Download Atomic Theory - Honors Chemistry

Atomic Theory
Greek Philosophers: 4
Elements – earth, air,
fire and water
Democritus: world made of two things –
empty space and tiny particles called
“atomos”

atoms smallest particles and each
substance had its own type of atom
 - wood atoms, air atoms, water atoms

Dalton:
1. all matter is made of tiny particles called
atoms
2. atoms can’t be broken down further
3. atoms of different elements differ
4. atoms of the same element are identical
5. atoms combine to form compounds in
specific ratios and can be rearranged to
make new compounds
Atomic Composition:
Nucleus surrounded by an electron cloud
nucleus contains protons and neutrons
Charges:
Protons =
positive (+)
Electrons = negative (-)
Neutrons = neutral (0)
Mass:
Protons =
1.6726 × 10-27 kilograms or 1 amu
Neutrons = 1.6749 × 10-27 kilograms or 1 amu
Electrons = 9.10938188 × 10-31 kilograms or 0
amu

Atomic Research

Standard Model of Particle Physics
Breaking Down the Nucleus:
Protons and Neutrons
Protons:
- the number of protons in the atoms of an
element is always the same
- changing the protons will change the
type of element
The Number of Protons = The Atomic
Number
The Number of Protons and Neutrons = The
Atomic Mass
Element Symbols
Symbol, Atomic Mass, Atomic Number
12
C
6
At. Mass
Element Symbol
At. #
- information: element, # of protons (and
electrons by implication), # of neutrons
Differences Among Atoms of the Same
Element
IONS:
in electrically neutral atoms, # protons = #
of electrons
loss or gain of electrons  ion
Cation = (+) ion – loss of electrons
Anion = (-) ion – gain of electrons
Ex: Sodium – atomic # = 11
11 protons, 11 electrons
- lose 1 electron = 11 p+ and 10 e- = (+1)
charge
Chlorine – atomic # = 17
17 protons, 17 electrons
- gains 1 electron = 17 p+ and 18 e- = (-1)
charge
Calculating Atomic Mass and
Charge
Atomic Number = # Protons
 Atomic Mass = Protons + Neutrons
 Atomic Charge = Protons + Electrons

EX:
An atom has an atomic number of 12, 13
neutrons and 12 electrons.
Identity:
Atomic Mass:
Charge:
EX:
An atom has an atomic mass of 35, 18
neutrons and 18 electrons.
Identity:
Protons:
Charge:
EX:
An atom has 20 protons, 21 neutrons and 18
electrons.
Identity:
Atomic #:
Atomic Mass:
Charge:

Page 113, # 64
ISOTOPES:
same number of protons, different numbers
of neutrons
changes the mass of the atom
ATOMIC MASS
Atomic mass = # of protons and neutrons
each has about the same mass which is
designated as an atomic mass unit (1 amu)
Determination of amu
one element chosen as a standard – Carbon 12
6 protons, 6 neutrons = 12 amu
therefore 1 amu = 1/12th Carbon atom
HOWEVER:
Is the Atomic Mass of an Element a
whole number? NOPE
WHY? The Atomic Mass on the Periodic
Table is the AVERAGE ATOMIC MASS
OF ALL THE KNOWN ISOTOPES AND
THEIR ABUNDANCE
Average Atomic Mass:
Mass Spectrometer:
Sample of gaseous element
Charged
Propelled by electromagnetic fields toward a
photographic plate which records how much is present
Less massive fall shorter
More massive fall farther
Allows for determination of relative mass
Masses are then combined and averaged  decimal

Mass Spectrometer
Calculating Average Atomic Mass:
[(# of atoms X mass of isotope A) + (# of atoms X
mass of isotope B) + . . .] divided by (total
number of atoms of all isotopes combined)
OR
(Mass of Isotope A X Relative Abundance) +
(Mass of Isotope B X Relative Abundance) + . . .
= Average Atomic Mass
Page 104: # 15 Average Atomic Mass of Boron
= (10.013 x .198) + (11.009 x .802) = 10.8 amu

Boron has two naturally occurring
isotopes: Boron-10 (abundance = 19.8%,
mass = 10.013 amu) and Boron-11
(abundance = 80.2%, mass = 11.009 amu.
Calculate the atomic mass of boron.

Calculate the atomic mass of
magnesium. The three magnesium
isotopes have atomic masses and relative
abundances as follows:
23.985 amu (78.99%)
24.986 amu (10.00%)
25.982 amu (11.01%)
Page 113, #s 66 - 68
Radioactive Decay
Unstable Nuclei:
Radioactive Atoms - Radioisotopes
atoms are unstable because they are high
in energy
atoms give off the energy (radiation) to
become more stable
Process of losing the energy is radioactive
decay

Atoms can actually become other elements
Radioactive Decay
Alpha particles – positively charged Helium
Nucleus
Alpha particle decay: Unstable nucleus loses
an alpha particle – result, atom loses two protons
and two neutrons
mass decreases by 4 amu
atom becomes another element
238
U 
92
234
Th
90
4
+ He
2
Radioactive Decay
Beta particles – negatively charge particle (electron)
Beta particle decay: results from the break down of a
neutron into a proton and an electron
- atom becomes another element
1
1
n 
0
234
90
Th 
0
p +
e
1
-1
234
Pa
91
0
+
e
-1
Radioactive Decay
Gamma waves – high energy (no matter, no
charge)
Gamma Ray Emission: following Beta
particle decay the nucleus still has energy so
the nucleus releases it as gamma rays
- both the atomic mass and number
remain the same
Means lots of energy
230m
Th
90
230

Th
90
Gamma symbol
+
γ
Radioactive Decay
Positron Emission – release of a positive
electron (positron)
the nucleus does not lose mass but does
decrease in atomic number – the mass
of a positron is nearly zero
13
N
7
0

e
+1
13
+
C
6
Radioactive Decay
Electron Capture – rare instance where an
electron runs into the nucleus
– the electron combines with a proton to
form a neutron
– therefore the mass does not change but
the atomic number does
41
0
Ca +
20
e
-1
41

K
19

Page 814, #s 6 - 9
Calculating Nuclear Decay:
Half-life
time required for half of the sample to decay to the products
Ex: Uranium-238 takes 4.47x109 years
If you had a 10.0g sample, it would take 4.47x109 years for it to
decay to 5.0g
It would take 4.47x109 more years to degrade to 2.5 g
It would take 4.47x109 more years to degrade to 1.25 g…….
Ex: Substance X has a half life of 10 years.
Half Life
0
1
2
3
4
Time
0
10
10 more
20 total
10 more
30 total
10 more
40 total
Percent
Remaining
100%
50%
25%
12.5%
6.25%
Fraction
Remaining
1
½
¼
1/8
1/16
Equations:
Amount Remaining = (Initial amount)(1/2)n
n = number of half lives that have passed
Amount Remaining = (Initial amount)(1/2)t/T
t = elapsed time
T = duration of half-life
Ex: Radioactive iodine-131 has a half-life of
8.04 days
1) If you have 8.2 ug (micrograms) of this
isotope, what mass remains after 32.2
days?
2) How long will it take for a sample of
iodine-131 to decay to 1/8 of its activity?

Page 819, #’s 17 -19