Download File - Mrs. Henderson

Chapter 4
Atomic Structure
Alpharetta High School
Dr. Sonha Payne
Why do scientists
use models?
 Models may be used to represent things that are difficult
to visualize.
 Scaled-down models allow you to see something too large
to see all at once, (solar system) or something that hasn’t
been built yet.
 Scaled-up models are used to visualize things that are too
small to see. (atoms)
1st Model of the
Atom
Just as a brick is basic
to the structure of a wall,
an atom is basic to the
structure of matter.
Democritus (400 BC)
atomos,
meaning uncuttable,
or cannot be divided
Foundations of Atomic Theory
Law of Conservation of Mass
(Lavoisier)
Mass is neither created nor destroyed
during ordinary chemical reactions.
Mass products = mass reactants
Foundations of Atomic Theory
Law of Definite Proportions (Proust)
A chemical compound contains the same
elements in the same percent by mass
regardless of the size of the sample or
the source of the sample.
Ex: H2O will always have the same percent
by mass, 11.2% H and 88.8% O.
Foundations of Atomic Theory
Law of Multiple Proportions
(Dalton)
When elements combine, they do so in
the ratio of small whole numbers.
CO2 and CO
H2O and H2O2
Principles of Dalton’s Atomic Theory of
Matter (1808)
1. All matter is composed of extremely small, indivisible
particles called atoms.
2. All atoms of a given element have identical properties
that differ from those of other elements.
3. Atoms cannot be created, destroyed, or transformed into
atoms of another element.
4. Compounds are formed when atoms of different
elements combine with one another in small wholenumber ratios.
5. The relative numbers and kinds of elements are constant
in a given compound.
Dalton’s Billiard Ball Model of the Atom
Billiard Ball Model
because the atom
is likened to a
billiard ball - it is a
single, complete
unit of matter.
A solid, indivisible,
indestructible sphere
No Charge Flows When the Glass Tube is Empty
(Vacuum, No Gas)
Voltage source
 Cathode Ray Tube: 2 metal plates sealed inside a
glass tube connected to a source of electricity
 Since glass is an insulator, no charge was
observed to flow when the tube was empty.
 Some means of conduction was needed if charge
was to flow.
Charge Flows (a Ray is observed)
in the Presence of ANY GAS
Voltage source
-
+
When a small amount of any gas was placed
in the glass tube and the power source was
turned on, a ray was observed striking the
phosphor-coated end of the tube and
emitting a flash of light.
For current to flow, mobile,
charged particles are required.
-
Voltage source
+
The current carriers in the gas are “invisible”
and are visualized by the phosphor-coated
screen which fluoresces as the current
carriers strike it.
 Observation: A fluorescent screen glows (emits a flash of
light) when struck by the “ray”.
 Conclusion: Invisible particles traveling from the cathode to
the anode are carrying electric current and are visualized
by the fluorescent screen.
 Observation: When a tiny object is placed in the middle of the tube,
a shadow is cast on the screen at the anode
 Conclusion: the particles travel in a straight line.
Cathode Rays Originate at the Cathode
(Negative Electrode)
 The rays were called cathode rays because
they originated at the negative electrode
(aka the cathode) and moved to the
positive electrode (aka the anode).
 The entire apparatus is now known as a
cathode ray tube (CRT).
 Observation: In the presence of a magnetic field, the
cathode rays bend.
 Conclusion: The cathode ray consists of charged particles.
 Observation: In the presence of an external electric field,
the cathode rays bent towards the positive plate.
 Conclusion: The particles are negatively charged.
ALL gases used in the tube were
found to produce identical rays, so
the negatively charged particles
of the cathode rays were
determined to be part of all
matter.
 Observation: A paddle wheel in the
middle of the cathode ray tube turns.
 Conclusion: The particles have mass.
JJ Thomson (1897)
 Credited with discovery of the electron
 Found the ratio of the mass of the particles to the charge
of the particles in cathode rays
m/e = -5.6856 x 10-9 g/C
 Conclusion: all cathode rays are composed of identical,
negatively-charged particles.
 Conclusion: these negative particles are fundamental
particles of matter. (electrons)
In 1897, JJ Thomson discovered that negatively
charged electrons were part of all matter.
Dalton Postulate 1: All matter is composed of extremely small,
indivisible particles called atoms.
This is no longer valid.
Atoms are NOT indivisible particles, but CONTAIN ELECTRONS.
Millikan Determined the Charge on an Electron
by Examining the Motion of Tiny Oil Drops (1909)
 Small drops of oil with a negative charge are examined.
 The diameter of the3 drop is measured. From this, the volume is
determined (4/3pr ).
 The mass is calculated using the known density of the oil.
 By adjusting the electric field to balance the force of gravity, the
charge on the drop is calculated.
Charge in Coulombs
13.458 x 10-19
= 7 x (1.92 x 10-19)
15.373 x 10-19
= 8 x (1.92 x 10-19)
17.303 x 10-19
= 9 x (1.92 x 10-19)
15.378 x 10-19
= 8 x (1.92 x 10-19)
17.308 x 10-19
= 9 x (1.92 x 10-19)
28.844 x 10-19
= 15 x (1.92 x 10-19)
11.545 x 10-19
= 6 x (1.92 x 10-19)
19.214 x 10-19
= 10 x (1.92 x 10-19)
Observation
The charge of each drop is a whole number multiple of some
common number.
Conclusion
Each number is divisible by 1.92 x 10-19. This is the charge of
an electron. (modern value = 1.602 x 10-19 C)
Millikan Determined the Mass of an Electron
Knowing the charge, Millikan was able to use Thomson’s chargeto-mass ratio to determine the mass of an electron.
charge
1.6022  1019 C
28
mass of an electron =
=
=
9.10

10
g
8
charge / mass
1.76  10 C/g
Inferences from the
Properties of Electrons
 Atoms are neutral, so there must be positively
charged particles to balance the negatively charged
electrons.
 Electrons have a tiny mass, so atoms must contain
other particles that account for most of their mass.
 Even the lightest atom, H, has a mass of 1.7 x 10-24 g,
compared to a mass of only 9.1 x 10-28 g for an electron.
 That is, the lightest atom is almost 10,000 x heavier than an
electron.
 Thus, most of the mass of an atom had to come from
somewhere else.
Thomson’s Plum Pudding Model of the Atom
•Since electrons are negatively
charged, and atoms are neutral,
atoms must also contain a positively
charged substance.
•Thomson’s model of the atom: the
positively charged substance fills the
atom and the electrons are
embedded throughout the substance
like “raisons in plum pudding”.
Rutherford's gold foil experiment (1911)
Discovery of the Nucleus
 (helium nuclei) a-particles were known to be heavy positive
particles.
 When the particles hit the screen, brief flashes of light
were seen.
Rutherford’s Expected and
Actual Results
If plum-pudding
model was correct.
Actual results.
Rutherford’s Expected and Actual Results
• Observation: Most particles went straight through the foil
undeflected.
• Conclusion: The atom is mostly empty space.
• Observation: Some particles bounced back!
• Conclusion
• To deflect the energetic a-particles, most of the mass and
all of the positive charge must be in a dense central region
which he called the nucleus.
• For the fraction deflected to be small, the nucleus must be
small, relative to the overall size of the atom.
• Since atoms are neutral particles, the charge of the nucleus
must be equal to the sum of the negative charges of the
electrons.
1911
Rutherford’s
Gold Foil
Experiment
 Conclusion: An atom is mostly empty space occupied by
electrons, and centrally located within that space lies a tiny
region, which he called the nucleus, that contains all the
positive charge and essentially all the mass of the atom.
 Rutherford proposed that positive particles lay within the
nucleus and called them protons.
Summary: The Rutherford Experiment (1911)

Alpha particles (helium nuclei) fired at a thin sheet of
gold

Assumed that the positively charged alpha-particles were bounced
back if they approached a positively charged atomic nucleus head
on
(Like charges repel one another)

Very few particles were greatly deflected back from the
gold sheet

Atoms contain very small, very dense, positively-charged nuclei.

The electrons are in clouds surrounding the nucleus at relatively
large distances.

Most of the atom is empty space, and if the nucleus were the size
of a ladybug, it would be in an atom the size of the Georgia Dome.
Moseley,1913
X-rays, Atomic Number and the Proton
 A beam of electrons shot at a sample of an
element gives off x-rays.
 Observation
 The frequency of x-rays given off is unique to that
element.
 Higher energy rays are given off when the nuclear
charge is higher.
 Conclusion
 The atomic number can be determined from the x-ray
data
 Each element on the periodic chart differs from the
next by having one more positive charge in the nucleus
 These fundamental positive charge units are the
protons
Bohr’s Theory of the Hydrogen Atom
Visible Light
Spectrum
When white light is passed through a prism, a
continuous spectrum of colors results which
contains all of the wavelengths of visible light
The
Hydrogen
Line
Emission
Spectrum
When a sample of hydrogen gas is excited by
electricity and passed through a prism, only a few
lines are seen, each of which corresponds to a discrete
wavelength.
Every Element Has Its Own
Unique Emission Line Spectrum
Bohr Attributed the Emission of Radiation to the
Electron Dropping From a Higher Energy Orbit
to a Lower One
1. Electrons orbit the
nucleus in circular orbits.
2. Only orbits of certain radii
are permitted.
3. An electron in a permitted
orbit has a specific
energy.
4. Energy is emitted or
absorbed as a photon
when the electron
changes its orbit.
What color of light is emitted when an excited electron in the
hydrogen atom falls from:
a) n =5 to n = 2
b) n = 4 to n = 2
c)n = 3 to n = 2
The Hydrogen e- Visualized as a
Standing Wave Around the Nucleus
Only certain circular orbits have a
circumference into which a whole number
of wavelengths of the standing electron
wave will “fit”.
Circular orbits with any other
circumference produce destructive
interference of the standing electron wave
and are not allowed.
Electron Diffraction Experiments Demonstrated that
Electrons Exhibit the Wave Property of Interference
Experiments have shown that electrons do indeed possess wavelike
properties:
X-ray diffraction pattern of
aluminum foil
Electron diffraction pattern of
aluminum foil.
Quantum Mechanics
The Heisenberg uncertainty principle states that it is impossible to
know simultaneously both the momentum p and the position x of a
particle with certainty.
Δx is the uncertainty in position in meters
Δp is the uncertainty in momentum
Δu is the uncertainty in velocity in m/s
m is the mass in kg
If the position of a particle is known more
precisely, then it’s velocity measurement
must become less precise
In the picture, we
know the exact
location of the
cars, but we have
no idea how fast
they are moving.
If the velocity of a particle is measured more
precisely then the position must become
correspondingly less precise
In the
picture, we
know the
speed of the
cars, but we
have no idea
exactly
where they
are.
Which of the following statements regarding
Dalton’s atomic theory are still believed to
be true?
I. Elements are made of tiny particles called atoms.
II. All atoms of a given element are identical.
III. A given compound always has the same relative
numbers and types of atoms.
IV. Atoms are indestructible.
43
Modern Atomic Theory
1. All matter is made up of very tiny particles called
atoms.
2. Atoms of the same element are chemically alike.
3. Individual atoms of an element may not all have the
same mass. However, the atoms of an element have
a definite average mass that is characteristic of the
element.
4. Atoms of different elements have different average
masses.
5. Atoms are not subdivided, created, or destroyed in
chemical reactions.
Modern
Atomic
Theory
 An atom is an electrically neutral, spherical entity composed of a
positively charged central nucleus surrounded by one or more
negatively charged electrons. The “cloud” of rapidly moving, negatively
charged electron occupies virtually all of the atom’s volume and
surrounds the tiny nucleus.
 The nucleus is very dense as it contributes 99.97% of the atom’s mass
but occupies only about one ten-trillionth of its volume
-10 m) is about 10,000 times the diameter of
 An atom’s diameter
(~10
-14
it’s nucleus (~10
m).
The Modern Model of the Atom
The precise paths of electrons cannot be
determined accurately. Instead, the PROBABILITY
of finding electrons in a specific location can be
determined.
The location and energy of electrons can be
specified using three terms:
Shell (aka level)
Subshell (aka sublevel)
Orbital
An additional fourth term, spin number, indicates
whether the electron is spinning clockwise or
What are the particles that make up an
atom and where are they located?
Particle
Symbol
Charge
Relative Mass (amu)
Actual Mass (g)
Proton
p+
+1
1
1.7 x 10-24
Neutron
n
0
1
1.7 x 10-24
Electron
e-
-1
0
9.1 x 10-28
Atomic Mass Unit (amu): the Unit Used For
Masses of Atoms and Subatomic Particles
Atoms have an very small mass
(e.g. hydrogen 1.67 x 10–24 g)
It is hard to work with these small numbers
so a relative mass scale was introduced:
atomic mass unit (amu or u)
1 u = 1/12 the mass of a carbon-12 atom
Particle
Symbol
Charge
Relative Mass (amu)
Actual Mass (g)
Proton
p+
+1
1
1.7 x 10-24
Neutron
n
0
1
1.7 x 10-24
Electron
e-
-1
0
9.1 x 10-28
# Protons
Defines an
Atom
The number of protons distinguishes
atoms of one element from
atoms of all of the other elements.
Atomic number = # protons in the nucleus
What Distinguishes One Element from
Another Element? The # of Protons!
For a Neutral Atom,
# protons = # electrons
For a neutral atom,
the positive charges (+1 per proton)
and negative charges (-1 per electron)
must add up to zero.
Complete the table.
Element
Pb
Atomic
number
82
Protons Electrons
8
30
Mass Number = protons + neutrons
Always a whole number
The mass number is NOT on the periodic table!
Carbon-12
Nuclear Symbol and Hyphen Notation
are Interchangeable
Hydrogen-2
Atoms of the Same Element That Have
Different # of Neutrons are ISOTOPES
Isotopes of
elements have the
same # of protons
but different # of
neutrons.
Magnesium-24
Magnesium-25
Magnesium-26
An isotope of an element is
identified by the mass number
How many protons, electrons, and neutrons do
each of these isotopes have?
Determine # p+, # e-, # n0 for the following.
Name each isotope and write its symbol.
Element
Neon
Calcium
Oxygen
Iron
Zinc
Mercury
Atomic
number
10
20
8
26
30
80
Mass number
22
46
17
57
64
204
Complete each of the following
isotope symbols:
206
?
a.
84
197
Au
c.
?
224
Ra
b.
?
d.
84 ?
36
Average Atomic Mass
Atomic mass is the mass of an atom in atomic mass units (amu).
1 amu = 1/12 the mass of a carbon-12 atom
The average atomic mass on the periodic table represents the average
mass of the naturally occurring mixture of isotopes.
Isotope
Isotopic mass (amu)
Natural
abundance (%)
12C
12.00000
98.93
13C
13.003355
1.07
= 12.01 amu
Average mass (C) = (0.9893)(12.00000 amu) + (0.0107)(13.003355 amu)
Isotopes Used in Medicine
Some radioactive isotopes are useful in the medical field.
131
53
I
Used to
detect
thyroid
problems
60
27
Co
Used in
cancer
therapy to
kill
cancerous
tissue
32
15
P
Used in
leukemia
therapy
137
55
Cs
Used to
irradiate food
to kill bacteria
and other
organisms
Isotopes
Is it possible to have a hydrogen
isotope with 3 neutrons,
hydrogen-4?
No, only certain combinations of
protons and neutrons are
possible; the others are
unstable.
Not all elements have the same
number of isotopes.
Every element up to lead (Z = 82)
has at least 1 stable isotope.
Isotope Abundance
Some elements have only 1 isotope,
while others have many more.
Not all isotopes are present in equal amounts.
Oxygen-16 99.762%
Oxygen-17 0.038%
Oxygen-18 0.200%
Total
100.00%
Chlorine-35 75.77%
Chlorine-37 24.23%
Total
100.00%
Note: the sum of the percentages is 100%
The atomic mass of an element
(the number on the periodic table)
is the weighted average mass of all its isotopes
expressed in atomic mass units.
Chlorine (Cl)
How to determine
the atomic mass of an element
What is the atomic weight of chlorine?
List each isotope, it’s mass in atomic
mass units, and it’s abundance in nature.
Isotope
Mass (amu)
Isotopic Abundance
Cl-35
34.97
75.78% = 0.7578
Cl-37
36.97
24.22% = 0.2422
How to determine
the atomic mass of an element
Multiply the isotopic abundance by the mass
of each isotope, and add up the products.
The sum is the atomic weight of the element.
34.97 x 0.7578 =
26.5003 amu
36.97 x 0.2422 =
8.9541 amu
35.4544 amu =
4 sig. figs.
35.45 amu
Answer
4 sig. figs.
Calculating Atomic Mass Using Isotopes
Average atomic mass =
å(isotope %)(isotope mass)
Example
average atomic mass of Cl:
Chlorine-35 75.77%
Chlorine-37 24.23%
Total
100.00%
Most elements have two or more
isotopes that contribute to the
atomic mass of that element
Average
Atomic
Mass of
Mg
Calculating Atomic Mass Using Isotopes
Average atomic mass =
Example
average atomic mass of O
å(isotope %)(isotope mass)
Oxygen-16 99.762%
Oxygen-17
0.038%
Oxygen-18
0.200%
Total
100.00%
= (99.762%)(15.99491464) + (0.038%)(16.9991306) + (0.200%)(17.99915939) = 16.00 u
Calculating the Average Mass of an Element
Oxygen is the most abundant element in both Earth’s crust and the human
17
16
body. The atomic masses of its three stable isotopes,
O
(99.757
percent),
O
8
8
18
(0.038 percent),
O (0.205 percent), are 15.9949, 16.9991, and 17.9992 amu,
8
respectively. Calculate the average atomic mass of oxygen using the relative
abundances given in parentheses.
Strategy Each isotope contributes to the average atomic mass based on its
relative abundance. Multiplying the mass of each isotope by its fractional
abundance (percent value divided by 100) will give its contribution to the
average atomic mass.
Solution
(0.99757)(15.9949 amu) + (0.00038)(16.9991 amu) + (0.00205)(17.992 amu)
= 15.9994 amu
Think About It The average atomic mass should be closest to the atomic
mass of the most abundant isotope (in this case, oxygen-16) and, to four
significant figures, should be the same number that appears in the periodic
table on the inside front cover of your textbook (in this case, 16.00 amu).
The Modern Periodic Table
Groups contain elements with similar properties
and are arranged in vertical columns.
Periods are the horizontal rows of elements.
Two Numbering Systems for Groups
 The old method uses the letter A for the representative
elements (1A to 8A) and the letter B for the transition
elements.
 The new method numbers groups 1–18 from left to right.
Alkali Metals
Special Names of
Groups
Halogens
Shape of the Periodic Table with
Lanthanides and Actinides
Inserted Into the Main Section
Metals, Nonmetals, and Metalloids
The heavy zigzag line
separates
metals and nonmetals.
 Metals are located to the left.
 Nonmetals are located to the
right.
 Metalloids are located along
the heavy zigzag line between
the metals and nonmetals
(except for Al)
Metals
 solids at room temperature, except Hg
 reflective surface (shiny)
 conduct heat and electricity
 Malleable (can be shaped)
 Ductile (can be pulled into wires)
 lose electrons and form cations in
reactions
 About 75% of the elements are metals.
79
Nonmetals
 Exist as s, l, g
 poor conductors of heat and electricity
 Solids are brittle.
 gain electrons in reactions to become
anions
 upper right on the table
 except H
Sulfur, S(s)
Chlorine, Cl2(g
Bromine, Br2(l)
80
Metalloids
 show some properties of
metals and some of
nonmetals
 also known as
semiconductors
Properties of Silicon
shiny
conducts electricity
does not conduct heat
Properties of Metals, Nonmetals,
and Metalloids
Metals
Shiny (s)
Ductile,
malleable
Good
conductors
Metalloids
Nonmetals
Dull (s, l, g)
Brittle
Better conductors than Good insulators
nonmetals, but not as
good as metals