Chapter Menu The Structure of the Atom
Section 4.1 Early Ideas About Matter
Section 4.2 Defining the Atom
Section 4.3 How Atoms Differ
Section 4.4 Unstable Nuclei and
Radioactive Decay
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BELL RINGER! Describe the experiment proving the positive charge
in an atom resides in the nucleus
Rutherford’s gold foil experiment demonstrated that the positive charge in atoms resides in the nucleus. In this experiment, Rutherford set up an alpha particle emitter such that particles hit a very thin piece of gold foil. Alpha particles are the nucleus of the 4He isotope and have a charge of +2. In the experiment, most of the alpha particles passed directly through the gold foil but some were deflected by large angles
Niels Bohr
1885-1962
• Planetary Model 1913
– Nucleus surrounded by
orbiting electrons at
different energy levels
– Electrons have definite
orbits
• Utilized Planck’s
Quantum Energy theory
• Worked on the
Manhattan Project (US
atomic bomb)
Ernst Schrödinger 1887-1961
• Quantum Mechanical
Model 1926
– Electrons are in
probability zones called
“orbitals”, not orbits and
the location cannot be
pinpointed
– Electrons are particles
and waves at the same
time
– Developed quantum
numbers based on
theories of Einstein and
Planck
Werner Heisenberg 1901-1976
Quantum Mechanical Theory Electron in a Hydrogen atom
The Structure of the Atom
ATOM
nucleus
energy
levels
protons
neutrons
electrons
( p+ )
( n0 )
( e- )
CHEMICAL COMPOSITION SHORTHAND
Cl 35
17
MASS
NUMBER
ATOMIC
NUMBER
NUMBER OF
PROTONS
# OF PROTONS +
# OF NEUTRONS
D. The Size of Atoms Electron
Cloud 1. Nucleus – extremely
small volume but
massive
2. Protons & neutrons
have the same mass
& are found in the
nucleus
3. Electrons are 1836
times smaller in
mass than protons
& neutrons & are
found in the
electron cloud
(large in volume)
Counting Atoms Objectives:
• Explain what isotopes are
• Define atomic number and mass
number
• Determine the number of protons,
neutrons, and electrons of a nuclide
• Define mole in terms of Avogadro’s
number, and define molar mass
A. Atomic Number
1. Represents the # of protons in an atom
Ex: Hydrogen is atomic #1 has 1 proton
2. If the atom is electrically neutral,
then # of protons (p+) = # of electrons (e-)
A = P = E
C. Using A=P=E M-A=N
M - A = N
1. How many p+, e-, and no are in
chlorine-37?
17 p+, 17 e-, 20 no
2. How many p+, e-, and no are in
bromine-80?
35 p+, 35 e-, 45 no
CHEMICAL COMPOSITION SHORTHAND
Cl 35
17
MASS
NUMBER
ATOMIC
NUMBER
NUMBER OF
PROTONS
# OF PROTONS +
# OF NEUTRONS
• EVERY CHLORINE ATOM HAS 17 PROTONS, WITHOUT EXCEPTION, –HOWEVER, NOT EVERY CHLORINE
ATOM HAS 18 NEUTRONS.
– ATOMS WITH THE SAME NUMBER OF PROTONS BUT CONTAIN DIFFERENT NUMBERS OF NEUTRONS ARE CALLED ISOTOPES.
• BECAUSE ISOTOPES OF AN ELEMENT HAVE DIFFERENT NUMBERS OF NEUTRONS THEY HAVE DIFFERENT MASS NUMBERS.
ISOTOPES
• ISOTOPES ARE CHEMICALLY ALIKE BECAUSE THEY HAVE IDENTICAL NUMBERS OF PROTONS AND ELECTRONS
– IT’S THE ELECTRONS AND PROTONS THAT ARE RESPONSIBLE FOR CHEMICAL BEHAVIOR
proton
neutron
electron
BERYLLIUM ISOTOPES
IONS • AN ELEMENT’S ATOMS ARE NOT
ALWAYS NEUTRAL IN CHARGE. – WHEN AN ATOM LOSES OR GAINS
ONE OR MORE OF ITS ELECTRONS IT BECOMES ION.
• AN ION THAT HAS MORE ELECTRONS THAN PROTONS HAS A NEGATIVE ELECTRICAL CHARGE
• AN ION THAT HAS FEWER ELECTRONS THAN PROTONS HAS A POSITIVE ELECTRICAL CHARGE
NOTE: IT’S THE PROTONS THAT DEFINE THE TYPE OF ATOM IT IS, BUT THE ELECTRONS DEFINE THE ATOM’S CHARGE.
SOME ATOMS GAIN ELECTRONS
O
-
- -
-
-
-
-
-
O-2
-
- -
-
-
-
-
-
- -
ATOM’S IONIC CHARGE =
# PROTONS - # ELECTRONS
ATOMS, IONS, AND ISOTOPES
ATOMS
NEUTRAL AND ARE DEFINED BY THE # OF PROTONS IN THEIR
NUCLEUS
3 p+ = Li ATOM, ETC.
IONS
HAVE AN ELECTRICAL CHARGE DETERMINED BY
# PROTONS - # ELECTRONS
N-2 = 7 p+ - 9 e- ; ETC.
ISOTOPES
TWO ATOMS WITH THE SAME # OF PROTONS, BUT DIFFERENT #’S OF
NEUTRONS OR MASSES
CALCIUM-40 & CALCIUM-44
atomic number:
--
--
To find net charge on an atom, consider
____ and ____.
mass number:
# of p+ the whole number
on Periodic Table
Does not change
determines identity
of atom
(# of p+) + (# of n0)
10
Ne 20.1797
p+ e–
(It is NOT on ―the Table.‖)
protium
Isotopes: different varieties of an element’s atoms
--
-- have diff. #’s of n0; thus, diff. masses
some are radioactive; others aren’t
Isotope
H–1
Mass p+ n0 Common Name
H–2
H–3 tritium
deuterium
1
2
3
1 0
1 1
1 2
C–12 atoms C–14 atoms
All atoms of an element react the same, chemically.
6 p+ 6 n0
stable
6 p+ 8 n0
radioactive
Proton
s
Complete
Atomic
Designation
92
Neutron
s
Electro
ns
34
11
146
45
12
92
36
10
59 27
3+ Co
37 17
1– Cl
55 7+ Mn 25
238 92
U
23 11
1+ Na
79 34
2– Se
20 18 17
30 18 25
32 24 27
Two Methods to Identify Isotopes
1. Hyphen notation - mass number is written with a hyphen after the name of the element
hydrogen-3, uranium-235
2. Nuclear symbol – shows composition of nucleus
232U
Superscript = mass number
Subscript = atomic number
Nuclide – general term for any isotope
of any element
92
Atomic mass unit – (amu) 1 amu is equal to
1/12 mass of a carbon –12 atom – used to express the masses of atoms
* carbon-12 is the standard for determining atomic mass
* depends on masses and relative abundance of the isotopes
Average atomic mass – weighted average of the atomic masses of the naturally occurring isotopes of an element
* may not equal the mass of any of its isotopes
* atomic mass listed in the periodic table is the average atomic mass
• An element’s atomic mass is
usually rounded to TWO decimal
places before it is used in a
calculation.
Average Atomic Mass (Atomic Mass, AAM)
This is the weighted average mass of all atoms of
an element, measured in amu.
For an element with
isotopes A, B, etc.:
AAM = Mass A (% A) + Mass B (% B) + …
(use the decimal form of the %;
e.g., use 0.253 for 25.3%)
% abundance
Ti has five naturally-
occurring isotopes
Lithium has two isotopes.
Li-6 atoms have mass 6.015 amu;
Li-7 atoms have mass 7.016 amu.
Li-6 makes up 7.5% of all Li atoms.
Find AAM of Li.
AAM = Mass A (% A) + Mass B (% B)
AAM = 6.015 amu (0.075) + 7.016 amu
AAM = 6.94 amu
(0.925)
AAM = 0.451 amu + 6.490 amu
Li batteries
** Decimal number on Table refers to…
molar mass (in g) OR AAM (in amu).
6.02 x 1023
atoms
1 ―average‖
atom
19
Description Net
Charge
Atomic
Number
Mass
Number
Ion
Symbol
15 p+
16 n0
18 e– 38 p+
50 n0
36 e–
18 e– 1+ 39
Te2– 128
K1+
Sr2+
P3–
88
31 15 3–
38
52 2–
2+
76 n0 52 p+
54 e–
20 n0 19 p+
a- or b-particles, g rays
Radioactive Isotopes:
Nucleus attempts to attain a lower
energy state by releasing extra
energy as __________.
e.g.,
half-life: the time needed for
½ of a radioactive
sample to decay
into stable matter
have too many or too few n0
radiation
e.g., C–14: -- half-life is 5,730 years
-- decays into stable N–14
Years
from now
0
g of N–14
present
5,730
11,460
120
60
30
15
7.5
0
Say that a 120 g
sample of C-14 is
found today.
g of C–14
present
17,190
22,920
60
90
105
112.5
= C–14
= N–14
Section 4.4 Unstable Nuclei and
Radioactive Decay
• Explain the relationship between unstable nuclei and
radioactive decay.
element: a pure substance that cannot be broken
down into simpler substances by physical or chemical
means
• Characterize alpha, beta, and gamma radiation in
terms of mass and charge.
Section 4.4 Unstable Nuclei and
Radioactive Decay (cont.)
radioactivity
radiation
nuclear reaction
radioactive decay
alpha radiation
Unstable atoms emit radiation to gain stability.
alpha particle
nuclear equation
beta radiation
beta particle
gamma rays
Radioactivity
• Nuclear reactions can change one element
into another element.
• In the late 1890s, scientists noticed some
substances spontaneously emitted radiation,
a process they called radioactivity.
• The rays and particles emitted are called
radiation.
• A reaction that involves a change in an atom's
nucleus is called a nuclear reaction.
Radioactive Decay
• Unstable nuclei lose energy by emitting
radiation in a spontaneous process called
radioactive decay.
• Unstable radioactive elements undergo
radioactive decay thus forming stable
nonradioactive elements.
Radioactive Decay (cont.)
• Alpha radiation is made up of positively
charged particles called alpha particles.
• Each alpha particle contains two protons and
two neutrons and has a 2+ charge.
Radioactive Decay (cont.)
• The figure shown below is a nuclear
equation showing the radioactive decay of
radium-226 to radon-222.
• The mass is conserved in nuclear equations.
Radioactive Decay (cont.)
• Beta radiation is radiation that has a
negative charge and emits beta particles.
• Each beta particle is an electron with a 1–
charge.
Radioactive Decay (cont.)
• Gamma rays are high-energy radiation
with no mass and are neutral.
• Gamma rays account for most of the energy
lost during radioactive decay.
Radioactive Decay (cont.)
• Atoms that contain too many or two few
neutrons are unstable and lose energy
through radioactive decay to form a stable
nucleus.
• Few exist in nature—most have already
decayed to stable forms.
Section 4.4 Assessment
A reaction that changes one element into
another is called what?
A. chemical reaction
B. beta radiation
C. nuclear reaction
D. physical reaction
Section 4.4 Assessment
Why are radioactive elements rare in
nature?
A. They do no occur on Earth.
B. Most have already decayed to a
stable form.
C. They take a long time to form.
D. They are too hard to detect.
• Mole (mol) – amount of a
substance that contains as many
particles as there are atoms in
exactly 12 g of carbon
* counting unit
• Avogadro’s number – 6.022 X 1023 number of particles in 1 mole of a
pure substance
• Molar mass – mass of 1 mole of a
pure substance
* numerically equal to the
average atomic mass of an
element
* must know the average atomic
mass to determine the molar mass