Download Chapter 4 – Structure of the Atom

Chapter 4 – Structure of the Atom
Early Theories
Greek philosopher, Democritus (460370 B.C.) proposed that the world was made up of
empty space and tiny particles called “atoms”.
Aristotle, a most influential philosopher (384 322 B.C.) proposed that matter was
continuous and this theory was widely accepted until the 19th century.
In the early 1800’s, John Dalton (1766-1844) studied many chemical reactions and
proposed Dalton’s Atomic Theory.
a) All matter is composed of very small particles called ATOMS.
b) Atoms of the same elements are identical.
c) Atoms cannot be created, divided or destroyed.
d) Different atoms combine in simple whole number ratios to form compounds.
e) In a chemical reaction, atoms are separated, combined or rearranged. (explains
Law of Conservation of Mass)
Questions:1. Which statements in Dalton’s original atomic theory are now considered
to be incorrect?
2. Why was Democritus’s ideas rejected by other philosophers of his time?
Definition of an Atom
The smallest particle of an element that retains the properties of the element is called an
ATOM.
Discovering the Electron
J. J. Thomson and William Crooke experimented with cathode ray tubes. These were
glass tubes that had been evacuated (that is most of the air is removed) and electricity was
passed through this near vacuum using two electrodes, an anode and a cathode.
William Crooke was the first person to observe a CATHODE RAY, which was actually a
stream of negatively charged particles called ELECTRONS.
J. J. Thomson (1856-1940) measured the charge/ mass ratio of an electron. He concluded
that the mass of the charged particle was much less than a hydrogen atom (the lightest
element). Therefore, there were particles smaller than the atom, the ELECTRON.
J. J. Thomson proposed the plum pudding model of atomic theory. Spherical atom
composed of uniformly distributed positive charge within which reside negatively
charged electrons.
Rutherford and the Nucleus.
Discovery of Protons and Neutrons
Rutherford concluded that the nucleus contained positively charged particles called
PROTONS. A proton is a subatomic particle carrying a charge equal and opposite to the
charge on an electron.
James Chadwick (1891-1974) discovered another particle in the nucleus – a neutral
particle that he called the NEUTRON. It has a mass similar to that of a proton and has no
electrical charge.
Subatomic Particles
Particle
Symbol
ELECTRON
e-
PROTON
p+
NEUTRON
no
Location
Relative
Electrical
Charge
Relative
Mass
Actual Mass
(g)
The nucleus is extremely dense and consists of neutrons and protons. The nucleus is
always positively charged due to the presence of protons.
Since the atom is electrically neutral:# of protons (p+) = # of electrons (e-)
The nucleus contains virtually all of the mass of the atom (neutrons + protons). Most of
the atom consists of empty space around the nucleus. Fast moving electrons travel
through this empty space. They are held within the atom by the attraction of the positive
nucleus.
Atomic Number
Each element has its own atomic number.
Atomic number = number of protons.
Since the atom is electrically neutral
number of protons
= number of electrons
p+ = eMass Number and Isotopes
Atoms that have the same number of protons but different numbers of neutrons are called
Isotopes.
The Mass Number
=
Number of PROTONS
Consider Potassium, which has three isotopes
potassium –39
Protons
19
Neutrons
20
Electrons
19
+
potassium –40
19
21
19
Number of NEUTRONS
potassium –41
19
22
19
Notation
Relative
percentage
abundance
Notation
Mass number
X
Atomic number
Practice Problem 14:
Element
Isotope
b)
calcium-46
c)
oxygen-17
d)
iron -57
e)
zinc -64
f)
mercury -204
Atomic
Number
Mass
Number
p+
Neutrons
Symbol
Mass of Individual Atoms
The mass of an electron is about 1/1840th of the mass of a proton or neutron. Since the
masses of p+, no and e- are so small, chemists have developed a more convenient method
or measuring the mass of an atom relative to the mass of the carbon –12 atom which is
chosen as the standard, and is given an atomic mass of 12 a.m.u. (atomic mass unit)
Therefore one atomic mass unit (a.m.u.) is defined as equal to 1/12th the mass of a carbon
–12 atom.
The atomic mass of an element is the weighted average mass of the isotopes of that
element, e.g. chlorine exists naturally as about 23% 3717Cl and about 75% 3517Cl and so
the weighted average atomic mass number will be about 35.5 a.m.u.
Atomic mass
Percent Abundance
Mass contribution
35
17Cl
37
17Cl
34.969 a.m.u.
36.966 a.m.u.
75.770%
24.230%
Weighted average atomic mass for chlorine =
Practice Problem #15, page 104
Boron has 2 isotopes:Boron –10
Boron –11
For
abundance 19.8%, mass = 10.013 a.m.u.
abundance 80.2%, mass = 11.009 a.m.u.
10
B, mass contribution =
11
B, mass contribution =
Atomic mass of B =
Try practice problem #17.
For
24
Mg: mass contribution =
25
Mg: mass contribution =
26
Mg: mass contribution =
Atomic mass of magnesium =
Page 113, Qu. 68
NEWS RELEASE FROM VILLA WALSH CHEMISTRY DEPARTMENT
A substance thought to be a new element has been found on the salad bar in the café.
Students have named the element “Beanium”.
Students in the Chemistry class are being asked to examine the material and determine
some of its properties.
Since “Beanium” atoms are visible to the naked eye, this research will be fairly easy (no
other known element has atoms that are visible to the naked eye or even the most
powerful microscopes).
Your task: To determine the average atomic mass of Beanium.
Procedure:
1.
2.
3.
4.
5.
6.
7.
Obtain a sample of Beanium atoms.
Construct a data table to record the following:
Determine the total mass of Beanium atoms in your sample.
Sort the atoms into groups. Each group represents a different isotope.
Determine the mass of each isotope.
Record the total number of each isotope in your sample.
Find a property that you can use to identify each isotope.
Analysis:
Atomic masses are determined by comparing the mass of an atom with a standard mass,
the mass of a carbon-12 atom. In this experiment, the atoms are so big that we must use
another standard – the mass of 1mL of water or the gram. Make all measurements in
grams.
Determine the mass of each isotope in the following manner.
A. Take the total mass of each isotope.
B. Divide the total mass of each isotope by the number of atoms of that isotope.
C. Determine the percentage abundance of each isotope by dividing the number of
atoms of the isotope by the total number of atoms in the sample you received.
D. Determine the average atomic mass using the following formula:
Atomic mass =
% of isotope (1) x mass of 1 atom of isotope (1) +
% of isotope (2) x mass of 1 atom of isotope (2) +
% of isotope (3) x mass of 1 atom of isotope (3)
Conclusion:
Based on your individual research what is the average atomic mass of Beanium?
Obtain the results from all of your classmates and compute the class average:
Application:
1. Neon 20 has a mass of 19.9924 amu, Neon 21 has a mass of 20.9940 amu, and
Neon 22 has a mass of 21.9914 amu. The abundance of Neon 20 is 90.92%,
Neon 21 is 0.257% and Neon 22 is 8.882%.
Calculate the average atomic mass of Neon, expected to be found on the
periodic table.
2. What is the atomic mass of silver if 51.83% of the silver atoms occurring in
nature have a mass of 106.905 amu and 48.17% of the atoms have a mass of
108.905 amu.
Show all work.