Download The Atom Powerpoint

Bell Ringer
List 3 facts about atoms, which you’ve learned in
past courses (do not talk with neighbors… come
up with your own ideas!)
CHAPTER 3
Protons, Neutrons, Electrons
Atomic Mass, Ions and Isotopes
Atom History
Atom Models
Atom
•
Greek for “Atomos” which means “indivisible”
•
Smallest unit that can exist alone or in combination with other atoms
Structure of an Atom
1. Nucleus – contains protons (p+) and neutrons (n0)
– dense, small
2. Electron Cloud – contains electrons (e-)
– surrounds nucleus, mostly empty space
– arranged in shells or energy levels
Atom itself is neutral: Why?
Subatomic Particles
Phet Interactive
• Watch for patterns:
– What happens when a proton, electron or
neutron is added to the charge? How about the
mass?
– What element has which particles in it?
– Which atoms are stable? Which are not?
• List the parts of an atom that have mass.
• List the charges of each subatomic particle.
Subatomic Particles
Particle
Symbol
Charge
Mass
Number
Proton
p+
+1
1
Neutron
n0
0
1
Electron
e-
-1
0
Atomic Number – number of p+ in
the nucleus of an atom (always
equal to number of e-)
‘Big Blue’
Atomic mass – the average mass of
the isotopes
‘little red’
Mass Number – number of p+ and
n0 in the nucleus of an atom
The mass number is the atomic
mass rounded off to a whole
number!!
Periodic Table
How many protons, neutrons and electrons are
in aluminum?
p+ = Atomic Number (Big Blue)
n0 = Mass Number – p+
Mass Number is the Atomic Mass (little red) rounded to a whole number
Mass Number = p+ + n0
e- = p+ because atoms are electrically neutral
p+ = 13
n0 = 27-13 = 14
e- = p+ = 13
Atomic Mass
Atomic #
Mass #
# p+
# e# n0
Try Sodium (Na):
Atomic Mass =
Atomic # =
Mass # =
# p+ =
# e- =
# n0 =
=
=
=
=
=
=
(Avg. mass of isotopes)
(number of p+ or e-)
(Atomic mass rounded)
(same as Atomic #)
(same as Atomic #)
(Mass # - Atomic #)
________
________
________
________
________
________
=
=
=
=
=
=
32.065
16
32
16 p+
16 e32 – 16 = 16 n0
Find the number of protons, neutrons and electrons
for potassium.
1. 19
39.098
K
2. Atomic Number = 19
Atomic Weight = 39.098
Mass Number = 39
Proton =
Neutrons =
Electrons =
Find the number of protons, neutrons and electrons.
1. 19
39.098
K
2. 84
36
Kr
3. Uranium -235
Protons Neutrons Electrons
K 19
20
19
Kr 36
48
36
U 92
143
92
Bell Ringer
Define:
Atomic Mass (also known as atomic weight)
Mass Number
Atomic Number
More on Structure
Atom is very, very small overall
• Measure in Angstroms (1 Å = 1 x 10-10 m)
•
•
Nucleus is extremely tiny compared to the
electron cloud
How much smaller???
Nucleus is a marble in a football stadium.
Nucleus “marble”
Electron Cloud “football stadium”
Particles in the Atom
Nucleus – protons
– neutrons
Cloud – electrons
mass: 1 amu
mass: 1 amu
mass: “0” amu (5.45 x 10-10 amu)
1 amu (atomic mass unit) = 1.673 x 10-24 g
0.000000000000000000000001673 g
amu – special unit for the mass of an atom
– 1/12 mass of carbon-12 atom
Similar to:
• 2000 lbs = 1 ton
• 1 amu (atomic mass unit) = 1.673 x 10-24 g
• Ion: When the # of p+ ≠ # e– This changes the charge of the atom
• Isotope: When two atoms of one element
have different masses.
Mass Number v Atomic Weight
Atomic weight: The average atomic mass of all the
isotopes for an element
– Measured in “amu” (atomic mass units)
– AKA atomic mass
Mass Number: The mass of a single atom for one
isotope of an element
- measured in amu
Bell Ringer
• What happens when you change the protons,
neutrons, or electrons in an atom?
Example of an Average Mass
You have a bag of marbles. Half of the marbles have a mass of 5 g, 30%
of the marbles have a mass of 2 g, and 20% of the marbles have a
mass of 1 g. What is the average mass of a marble in your bag?
Mass of marble
%
5g
.5
2g
.3
1g
.2

Avg.  (5g  .5)  (2g  .3)  (1g  .2)
Avg.  2.5g  0.6g  0.2g  3.3g

Example: The 4 Isotopes of Chromium
(amu x %)
Isotope
Atomic Mass
Units
%
Cr - 50
49.9461 amu
.0435
2.17 amu
Cr - 52
51.9405 amu
.8379
43.52 amu
Cr - 53
52.9407 amu
.0950
5.03 amu
Cr - 54
53.9389 amu
.0236
1.27 amu
2.17amu  43.52amu  5.03amu 1.27amu  51.99 amu
Aristotle – 350 BC
• All matter composed of
4 elements: earth, air,
fire, water
• People believed for
over a thousand
years…into the middle
ages
Democritus – 400 BC
• Greek Philosopher
• The idea of the atom
• Not scientific, based
on philosophy not
evidence
• First to use ‘atomos’
the smallest particle
of matter that cannot
be divided
Antoine Lavoisier - 1782
• With the recent invention
of the balance - he did a lot
of quantitative
measurements
• Leading him to the Law of
Conservation of Mass –
which states mass cannot
be created nor destroyed
• Mass of reactants will equal
mass of products
Joseph Proust - 1799
• Law of constant
composition–a compound
is composed of exact
proportions of elements by
mass regardless of how the
compound was created
NaCl by mass, is always:
39.34% Na and 60.66 % Cl
John Dalton – 1803
• English Chemist/schoolteacher
• Law of multiple proportions – if
compounds are composed of the
same elements, the masses of
the elements can be expressed
as ratios of whole numbers.
Carbon monoxide CO
C = 12 g
O = 16 g
C = 12 g
O = 32 g
Ratio 1 : 1
Carbon dioxide
CO2
Ratio 1 : 2
Same elements but different ratios make different compounds
Dalton’s Atomic Theory
1.
All matter is composed of
extremely small particles.
2. Atoms of a given element are
identical in size, mass, and other
properties; atoms of different
elements differ in size, mass, and
other properties.
3. Atoms cannot be divided, created,
nor destroyed.
4. Atoms combine in whole number
ratios to form compounds.
5. Atoms are combined, separated, or
rearranged in chemical reactions.
What ideas did each of these
people have about the atom?
Democritus
Aristotle
Lavoisier
Dalton
Joseph John Thomson - 1877
• Discovered electrons had a negative charge
• Cathode Ray Tube
How did we figure out all this stuff about the atom??? Because the atom is so small !!!!!
Joseph Thompson
THE ELECTRON!
Used a cathode ray tube to prove there were
negative charged particles (now known as
electrons) in an atom. This opened the way
to the idea that an atom was not just a solid
sphere not able to be broken down
anymore.
1897
Robert Millikan - 1909
• Oil Drop Experiment
• Figured out the
charge of the
electrons
Ernest Rutherford - 1911
• Gold Foil Experiment
• Nucleus, very small
and very dense
• Positive charge part
was a proton
Rutherford’s Model
• There was a nucleus – very small, very
dense
• It is positively charged
• Atoms are mostly empty space
James Chadwick - 1932
• English physicist
• Worked with
Rutherford
• Discovered the
neutron
Niels Bohr - 1913
• Danish physicist
• Worked with
Rutherford
• Theory explained how
the electrons moved
around the nucleus
• Electrons moving around
nucleus can only follow
certain orbits
What are the four main models of
the atom?
Name each, and draw a picture to represent it
The Evolving Atomic Model
Dalton Model
Thomson Model
1803 – John Dalton
believed that an atom
was an indestructible
particle with no
internal frame.
1897 – J.J. Thomson
discovers the electron.
He believed electrons
were embedded in
positive charge sphere.
(Solid Ball Model)
(Plum pudding Model or
chocolate chip cookie
dough)
Rutherford Model
Bohr Model
1911 – Ernest
Rutherford discovers
that there is a dense,
positively charged
nucleus. Electrons
go around the
nucleus.
1913 – Niels Bohr
enhances Rutherford’s
model by having
electrons move in a
circular orbit at fixed
distances from the
nucleus.
(Nuclear Model)
Bohr Model
Nucleus
Shells
K
L
e- = 2
e- = 8
M
N
p+
n0
e- = 18
e- = 32
Fill shells inside to outside
Bohr Model of the Atom
– Similar to concert seating!
Put the number of p+ , n0 , and e - in the diagram as shown…
Lets diagram Sulfur:
p+ = Atomic # = 16
e - = Atomic # = 16
n0 = Mass # - Atomic # = 32 – 16 = 16
Maximum electrons
= 2
(seats)
p+ = 16
n0 = 16
e- = 2
8
e- = 8
18
32
e- = 6
e- =
Now try Bromine…
p+ =
n0 =
e- =
e- =
e- =
e- =
Now try Phosphorus…
p+ =
n0 =
e- =
e- =
e- =
e- =
Bell Ringer
1. What is the difference between an ion and
an isotope?
Please be checking side board for assignment
updates