves are in
motion.
– Wavelength () is the distance between two identical points
of the wave.
– Frequency () is the number of wave crests passing a point
per second.
– The product of the wavelength and frequency is equal to the
speed of light, c.
=c
c=3.00  108 meters/second
Wave-like Properties of Light
• The visible region of the
electromagnetic spectrum ranges
from about 4.0  10-7 m to 7.5 
10-7 m.
– A very tiny slice of the spectrum.
• Demo: Light through a prism.
White light contains many
wavelengths (all visible).
– The spectrum through a prism is
called a _____ _____ because it
contains a continuous range of
wavelengths.
Wave-like Properties of Light
• Some problems
– What is the wavelength, in angstroms, of
electromagnetic radiation having a frequency of
3.00  1020 Hz (cycles per second)?
– What is the corresponding frequency of
electromagnetic radiation that has a wavelength
of 15 nanometers?
1 nanometer = 1.0  10-9 meters
Particle-like Properties of Light
• Light can also exhibit particle-like properties
under certain conditions. The particles that
compose light are referred to as _______. Each
_____ has its own frequency, and, therefore, a
unique energy.
– Energy of a photon can be expressed as E = h or
hc/.
• h (Planck’s constant) = 6.6260755  10-34 Joulessecond
Energy is _____ proportional to the frequency and
_____ proportional to the wavelength.
Particle-like Properties of Light
• Some problems
– My cordless phone operates at 900 MHz. Calculate
the energy of a photon at this frequency.
– Red light has a wavelength near 7500 Å. Calculate
the energy of a photon of red light.
Increasing the intensity of light only increases the
number (or flux) of photons.
For a specific color to be observed the photons
emitted have to posses a specific energy. How did
this relate to the atom?
Quantization of the Electron in
the Atom
• The Bohr Atom
– When electric current is passed through hydrogen gas only
several emission lines are observed in a spectrum (only
certain colors).
• DEMO: Emission spectrum of hydrogen
– For the hydrogen atom, Balmer and Rydberg illustrated that
the observed wavelengths are related to the following
expression:
 1
1
1 
 R  2  2 

 n1 n 2 
Where R has a value of 1.097  107 m-1 (Rydberg constant).
The n’s are positive integers, with n2 greater than n1.
Quantization of the Electron in
the Atom
• The Bohr atom
– Niels Bohr proposed that the electron energy is
quantized; meaning that only certain values of
energy are allowed/possible for the electron.
Energy will be absorbed or emitted in specific
amounts as the electron moves from one energy
level to another.
– Even though the Bohr atom model has flaws, it
illustrates the quantization of the electron.
• Look at Fig. 5-16 in book and discuss.
Quantization of the Electron in
the Atom
• What does this mean in summary?
– An atom possesses a number of energy levels in
which electrons can exist.
• The lowest possible energy level is called the
ground state (n=1). The higher energy levels are
called excited excited states.
– In order for an electron to move from a lower
energy level to a higher energy level, it must
____ an amount of energy equal to the
difference between the two levels.
Quantization of the Electron in
the Atom
• What does this mean in summary?
– The electron may stay located in an energy level
without emitting or absorbing energy.
– Movement of an electron to a lower energy level
will emit energy equal to the difference between
the two levels.
E 2  E 1  E  h  
hc

Conservation of energy is observed in this process.
Atomic Emission
• Emission spectrum –
certain wavelengths of
light are emitted from a
substance
– Light can be separated by a
prism or grating. This is
called a bright-line
spectrum.
– The lines are caused by
electron transitions from
higher energy levels to
lower levels.
DEMO: Gratings on a few
gases.
Atomic Absorption
• Absorption spectrum certain wavelengths of light
are absorbed by a
substance.
– These wavelengths
correspond to electron being
promoted to higher levels.
The wavelengths are
removed.
– The wavelengths that are
absorbed are usually given
off in the emission spectrum.
Electrons Can Behave Like Waves
• Louis de Broglie illustrated that small particles,
such as electrons, can exhibit wave-like
properties. The wavelength can be expressed as

=h/m

h(Planck’s constant) = 6.626  10-342 Joulesec
kg  m
In units, J  sec 2
Therefore, Planck’s constant can also be written as
h  6 .626  10 34
kg  meter
sec
2
Electrons Can Behave Like Waves
• Some problems
– What is the wavelength of a proton moving at 2.50 
107 m/s?
– What would be the wavelength of a car (mass 1500 kg)
moving at a velocity of 15.2 m/s?
• For electrons in atoms, classical interpretations do
not work. The electron must be treated as having
wave-like properties. This is the basis for
quantum mechanics.
DEMO: Diffraction limit related to the optical
microscope versus the electron microscope
(calculation).
Basic Ideas of Quantum Mechanics
• Atoms or molecules can only exist in certain
energy states.
• Atoms or molecules emit or absorb energy
when they change their energy state.
– Equal to the difference between the two energy
states.
• The allowed energy states of atoms or
molecules can be described by a set of
quantum numbers.
Quantum Mechanics and the Atom
• The electron in the atom can be treated as a
standing wave. Only certain “waves” are
allowed/permitted for an electron around the
nucleus. Each wave corresponds to an energy
state.
• Schrödinger developed an equation that utilizes
standing waves to calculate the energies of the
electron in the hydrogen atom.
  2  2  2 
 2  2  2  2   V   E
8 m  x
y
z 
h
Quantum Mechanics and the Atom
• The equation can be solved for the hydrogen atom
to produce the values that are allowed for the
electron in the hydrogen atom. Each solution (or set
of values) can be described by a set of quantum
numbers. More complex equations are required for
atoms with many electrons.
• Four quantum are necessary to describe the energy
state of an electron in the atom.
– Principal quantum number, angular momentum quantum
number, magnetic quantum number, and spin quantum
number
Quantum Numbers (#1)
• Principal quantum number, n
– The shell number or main energy level. It may
be any positive integer.
– Primary indicator of energy
– Distance from the nucleus increases with n
• The further the electron from the nucleus the greater
the energy. Therefore, energy increases with n.
Quantum Numbers (#2)
• Angular momentum quantum number, l
– Corresponds to the subshell label
– Designates the shape of the region of space that
the electron occupies (orbital)
– The number of subshells (or values of l) is
equal to n, the principal quantum number
• l = 0, 1, 2, … (n-1)
What are the possible values of l if n is equal to 3?
• s=0, p=1, d=2, f=3
Quantum Numbers (#3)
• Magnetic quantum number, ml
– Specifies the orbital in a subshell that the electron
is assigned
– This number specifies the orientation of that
orbital in space
– The number of orbitals in a subshell increases
with l
• ml = (-l), …, 0, …., (+l)
• Number of individual orbitals = 2l+1
What are the possible values of ml if l = 2 and 3?
Quantum Numbers (#4)
• Spin quantum number, ms
– Refers to the spin of an electron
– For every value of ml (or orbital), ms can have a value of
+1/2 or –1/2
– Each orbital can accommodate only ___ electrons?
What are the possible values of ms if ml equals 2? How about
if ml equal 55?
Look at Table 5-4 for allowed values of quantum numbers.
If n=3, what are the allowed quantum numbers? These values
are the possible energy states for an electron in the 3rd shell.
Atomic Orbitals
• An atomic orbital is a region of space in
which the probability of finding an electron
around the nucleus is high (usually > 75%).
• The electrons must occupy an orbital.
Electrons in orbitals do not orbit or circle
around the nucleus.
– Orbitals contain diffuse clouds of electrons that
are at significant distances away from the
nucleus.
Atomic Orbitals
• Three quantum numbers are used to specify the
orbital in which the electron is located
– n, l, and ml
– The fourth quantum number, ms, is used to specify the
spin of the electron in the orbital.
• The shell number, n, of an atomic orbital has a
value of 1, 2, 3, …..
– Indicates the energy of the electron in that orbital
– Each shell is further away from the nucleus (higher
energy)
– Each shell has a capacity for 2n2 electron
Atomic Orbitals and Subshells
• s subshell
– For the s subshell, l = 0.
• Every shell has an s subshell that contains ___
orbital (ml =0). How many electrons can each
subshell hold? What are the values?
• These subshells are designated as 1s, 2s, 3s, etc.
How many subshells for n=1? n=3?
• The s orbital is spherically symmetric (illustrate)
• The diameter of s orbital increases with shell
number.
– The energy of an electron in the 4s orbital is higher than
the energy of an electron in the 2s orbital. Why?
Atomic Orbitals and Subshells
• s subshell
– There are regions of space where the electrons
are not allowed to be for s orbitals having a
shell number greater than 1. These regions are
called _____.
• DEMO: Show with orbital software
• The number of nodes increases with shell number.
Which has more nodes, 3s or 1s?
Atomic Orbitals and Subshells
• p subshell
– All shells equal to 2 or higher possess a p subshell
– For any p subshell, l = 1
– The three p orbitals are mutually perpendicular and shaped
like dumbbells
• The subscript indicates the Cartesian scale (px, py, and pz)
• ml = __ __ __
– With increasing n, the p subshells increase in energy and the
number of nodes. The distance from the nucleus also
increases.
DEMO: Illustrate the 2p and 4p
Atomic Orbitals and Subshells
• d subshells
– All shells equal to 3 or higher possess a d subshell
– For any d subshell, l = 2
– Each subshell consists of five d orbitals
• ml = __ __ __ __ __
• There are four clover-leaf-shaped orbitals and 1 orbital
shaped like a peanut with a donut around the center
DEMO: Illustrate orbitals with software
Atomic Orbitals and Subshells
• f subshell
– All shells equal to 4 or higher possess and f subshell
– For any f subshell, l = 3
– Each subshell consists of seven orbitals
• ml = __ __ __ __ __ __ __
– Shapes are complicated
DEMO: Illustrate a 4f orbital with software
Atomic Orbitals and Subshells
• Recall that a particular orbital is described by a
unique set of n, l, and ml values.
– Each orbital can hold two electrons. These two
electrons are described by the ___ quantum number
• Values equal __ and __
• The electrons are spin-paired (opposite spin). One
electron has an up spin, , and one electron has a down
spin, . Spin pairing produces lower energy.
Atomic Orbitals and Subshells
• Summarizing
– Number of subshells in a shell is equal to __
– Number of orbitals in a shell is equal to __
– Number of electrons in a shell is equal to __
Relation of atomic orbitals is illustrated nicely in
Table 5-4 (also on demo. CD, screen 7-12)
Handouts
Problems
Electron Configurations –
Electron Buildup
• For atoms with more than one electron, the
Schröndinger equation becomes much more
complicated. There are not exact solutions. For
these atoms, the electrons are approximated to
occupy atomic orbitals similar to those
developed for the hydrogen atom. This is
called the orbital approximation. With
appropriate modifications, this can be a very
good approximation.
Electron Configurations –
Electron Buildup
• How electrons are placed into individual orbitals is
best understood by examining the electron
configurations as a function of increasing atomic
number.
• Aufbau Principle – electrons will be added to orbitals
in such a way to acquire the lowest energy for the
atom.
– Orbitals increase in energy with increasing, n, or shell
number.
– Orbitals increase in energy with increasing value of l or
subshell.
• In the 2nd shell, the 2s would fill before the 2p
Electron Configurations –
Electron Buildup
• General guidelines:
– Electrons are assigned in orbitals of increasing value
of (n + l) or (shell + subshell).
– For subshells with the same value of (n+l), electrons
are first assigned to the subshell with lower n (shell
number).
Of the two subshells, 3p and 3d, which will fill first?
How about 4s and 3d?
What about 4s and 3p?
Electron Configurations –
Electron Buildup
• Paul Exclusion Principle – No two electrons in an
atom may have the same set of four quantum
numbers.
• Shell 1 – the first period or row in the periodic
table
– How many subshells? How many orbitals?
– Look at Figure 5-28
– The electron configuration for H and He
Electron Configurations –
Electron Buildup
•
Shell 2 – the 2nd row or period in the periodic table
–
–
There are __ subshells and __ orbitals in the 2nd shell
Li and Be, What happens with B?
•
–
Write with shortened electron configurations
The electron configurations for C and N.
•
The three p orbitals are degenerate. What does this mean?
Hund’s Rule: Electrons occupy all the orbitals of a given
subshell singly before pairing (parallel spins).
– The electron configuration for O.
Paramagnetic – substances with unpaired electrons are weakly
attracted into magnetic fields. Diamagnetic substances have
paired electrons are weakly repelled by magnetic fields.
Electron Configurations –
Electron Buildup
• Shell 3 – the 3rd row or period in the periodic table.
– How many subshells and orbitals?
– The full and shortened electron configuration for Cl
– What about the 3d subshell? Why doesn’t it fill before the
4s subshell?
• Shell 4 – the 4th row or period in the periodic table
starts to fill.
– How many subshells and orbitals?
– The electron configurations for Cr
– Be sure to fill the subshells according to increasing energy
(careful!!, Figure 5-28) I have handouts on my website.
Electron Configurations –
Electron Buildup
• Not all electron configurations follow the expected
trends. Some subshells are very close in energy. When
electrons start to fill these orbitals, the subshell can shift
in energy. The slight shift in energy can change the
electron buildup.
– Cr is [Ar]4s23d5 not the expected [Ar]4s23d4. The former
electron configuration has slightly lower energy since the
electrons are all unpaired.
– Cu has an electron configuration of [Ar]4s13d10
There appears to be stability in filled or half-filled subshells (d
and f).
Electron Configuration in the
Periodic Table
• The periodic table is arranged according to
how electrons fill specific orbitals in shells
and subshells. The periodic table is divided
into columns or groups.
• In the A groups the s or p orbitals are being
filled
– In columns 1 and 2, the s orbitals are being filled
(the s block)
– In columns 3-8, the p orbitals are being filled (the
p block)
Electron Configuration in the
Periodic Table
• In the B groups (center) the d orbitals are
being filled (the d block)
• In the rows listed below, the f orbitals are
being filled (the f block).
Examine Figure 5-31 and understand well.
Obtain my handouts if needed on my website.
Electron Configuration in the
Periodic Table
• Elements in the same group or column have
similar electron configurations. This the the
reason why they have similar chemical and
physical properties.
– Na and K; Ne and Ar
Interestingly, the periodic table was developed
before knowledge of electron configurations.
Table 5.5 and Appendix B exhibit the exact electron
configurations for the elements.