Download Chapter 3: Atomic Structure

Chapter 3: Atomic Structure
3.1 Early Models of the Atom
Since ancient times, humans have used chemical changes to their advantage.
 ores to produce metals
 embalming fluids in Egypt
 the Greeks were the first to try explaining chemical changes, they proposed that all
matter was made of four fundamental substances (elements): fire, earth, water and air.
 the next 2000 years were dominated by alchemy (see insert).
 Democritus philosophized that a pure substance could only be subdivided so many
times and then it was “atomos” or “indivisible”
An atom is the smallest particle of an element that retains the chemical identity of that element.
In the 17th and 18th centuries scientist started to base their ideas on observation and experiments
such as:
1. Law of Conservation of Mass (Antoine Lavoisier, 1789)
Matter is neither created nor destroyed under normal chemical reactions
2. Law of Definite Proportion (Joseph Proust, 1799)
Atom ratio is fixed, so mass must be constant.
3. Law of Multiple Proportions (John Dalton, 1803)
Applies where two elements, A and B, form more than one compound.
Dalton concluded from these three laws that the properties of matter could be explained in
terms of atoms. Dalton’s ATOMIC THEORY OF MATTER was based on the following postulates
(see insert).
1. Each element is composed of extremely small particles called atoms.
2. All atoms of a given element are identical, but they differ from those of any other element.
3. Atoms are neither created nor destroyed in a chemical reaction.
4. A given compound always has the same relative numbers and kinds of atoms.
Some of his postulates now have exceptions.
All matter can be broken down chemically into about 100 elements or 100 different kinds of atoms.
These atoms combine to form each of the vast substances that make up the world around us. Just like
the English language has thousands of words composed of 26 letters. These elements (atoms) vary in
abundance (see transparencies).
Scanning tunneling microscope will provide a blurred picture of an atom. It reveals very little about the
atom itself. Thus one must understand the microscopic world of chemistry, even sub-subatomic
particles (quarks and leptons).
Homework
Write the letter of the chemist who proposed each of the ideas listed below. Each letter maybe used
once, more than once, or not at all.
a. Democritus
b. Lavoisier
c. Proust
d. Dalton
_____ 1. Matter is neither created nor destroyed in chemical reactions.
_____ 2. A given compound always has the same relative numbers and kinds of atoms.
_____ 3. All atoms of a given element are identical, but they differ from those of any other element.
_____ 4. A given compound always contains the same elements in the same proportions by mass
_____5. Each element is composed of extremely small particles called atoms.
_____ 6. All matter is composed of tiny indivisible particles.
_____ 7. Atoms are neither created nor destroyed in chemical reactions.
If the statement is true, write “true.” If it is false, change the underlined word or words to make the
statement true.
__________ 8. There are about 5000 elements, which combine to form the vast number of different
substances in the world around us.
__________ 9. It is possible to actually “see” atoms using a scanning tunneling microscope.
__________ 10. The submicroscopic world of the atom includes exotic particles called quarks
gluons.
__________ 11. Like other well-known Greek philosophers of the time, Aristotle agreed with
Democritus’ ideas about atoms.
__________ 12. The study of atoms has lead to technological advances such as television and
computers.
and
3.2 Discovering Atomic Structure
In 1839, Michael Faraday suggested that the structure of atoms was somehow related to electricity.
The word electricity comes from the Greek word, “elektron” which means amber. Amber when rubbed
creates static electricity. DEMO
Electron
J.J. Thomson, 1897, discovered the electron by using a gas discharge tube. He proved that the atom of
any element can be made to emit tiny negative particles. (He knew they were negative because they
were attracted to the + plate and repelled by the negative plate.) Thomson named these negative
particles electrons. PLUM PUDDING MODEL of the atom.
In 1909 Robert Millikan obtained the first accurate measurement of an electron’s charge. He used an
experiment which became to be known as “MILLIKAN’S OIL DROP EXPT.” By atomizing the oil
droplets and using x-rays to transfer electrons to these oil droplets and by adjusting the electric force to
overcome the gravitational force. Millikan was able to calculate the charge on the droplet. These
charges varied per droplet, however, they were multiples of a small charge of 1.6 x 10-9
Electron 
symbol ehas a unit charge of 1mass that is 1/1837 the mass of a proton or 9.1 x 10-28g
Proton
1911, Ernest Rutherford discovered the nucleus of an atom and the subatomic particle the proton.




GOLD FOIL EXPERIMENT
cathode ray tube which shot alpha particles (+ radiation)
most alpha particles passed straight through however 1 in 8000 were deflected (+ charge
approach + charge)
thus he concluded an atom is mostly empty space with a small + center called a nucleus
the size of the nucleus is about 1/10000 of that of the atom (fly in Yankee stadium)
Nucleus 
center of the atom where the positive charge is concentrated
Contains protons and neutrons
Further experimentation suggested the nucleus consisted of still smaller particles. One is the proton
Proton 
symbol p+
1.67 x 10-24g
proton explains the positive charge of the nucleus
Rutherford realized that protons and electrons did not account for the total mass of an atom. He
predicted the existence of another particle which would have a neutral charge.
NUCLEAR ATOM MODEL of the atom.
Neutron
1932, James Chadwick discovered the missing particle, the neutron.
Neutron 
subatomic particle that is more massive than the proton and electron
1.675 x 10-24g
no charge, neutral
found in the nucleus
Since atoms are not negatively or positively charged. The # of (+) particles must equal the # of (-)
particles. Thus, # e- = # p+
Evolution of the Atom
John Dalton Model  Billard Ball Model
1) J.J. Thomson Model or Plum Pudding Model
 1904
 the electrons are like raisins dispersed in a pudding (positive charge cloud)
 seeds in a watermelon
 blue-berry muffin
2) Rutherford Model or Nuclear Atom Model
 1911
 positive dense center called the nucleus
 rest of atom empty space (electrons reside)
3) Bohr Model or the Planetary Model
 1913
 pictured the atom as a small positive nucleus with electrons orbiting around it in curved
circular pathways
4) Wave Mechanical Model or the Electron Cloud Model
 1920’s
 Louis Victor de Broglie and Werner Schrodinger suggested that because light seems to
behave both as a wave and as a stream of particles, then the electron should exhibit both
of these characteristics
 Orbitals (electron states) are nothing like orbits
 Similar to a cloud or firefly analogy (highest probability)
 Heisenberg Uncertainty Principle  we will never know simultaneously the exact
momentum and position of an electron
Section 3-2 Hwk: Structure of Atoms
In the blank at the left of each word or phrase, write the letter of the expression
on the right that is most closely related.
______ 1. alpha particle
______ 2. anode
______ 3. atomic number
______ 4. cathode
______ 5. Coulomb’s law
______ 6. electron
______ 7. proton
______ 8. isotope
______ 9. mass number
______10. neutron
______11. nucleus
12.
a. the electrode attached to the positive terminal
of a voltage source
b. the electrode attached to the negative terminal
of a voltage source
c. a subatomic particle that has a negative charge
d. an atom’s central region, which is made up of
protons and neutrons
e. a subatomic particle that has a positive charge
and that composes the nucleus of an atom; the
number of these particles determines the identity
of an element.
f. the number of protons that compose the
nucleus of an atom; this number is the same for all
atoms of an element.
g. a subatomic particle that has no charge and
that composes the nucleus of an atom
h. a small, positively charged particle, which
Rutherford directed at thin, gold foil
i. the sum of the number of protons and neutrons
of the nucleus of an atom
j. states that the closer two charges are, the
greater the force between them; in fact, the
force increases by a factor of 4 as the distance is
halved.
k. an atom that has the same number of protons
(atomic number) as other atoms of the same element
but has a different number of neutrons (atomic mass)
Use the appropriate term from the list below to fill in the blanks. Use each term only once.
volume, nucleus, small, alpha, positive, deflected, mass, undeflected
In the Rutherford gold foil experiment, positively charged _____________ particles were directed
at a thin gold foil. It was found that most of the particles passed through the foil ____________.
However, a small number of particles were ______________ , some even backward. These two
observations suggested that most of the ____________ of an atom is empty space but that there
was a central core with a charge that repelled the _____________ particles. This core is a very
____________ part of an atom. It contains most of the ___________ of the atom and is called the
_______________.
Complete the following table.
Isotope
Number of
protons
Number of
electrons
Number of
neutrons
Number of Particles
in nucleus (Mass #)
Symbol for
isotope
Helium-3
Lithium-7
Boron-11
Iodine-125
Potassium-39
Iron-56
Complete the Table
Subatomic
Particle
Location
Charge
Inside nucleus
1+
Mass (amu)
Neutron
1839
1
11.
2.
3.
4.
5.
6.
7.
How many protons and electrons are present in an argon atom?
How many protons and electrons are present in a nitrogen atom?
How many protons and electrons are present in a platinum atom?
What is the name of the element that has atoms that contain 5 protons?
What is the name of the element that has atoms that contain 25 protons?
What is the name of the element that has atoms that contain 82 protons?
Write the chemical symbol for the ion with 12 protons and 10 electrons.
8. Write the chemical symbol for the ion with 74 protons and 68 electrons.
9. Write the chemical symbol for the ion with 33 protons and 36 electrons.
10. How many protons, neutrons, and electrons are present in the
59
28
Ni 2+ ion?
11. How many protons, neutrons, and electrons are present in the
91
40
Zr 4+ ion?
12. Write the complete chemical symbol for the ion with 27 protons, 32 neutrons, and 25 electrons?
13. Write the complete chemical symbol for the ion with 84 protons, 125 neutrons, and 80
electrons?
3.3 Modern Atomic Theory
Electron Configurations
Introduction
In this chapter we will look at atomic structure in more detail. In particular, we will develop a
picture of the electron arrangements in atoms. In this chapter we will see that it is the way the electrons
are arranged in various atoms that accounts for these facts. However, before we examine atomic
structure, we must consider the nature of electromagnetic radiation, which plays a central role in the
study of the atom’s behavior.
Radiant Energy
Much of the understanding of how electrons behave in atoms comes from studies of how light
interacts with matter. To make sense of these studies, you must first consider some aspects of light.
Light travels through space and is a form of radiant energy. It is this energy that causes you to feel hot
when you stand in bright sunlight. But how does light carry energy through space?
Throughout the past 200 years there has been considerable scientific debate over the nature of
light. Until the 1800s, scientists believed that light was a beam of energy moving through space in the
form of waves, much like the waves you see on the surface of a lake. But in the 1900s, scientists
observed another side to the nature of light. They found that in some experiments light behaved like a
stream of extremely tiny, fast-moving particles. Today scientists recognize that light has both the
properties of waves and the properties of particles, and they use both models to describe it. This is
often called the wave-particle duality theory.
Waves
Light waves are electromagnetic waves and light is a form of electromagnetic radiation (A form
of energy called radiant energy that travels at the speed of light with wave-like behavior). X-rays,
gamma rays, and radio waves are other forms of electromagnetic radiation. An electromagnetic wave
consists of electric and magnetic fields oscillating at right angles to each other and to the direction of
motion of the wave. All waves, whether they are water waves or electromagnetic waves, can be
described in terms of four characteristics-amplitude, wavelength, frequency, and speed.
The amplitude of a wave is the height of the wave measured from the origin to its crest, or
peak. The brightness or intensity, of light depends on the amplitude of the light.
The wavelength is the distance between successive crests of the wave. It is the distance that the
wave travels as it completes one full cycle of upward and downward motion. The light, or
electromagnetic radiation, your eyes can see-visible light-has wavelengths in the range of 400 to 750
nanometers.
The frequency of a wave tells how fast the wave oscillates up and down. The frequency of
light is measured by the number of times a light wave completes a cycle of upward and downward
motion in one second. Thus, the unit for frequency is cycles per second. Because it is understood that
cycles are involved, frequency is commonly expressed simply as "per second," which is written as s-1
or 1/s. A cycle per second is also called a hertz (Hz): 1 Hz = 1 s-1.
Light, regardless of its wavelength, moves through space at a constant speed of 3.00 x 108
meters per second (m/s), which is the speed of light. Because light moves at a constant speed, there is
a relationship between its wavelength and its frequency. As Figure 4-3 shows, the shorter the distance
between the crests of the wave, the faster the wave oscillates up and down. That is, the shorter the
wavelength, the greater the frequency. This relationship can be expressed in a simple equation. Using
the symbol λ (the Greek letter lambda) for wavelength, v (the Greek letter nu) for frequency, and c for
the speed of light, the relationship between wavelength and frequency is
C = λν
Electromagnetic Spectrum
You have probably seen a glass prism spread ordinary sun- light into a rainbow of colors. The
same phenomenon happens when sunlight passes through raindrops, creating a familiar sight: a
rainbow. At the inner curve of the rainbow is the color violet and at the outer curve is the color red. In
between are the colors indigo, blue, green, yellow, and orange, each color gradually fading into the
next. This array of colors is called the visible spectrum. The visible spectrum is an example of a
continuous spectrum because one color fades gradually into the next color. The different colors have
different wavelengths (and therefore also different frequencies). Violet has the shortest wavelength
(and the highest frequency). Red has the longest wavelength (and the lowest frequency).
Visible light constitutes a very small portion of the total electromagnetic spectrum. The rest of
the electromagnetic spectrum is invisible to the human eye. You are probably familiar with the words
microwaves, radio waves, and X-rays. These are some other examples of electromagnetic radiation.
Radio waves and microwaves have longer wavelengths than visible light. X-rays on the other hand,
have wavelengths that are much shorter than the wavelengths of visible light.
Using visible light as the point of reference, notice how the wavelengths of the various kinds of
electromagnetic radiation differ. Beyond the red portion of the visible spectrum, at longer wavelengths,
is infrared radiation. Infrared radiation is also known as radiant heat. A heat lamp produces infrared
radiation. As wavelengths continue to increase, the radiation becomes microwaves and then radio
waves. Microwaves are used for quickly heating food in a microwave oven. The microwaves heat only
the food by transferring their energy to the moisture present in the food. On the other side of the visible
spectrum, beyond the violet portion, is ultraviolet or UV radiation. This form of radiation is
responsible for sunburns and skin cancer. X-rays and gamma rays are other forms of electromagnetic
radiation with even shorter wavelengths.
Electrons in Atoms
The Four Quantum Numbers
The locations described by the Schroedinger equation are not circular orbits, as they are in the Bohr
model of the atom, but orbitals which may have a more complex geometry. Orbital geometry is taken
up in later sections; for the moment we restrict ourselves to the permitted energy levels of electrons
within atoms.
Solutions to the Schroedinger equation describe these energy levels in terms of four quantum
numbers. Each electron in an atom is described by these quantum numbers. The quantum numbers are
all integers, which is not surprising if electrons have wave properties and the levels they occupy must
correspond to an integral number of wavelengths. The four quantum numbers are called the principal
quantum number, the momentum or subshell quantum number, the magnetic quantum number, and the
spin quantum number.
The integer values which one quantum number can have depend upon the values of other quantum
numbers describing the same electron. The permitted or allowed values of the subshell quantum
number depend upon the value of the principal quantum number with which it is associated. The
allowed values of the magnetic quantum number depend upon the value of the subshell quantum
number, and thus also the principal quantum number, with which it is associated.
Principal Quantum Number
The principal quantum number, generally symbolized by n, denotes the major shell in which the
electron is located. Usually this is simply called the shell. It corresponds to the orbit in the Bohr model
of hydrogen. The values it can take are those of any integer greater than zero, which is expressed as a
series: 1, 2, 3, 4, ... or K, L, M, N, ...
Subshell Quantum Number
The subshell quantum number, generally symbolized by l, denotes the subshell in which the electron is
located. The values it can take are those of any integer from 0 up to and including l = n-1. When n = 1,
l = 0; when n = 2, l = 0 and 1; when n = 3, l = 0, 1, 2 and so forth. The values of subshell quantum
numbers may be expressed as the series 0, 1, 2, 3, 4..., or s, p, d, f, g...
Magnetic Quantum Number
The magnetic quantum number, generally symbolized by m, denotes the energy levels available within
a subshell. It has the value of any positive or negative integer whose absolute magnitude does not
exceed l. This interdependence is summarized in the Table below.
Table: Interrelationships of the Quantum Numbers
Orbital Values
Number of Values
s
l=0, m=0
(1)
p
l=1, m=-1,0,+1
d
l=2, m=-2,-1,0,+1,+2
f
l=3, m = -3,-2,-1,0,+1,+2,+3 (7)
g
l=4, m = -4,-3,...,+3,+4
(3)
(5)
(9)
The different values of the magnetic quantum number do not correspond to different values for the
energy of an electron in the absence of an external magnetic field. Experimentally, electrons described
by different principal or subshell quantum numbers are found to have different energies regardless of
the presence of an external magnetic field, while those with different magnetic quantum numbers and
the same principal and subshell quantum numbers show differences only in the presence of such a
field.
Spin Quantum Number
The spin quantum number, generally symbolized by s, denotes the direction of the electron spin. It can
take either of two values, traditionally given as +1/2 or -1/2, sometimes as upspin and downspin. Any
other two-valued assignment, such as male-female, would have done as well.
Many atoms have an odd number of electrons or an arrangement of electrons in which the number of
positive and negative spins are not the same. These atoms or electrons are said to have unpaired spins.
The presence of unpaired spins in the electronic structure of atoms can be detected in various ways,
and chemists can make use of it to study the structures of molecules which contain these atoms.
3-4 Counting Atoms
A. Atomic Mass
The mass of one atom is about 10-23 grams.
atomic mass
the weighted average of the masses of the existing isotopes of an element
weight listed on the Periodic Table for each kind of atom
unit is the amu
amu
atomic mass unit
special unit used to describe the mass of one atom
1 amu is equal to 1/12 the mass of carbon-12
mass of helium =
mass of bromine =
mass of iron =
B. What is a Mole?
Molecular and formula masses are in atomic mass units. (a.m.u.). An atomic mass unit is only 1.66 x
10-24 grams. The mass of a single molecule is so small that it is impossible to measure in the
laboratory. So for everyday use in chemistry, a larger unit, such as a gram is needed.
We would like to choose a number of atoms that would have a mass in grams equivalent to the mass of
one atom in atomic mass units (a.m.u.). This same number would fit ALL elements, because equal
number of different atoms ALWAYS have the same mass ratio. Chemists have found that 6.02 x 1023
atoms of an element have a mass in grams equivalent to the mass of one atom in a.m.u. For example,
one atom of oxygen has a mass of 16 a.m.u.; 6.02 x 1023 atoms of oxygen have a mass of 16 grams.
The MOLE is a unit used to measure the number of particles of ANY KIND. One mole is equal to
6.02 x 1023 particles. This number is called Avogadro's number (abbreviated N)(named after
Amadeo Avogadro). So one mole of particles (atoms, ions, molecules) has a mass in grams equivalent
to that of one particle in atomic mass units.
C. What is in a Mole?
It is important to note that;
one mole of atoms = 6.02 x 1023 atoms = molar mass in grams
ex) 1 mole of Cu = 6.02 x 1023 atoms of Cu = 64 g of Cu
Mole Conversions
Calculating Mass from Moles and Calculating Moles from Mass for Elements
In other words,
given the mass of a sample, calculate the number of moles
given the number of moles of a sample calculate the number of grams.
These problems can be solved in TWO ways:
Factor Label Method (conversion factors)(proportions)
We know that:
1 mole of a element = 6.022 x 1023 atoms of that element
= molar mass (atomic weight) in grams for that element
Given (units given)
1
x Conversion ( units of unknown) = desired answer
factor (units of given)
Using Formulas
Molar Mass = mass ÷ moles
Mass (grams of the substance) = (moles) (Molar Mass)
Moles = grams of the substance ÷ Molar Mass
Number of Particles = (Moles) (Avogadro’s #)
Ex) What is the mass in grams of 0.586 mol Zn atoms?
0.586 mol Zn X 65.38 g Zn
1
1 mol Zn
= 38.31 g Zn
OR
Mass = (moles)(molar mass)
= (0.586 mol Zn)(65.38 g Zn)
= 38.31 g Zn
Sample Problems
Ex) How many atoms are in 2.83 mol Cr?
Ex) How many moles are in 6.195 g of K?
Ex) How many grams are in 8.76 x 1024 atoms of Ba?
Ex) see page 104
Homework
How many moles are in 600. g of calcium?
How many grams are in 14 moles of chromium?
How many atoms are in 7.5 moles of nitrogen?
How many moles are in 1.806 x 1023 atoms of nickel?
How many atoms are in 8.0 grams of magnesium?
How many grams are 3.01 x 1023 atoms of silver?
For Additional Problems see page 108 42- 47