Download Atomic theory & structure

Introduction to the Atom
and Atomic Models
DEMOCRITUS
400
BC
Democritus believed all things
consisted of tiny indivisible units.
He called these tiny units he
called atomos. The Greek word
for “can not be cut” or
“indivisible”
Ancient philosopher: Father of the Atom
John Dalton
(1799)
 Developed what is considered to
be the 1st Atomic Theory
 Was born into a modest Quaker
family in England
 Began lecturing in public at the age
of 12
Dalton’s Model (1799)
• Dalton's model was that the atoms were tiny,
indivisible, indestructible particles and that each one
had a certain mass, size, and chemical behavior that
was determined by what kind of element they were.
Dalton’s Model
Dalton’s model of the
atom was similar to a
tiny billiard ball.
Dalton’s model of the
atom was solid and had
no internal structure.
Dalton’s Atomic Theory
• elements consisted of tiny particles called
atoms.
– all atoms of an element are identical
– atoms of each element are different from one
another; they have different masses.
– compounds consisted of atoms of different
elements combined together.
– chemical reactions involved the rearrangement of
combinations of those atoms.
Flaws in Dalton’s Model
• Dalton’s falsely believed that the atom
was the most fundamental particle.
– We now know the atom is made up of even
smaller particles we call the proton, neutron
and electron.
• Dalton’s theory could also not account for
the formation of ions (charged particles)
Daltons Atomic Model Summary
• Called: Billiard Ball Model
• Could account for
–Atoms of different atomic masses
–Elements were tiny particles
• Could NOT account for
–Though atom was smallest particle
–Did not have an internal structure
–The formation of charged particles
John J. Thomson (1897)
• Discovered the electron using
the Cathod Ray Tube (CRT)
• Thomson found that the beam
of charge in the CRT was
attracted to the positive end of
a magnet and repelled by the
negative end.
Thomson’s Hypothesis
• Concluded that the cathode beam was a stream of
negative particles (electrons).
• He tested several cathode materials and found that all of
them produced the same result.
• He also found that the charge to mass ratio was the same
for all electrons regardless of the material used in the
cathode or the gas in the tube.
• Thomson concluded that electrons must be part of all
atoms.
Thomson’s atomic model
• Called the “plum-pudding” it was the most
popular and most wildly accepted model of
the time.
Thompsons atomic model could account
for…..
• the atom having an
internal structure
• Light given off by atoms
• Atom with different
atomic masses
Thompsons atomic model could NOT
account for…..
• Empty space (had atom
filled with positive pudding)
• Formation of ions
Gold Foil Experiment
• Conducted by students of Rutherford.
• Proved that all atoms had a tiny, positively
charged center.
• Confirmed that atom’s were mostly empty
space.
Rutherford ~ early 1900s
• α-particle interaction with matter
studied in gold foil experiment
Rutherford's Nuclear Model
• 1. The atom contains a tiny dense center
– the volume is about 1/10 trillionth the volume
of the atom
• 2. The nucleus is essentially the entire mass
of the atom
• 3. The nucleus is positively charged
– the amount of positive charge of the nucleus
balances the negative charge of the
electrons
• 4. The electrons move around in the empty
space of the atom surrounding the nucleus
Rutherford’s atomic model (1911)
• Could account for:
– Empty space
– Ions
– Internal structure
– Light given off when heated
to high temperature.
• Could not account for:
– Stability
Philipp Lenard (1903)
• Aluminum foil experiment
• Lenard found that a beam of electrons was
able to pass through a sheet of Al foil with
almost no deflection.
• Lenard correctly concluded that the majority
of an atom’s volume is empty space.
Lenard’s atomic model
• Lenard’s model was composed of
dynamids.
• Lenard calculated the size of a
dynamid based on his
experimental results and found it
to be 1 billionth (1/1,000,000,000)
the size of the atom.
Lenard’s model could account for:
• Different atomic masses (based on the number
of “dynamids”).
• The internal structure of an atom
• The fact that most of the atom was empty space.
Flaws in Lenard’s model
• Formation of Ions (gaining or losing charge)
• The light given off by materials when heated to
a high temperature.
Hantaro Nagaoka
• First to present an atomic model close
to the presently accepted model.
• He came up with his model in 1903.
Nagoka’s Model
• Nagoka’s model of the atom was unstable.
– According to the laws of planetary motion, the
atom would collapse over time.
Planetary model
• Planetary model used to explain
electrons moving around the tiny, but
dense nucleus
• Nucleus contains
– Protons- existence proposed in 1900s
– Neutrons- existence proposed in 1930s
Successes of Nagoka’s model
• Atoms were able to give
off electrons to form ions.
• Accounted for the
experimental fact that
atom’s were mostly
empty space
• Explained different atomic
weights.
• Explained the light given
off when heated to high
temperatures.
Bohr
• Questioned ‘planetary model’ of atom
– Electrons located in specific levels from
nucleus (discontinuous model)
• Proposed electron cloud model based on
evidence collected with H emission
spectra
Bohr’s Atomic Model (1913)
• Bohr was a student of Rutherford.
• Improved Rutherford’s model by proposing electrons are
found only in specific fixed orbits.
• These orbits have fixed levels of energy
• This explained how electrons could give off light (gain or
lose energy)
BOHR MODEL
• Electrons are placed in energy levels
surrounding the nucleus
8e8eNucleus
(p+ & n0)
2e-
Bohr’s Atomic Model
• Could account for
– Internal structure
– Atoms of different masses
– Atom being mostly empty space
– Light given off
– Formation of positive ions
• Flaws
– Only really worked for Hydrogen
Chadwick (1932)
• Discovered the neutron by bombarding Be with
beta radiation.
• Nuclear fission released a neutron.
chart
Subatomic particle summary
Particle
Discovery by
Year
Proton
Rutherford
1911 Gold Foil Experiment
Electron
Thompson
1887 The response of
cathode ray tube to a
magnetic and electric
fields
1932 Bombarded Be with
beta radiation and a
neutron was released
Neutrons Chadwich
experiment
Review
• Describe each of the 4 different atomic
models. Give the
– Scientist Name
– Name of model
– What they could account for
– What they could not account for (flaws)
Subatomic Particles
Name
Relative
Symbol Charge mass
Location
Electron
e-
-1
Proton
p+
+1
1
Nucleus
Neutron
n0
0
1
Nucleus
1/1840 Electron cloud
Subatomic Particles (cont.)
• All atoms of an element have the same # of
protons
protons identify an atom  atomic #
• Atoms are electrically neutral
#p = #e-
• Only neutrons and protons contribute to an
atoms mass
#n + #p = atomic mass
Atom Basics
• How many p+, n0, e- do the following elements
have?
e-
Mass Number
n0
Element
p+
Sodium
11 11 23-11=12
Hydrogen 1
1
1-1=0
Chlorine
17 17 35-17=18
Tin
50 50 119-50=69
Element Symbol
Atomic Number
ISOTOPES
= atoms with the same
number of protons
but DIFFERENT
numbers of neutrons
Mass Number
Atomic Number
Element Symbol
23Na
Na-23 or Sodium-23 or 23 Na or 11
C-14 or Carbon-14 or 14C or 14C
6
B-10 or Boron-10 or 10B or 10
5B
IF you are given the number of neutrons assume it
is an ISOTOPE and add up protons and neutrons to
determine atomic mass.
Ex.
Isotope Practice
Element has
a. 6 p+, 8 n and 6 eb. 6 p+, 6 n and 6 ec. 19 p+, 21 n and 19 e29
36 n and __
29 ed. __p+,
__
1
1 n and __
1 ee. __p+,
__
Symbol
a. . C-14
b. . C-12
c. . K-40
d. Cu-65
e. H-2
IF you are given the number of neutrons assume it
is an ISOTOPE and add up protons and neutrons to
determine atomic mass.
Ions
• Ion is an atom that has gained or lost one or
more electrons and now has a charge
• Metals always lose electrons to form POSITIVE
ions
• Non-metals always gain electrons to form
NEGATIVE ions
• The charge of the element is show on the right
side of the symbol as a super script
• Example: Na+1, Zn+1, S-2, N-3
Ions practice
Element
Ca+2
Br-1
Al+3
I-1
p+
20
35
13
53
n0
40-20=20
80-35=45
27-13=14
127-53=74
Number of
electrons
20-2=18
35+1=36
13-3=10
53+1=54
D. Average Atomic Mass
• weighted average of all isotopes
• on the Periodic Table
• round to 2 decimal places
Avg.
Atomic
Mass
(mass)(% )  (mass )(% )

100
Average
= (mass)(% as decimal) + (mass)(% as decimal) +
Atomic Mass
D. Average Atomic Mass
• EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O, and
0.20% 18O.
Avg.
Atomic
Mass
(16)(99.76 )  (17)(0.04)  (18)(0.20) 16.00


100
amu
OR
Avg
Atomic= 16 (0.9976)+17(0.0004)+18(0.0020) = 16.00 amu
mass
D. Average Atomic Mass
• EX: Find chlorine’s average atomic mass if
approximately 8 of every 10 atoms are chlorine-35
and 2 are chlorine-37.
Avg.
35.40 amu
(35)(8)  (37)(2)
Atomic
Mass


10
OR
Avg
Atomic= 35(0.80)+37(0.20) = 35.40 amu
mass
What can you EXPECT on Unit 3 test
• Using Periodic table or isotope information to find
– # of protons
– # of electrons
– # of neutrons (mass – atomic number)
– Molar Mass (protons + neutrons)
•
•
•
•
Average Atomic Mass
Atomic Models (who created it, what the said)
Major experiments discussed in class
Experiments/people who discovered subatomic
particles