Transcript Average Atomic Mass

Chapter 4: Atomic Structure
• Democritus believed that matter was made up
of particles.
• he called nature’s basic particle an “atom”.
Aristotle believed that everything was made
up of 4 substances:
Fire
Air
Water
Earth
The “People’s Choice”…… Aristotle’s idea was accepted
for nearly 2000 years! (poor Democritus )
Foundation of Atomic Theory
Basic laws of chemistry:
1. Law of conservation of mass: states that mass is
neither created nor destroyed during ordinary
chemical or physical reactions.
2. Law of definite proportions: a chemical compound
contains the same elements in exactly the same
proportions by mass regardless of the sample size.
3. Law of multiple proportions: If two or more
different compounds are composed of the same
two elements, then the ratio of the masses of the
second element combined with a certain mass of
the first element is always a ratio of small whole
numbers.
Dalton’s Atomic Theory
1. All matter is composed of extremely small
particles called atoms.
(atom: the smallest particle of an element
that retains the chemical and physical
properties of that element.)
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 subdivided, created, or
destroyed.
4. Atoms of different elements combine in
simple whole-number ratios to form
compounds.
5. In chemical reactions, atoms are
combined, separated, or rearranged.
Note: Not all aspects of Dalton’s atomic theory are correct.
•
atoms ARE divisible into even smaller particles.
•
atoms of a given element CAN have different masses.
The Structure of the Atom
•
all atoms consist of two regions:
1. Nucleus: small region located at the center of an
atom that contains protons (positive particles) and
neutrons (neutral particles).
2. Outside nucleus: very large region surrounding the
nucleus that contains electrons (negative particles).
• Protons, neutrons, and electrons are all considered to be
subatomic particles.
Discovery of the Electron
•
JJ Thomson conducted experiments using
Cathode Ray Tubes (CRT).
•
Results: All cathode rays are composed of
identical negatively charged particles.
Charge and Mass of the Electron
•
Robert Millikan experimentally determined
the mass of an electron to be 9.109 x 10-31
kg.
also confirmed the following:
JJ Thomson
Cathode
Ray
Tube 1
Cathode
Ray
Tube 2
Millikan’s
Experiment
1. Electron’s have a negative charge.
2. Atoms are electrically neutral, thus they
must contain positive particles too.
3. The other particles in an atom account for
most of the mass.
Robert Millikan
JJ’s “Plum – Pudding”
Model of the Atom
JJ Thomson
Discovery of the Atomic Nucleus
Ernest Rutherford’s Gold Foil Experiment
• bombarded a thin, gold foil with alpha particles.
• he expected to see the heavier alpha particles
pass through relatively undisturbed.
• most particles did, however a few bounced
straight back.
Result: there must be a very densely packed
bundle of matter with a positive charge  nucleus
Particle
Symbol Charge Relative Location
Mass
Electron
e-
-1
0 amu
Outside
(e- cloud)
Proton
p
Neutron
n
+1
0
1 amu
1 amu
nucleus
nucleus
If the nucleus contains positive protons, what keeps the
nucleus together?
• neutrons provide the “glue” of the nucleus and are
responsible for nuclear forces (prevent the nucleus from
breaking apart)
Isotopes
• the identity of the atom is determined by the number of
protons.
ex: 1 proton atom  Hydrogen
• however, like many other elements, hydrogen atoms can
contain different numbers of neutrons  ISOTOPES
ex: 3 isotopes of H
protium  1 proton, 1 electron
deuterium  1 proton, 1 neutron, 1 electron
tritium  1 proton, 2 neutrons, 1 electron
Atomic Number: indicates the number of protons in an
atom.
Mass Number: indicates the total number of protons and
neutrons present in an atom.
Let’s look back at the isotopes of hydrogen.
Protium (1 proton, 1 electron)
Mass Number
Atomic Number
1
H
1
Hyphen notation
Hydrogen - 1
Nuclear
Symbol
Deuterium (1 proton, 1 neutron, 1 electron)
2
H
1
Hyphen notation
Hydrogen - 2
Tritium (1 proton, 2 neutrons, 1 electron)
3
H
1
Hyphen notation
Hydrogen - 3
My definition of an isotope:
Isotopes are atoms of the same element that have
different masses.
Example: “Tin” has 10 stable isotopes……the most of any
element!
Relative Atomic Mass
• because atoms have such a very small mass, chemists use
atomic mass units (amu) to describe the mass of atoms.
1 amu = 1/12 the mass of a carbon-12 atom
Average Atomic Mass
• since most elements occur naturally as mixtures of
isotopes, scientists refer to the average atomic mass of an
element.
•
the average atomic mass of an element depends on
two factors:
1. The abundance of each of the element’s isotopes.
2. The mass of the element’s isotopes.
Average atomic mass = (% abundance #1)(amu of #1) +
(% abundance #2)(amu of #2) + …
Calculate the average atomic mass for copper:
Atomic mass
Isotope
(amu)
Percent natural
abundance
Copper – 63
62.929 599
69.17
Copper – 65
64.927 793
30.83
(.6917 x 62.929 599) + (.3083 x 64.927 793)
Average atomic mass for copper (Cu) = 63.54 amu
Do now “Isotope Problems”
 How many protons, neutrons and electrons are in an
atom of carbon – 13?
6 protons
7 neutrons
6 electrons
 Write the nuclear symbol for oxygen – 16
16
O
8
 Write the hyphen notation for the element whose
atoms contain 7 electrons and 9 neutrons.
Nitrogen - 16
Pennium Lab
Calculations
#1. Determine the number of isotopes of Pennium.
<Hint> Group your data so that there is no more
than 5 isotopes.
#2. % abundance of each isotope.
% abundance = (total mass of the isotope) x 100
mass of 20 pennies
#3. Average atomic mass of each isotope.
Isotope mass = (total mass of each isotope)
# of pennies of that
isotope
#4. Atomic mass of Pennium
Atomic mass of Pennium = (% abundance #1)(amu of
#1) + (% abundance #2) )(amu of #2) + …