Chapter 4: Atomic Structure
Chapter 4: Atomic Structure
Early models of the atom:
Have you ever been asked to believe in
something you couldn’t see?
Love, faith, truth… Santa Claus?
Using your unaided eye, you cannot see
the tiny fundamental particles that make
up matter…
atoms.
Add to notes next to defining the atom:
Democritus’s Atomic Philosophy
400 BC
•Atoms are indivisible and
indestructible
•Did not explain chemical behavior
•Lacked experimental support,
because he didn’t use any type of
scientific method
Early philosophers and scientists could not observe
individual atoms, but they were still able to propose
ideas about the structure of atoms.
Dalton’s Atomic Theory (1766-1844) English Chemist and
Schoolteacher
1. All elements are composed
of tiny indivisible particles
called atoms.
2. Atoms of the same element
are identical. Atoms of any
element are different from
atoms of other elements.
3. Atoms of different
elements can chemically
combine in simple whole
number ratios to form
compounds.
4. Chemical reactions
occur when atoms are
separated, joined or
rearranged.
BUT, 1 atom can NOT
change into another.
The smallest particle of an element that
has all the properties of that element
Electron - Negative charge
Proton - Positive charge
Neutron - No charge (neutral)
Nucleus: protons & neutrons
Subatomic Particles: What makes up the atom
Rutherford’s Gold Foil Experiment DO NOT WRITE THIS IN NOTES:
Rutherford aimed his beam of alpha
particles at a sheet of gold foil surrounded
by fluorescent screen
73 through foil
12 deflected
9 bounced back
“ah hah”
LOTS of empty space In an atom!!!
Rutherford’s Gold Foil Experiment
Summarize on your note page.
Rutherford’s Gold Foil Experiment
Current Atomic theory:
•Atoms can be seen with high powered microscope
•Scanning, tunneling microscope
•Generate images of individual atoms
•“Nanoscale” technology = atomic scale
NanoTechnology:
Top: Manganese atom
Bottom:
aligning theoretical
Configuration with
the image
National Institute of Standards and Technology (NIST)
and the Naval Research Laboratory
http://www.sciencecodex.com/nist_researchers_put_a_new_spin_on_atomic_musical_chairs
Summary Time
Up and At ‘em... Atom Ant!
Mystery Box
Activity
Black Box Activity
Obscertainers
Ch 4.3
Distinguishing Among
Atoms
6 protons
=
atomic number of 6
Atomic Number: the number of protons
in the nucleus
Mass Number: The number of protons
and neutrons in the nucleus
6 protons and 6 neutrons
=
mass number of 12
Neutrons: The number of neutrons is the
difference between the mass number and
atomic number
12 items in nucleus
(mass number)
6 Protons (atomic
number)
# Neutrons = 12 – 6 = 6
Electrons: Equal to
the number of
protons
WHY?
ADD THIS ONE TO
YOUR NOTES…
Atoms are always
neutral, so + charges
have to equal - charges
# Protons = 6
SO
# Electrons = 6
Atomic number =
Mass # =
Protons =
Electrons =
Neutrons =
Practice!
Atomic number =
Mass # =
Protons =
Electrons =
Neutrons =
79
Au 197
2
He 4
197-79 = 118
2 4-2 =
79 79
79 197
2
2 4
2
Carbon – 13
Nitrogen – 15
Radium – 226
Practice! • How many neutrons?
16 O 8
32 S 16
108 Ag
47
At. Mass – At # =
16 – 8 = 8 At. Mass – At # =
At. Mass – At # =
At. Mass – At # =
At. Mass – At # =
At. Mass – At # =
226 – 88 = 138
108 – 47 = 61
15 – 7 = 8
32 – 16 = 16
13 – 6 = 7
Atom Electron
Proton Nucleus
# of electrons
= # of protons in neutral atom
Negative
charge Mass:
1200 e- to
equal 1
proton
Neutron
# Neutrons =
Mass Number –
Atomic number
No
Charge
# of Protons
= Atomic
Number
Positive
charge
Summary Time!
Isotopes
• Atoms that have the same
number of protons but
different number of
neutrons
• Because isotopes of an
element have different
number of neutrons
they have different
mass numbers
Isotopes of Carbon
Isotopes of Sodium:
23 Na 11
24 Na 11
Atomic Mass
AMU • Atomic Mass Unit
• 1 amu is defined as
1/12 of the mass of a
carbon-12 atom
Atomic Mass
• The weighted average of all possible masses of 1 kind of atom (isotopes)
• These must occur naturally
• Depends on how many isotopes and abundance of each type
• Chlorine-35 (75%) • Chlorine-37 (25%)
Calculating Atomic Mass
• The number of isotopes
• The mass of each
• The natural percent
abundance of each
What you need to know
Calculating Atomic Mass
Element X has 2 isotopes Isotope 10X mass = 10.012 amu
Relative abundance is 19.91% = 0.1991
Isotope 11X mass = 11.009 amu
Relative abundance is 80.09% = 0.8009
Example
Calculate
isotope mass x abundance 10X: 10.012 amu x 0.1991 = 1.993 amu 11X: 11.009 amu x 0.8009 = 8.817 amu
+
weighted Atomic mass = 10.810 amu
Practice!!
Summary Time!