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Chapter 4 Atomic Structure
I.
II.
Before Atoms
The accepted belief was that all things were made up of 4
Elements – Fire, wind, earth, and water
People believed in what they could see and they could not see
atoms
Atoms
– Discovery and study of atoms done by indirect observations,
logic, and scientific deduction
Democtritus – 4 century B.C. in Greece 1st to suggest atoms were tiny indivisible particles that made
up matter (philosopher)
2000 years later Dalton proposed atomic theory (performed experimental science)
Summarize Dalton’s atomic theory states:
Elements made up of submicroscopic indivisible particles called atoms
Atoms of same element identical
Atoms of different elements are different
Atoms of different elements can physically mix together
Atoms can chemically combine in simple whole-number ratios to form compounds
Chemical reactions occur when atoms are separated, joined, or rearranged – atoms do not change in
chemical reaction
An atom is the smallest particle of an element that retains the properties of that element
Atoms are so small that we could only see outlines with an STM for many years – But now we can
actually see individual atoms with atomic force microscopes. Look at STM and article AFM
Most of atomic theory accepted today – atoms are divisible
World’s strongest microscope powers up for the first time
Jan. 23, 2008
TEAM 0.5, the world's most powerful transmission electron microscope — capable of producing images with
half-angstrom resolution (half a ten-billionth of a meter), less than the diameter of a single hydrogen atom — has
been installed at the Department of Energy's National Center for Electron Microscopy (NCEM) at Lawrence
Berkeley National Laboratory.
Atom by atom in 3-D
In preliminary tests at the FEI Company, before the TEAM 0.5 was shipped, NCEM's Christian Kisielowski tested
the microscope's ability to resolve individual atoms and precisely locate their positions in three dimensions. He
made a series of images of two gold crystals connected by a "nanobridge" only a few dozen atoms wide. From
each exposure to the next, individual gold atoms could be seen changing positions.
Read Article on Atomic Force Microscope in Notes
Concept practice page 86 - 1 and 2
1. Democritus and Dalton both proposed that matter consists of atoms. Explain how their approaches to
reaching the same conclusion differed.
2. Which of these statements would John Dalton have agreed with? Use Dalton's law to explain your answer.
a. Atoms are the smallest particles of matter.
b. The mass of an iron atom is different from the mass of a copper atom.
c. Every atom of silver is identical to every other atom of silver.
d. A compound is composed of atoms of two or more different elements.
III.
Electrons, protons, and neutrons
Electrons, protons and neutrons are the three basic subatomic particles that make up atoms
1
There are other particles, but in chemistry we will only study the 3 basic particles
Electrons smallest particle 1/2000 the size of the other two, has negative charge,
Thomson discovered electrons using the cathode ray tube (p86)
Read famous experiment cathode ray notes
http://dev.physicslab.org/Document.aspx?doctype=3&filename=Magnetism_CathodeRays.xml
The second subatomic unit is positively charged particle called the proton – much more mass than the electron
Goldstein discovered the proton using cathode ray tube using canal rays – then his family went into furniture!
http://chemed.chem.purdue.edu/genchem/history/goldstein.html
The third primary particle the neutron was discovered by Chadwick
http://dev.physicslab.org/Document.aspx?doctype=3&filename=AtomicNuclear_ChadwickNeutron.xml
Neutrons are subatomic particles with no charge and have a mass about equal to the proton
Concept practice page 89, 3 and 4
Demo 4.2 cathode ray
3. Since all atoms have negatively charged electrons, shouldn't every sample of matter have a negative
charge? Explain.
4. What experimental evidence did Thomson have for the following ideas?
a. Electrons have a negative charge.
b. Atoms of all elements contain electrons.
IV.
Structure of an atom
It was first thought that the subatomic particles were evenly distributed throughout the atom
Rutherford fired alpha particles at gold foil to test theory he concluded that most of the mass of the
atom is in the dense center (nucleus) and the rest of the atom is filled with empty space
http://micro.magnet.fsu.edu/electromag/java/rutherford/
The nucleus is the central core of the atom composed of the protons and neutrons (very dense, most
mass, pea=250 tons)
The nucleus carries a positive charge and occupies a very small part of the volume of an atom
The small negative electrons move around the rest of the atom, which makes most of the atoms
volume empty space
If the atom was the size of the earth (radius 6000km) the nucleus would have a radius of 60 m
Average radius of an atom is 1e-8cm and a volume of 4e-24cm3
http://www.newton.dep.anl.gov/askasci/phy00/phy00599.htm
Average of 2.4e22 atoms in a penny and the world population is 6.7e9 people 6.7 billion
Demo 4.3 page 91 Concept practice page 92 - 5 and 6
5. How did the results of Rutherford's gold foil experiment differ from his expectations?
6. What is the charge, positive or negative, of the nucleus of every atom?
V.
Atomic number identifies an element (pass out tables)
The atomic number of an element is the number of protons in the nucleus of the atom of that element
The number of protons of an atom determines the kind of atom it is Ex C all have 6 protons O all have
8 protons
The atomic number identifies the atom of the element
Atoms (elements) are classified according to the number of protons in the nucleus – or atomic number
Atoms are electrically neutral because the (+1) protons and the (-1) electrons cancel each other out
The protons are equal to the electrons in a neutral atom, so the atomic number is equal to the number
of protons and electrons
Table 4.2 page 92
The periodic table is a listing of the elements, elements in the period table arranged according to their
atomic number
Page 93 Example 1, Concept Practice, Practice problem
Example:
The element nitrogen (N) is atomic number 7. How many protons and electrons are in a nitrogen atom?
The atomic number equals the number of protons or the number of electrons in an atom. Since the atomic
number is 7, a nitrogen atom has seven protons and seven electrons.
Practice Problems
7. Why is an atom electrically neutral?
8. What is the relationship between the number of protons and the
atomic number of an atom?
9. Use the periodic table to complete this table.
2
Page 5
Element
Symbol
Atomic number
Number of protons
Potassium
5
16
Y
VI.
Mass Number
The atomic mass number of an atom is equal to the number of protons plus the number of neutrons
Because these particles are found in the nucleus they are referred to as the nucleons
The composition of any atom can be determined by using:
Mass number = number of protons + number of neutrons
In some periodic tables the mass number is written as the superscript (it is always larger) and the
atomic number is written as a subscript (it is always smaller)
Page 94 example 2
How many protons, electrons, and neutrons are in the following atoms?
Atomic number
Mass number
a. Beryllium (Be)
4
9
b. Neon (Ne)
10
20
c. Sodium (Na)
11
23
Complete this table. , page 95 practice problem 10
Atomic
Mass
Number of
Number of
number
number
protons
neutrons
9
Number of
electrons
Symbol of
element
10
14
15
47
55
22
25
Page 95 example 3 and
How many neutrons are in the following atoms?
Recall that the superscript is the mass number and the subscript is the atomic number. The mass number minus
the atomic number equals the number of neutrons.
3
concept practice 11
11. An atom is identified as platinum-195.
a. What is the number 195 called?
b.
Write the symbol for this atom using superscripts and subscripts.
12. Determine the number of neutrons in each atom.
a. carbon-13
b. nitrogen-15
c. radium-226
** Ions – when an atom becomes charged by gaining or losing an
electron.
Gain electrons – becomes negative charge = to number gained
Lose electron – becomes positive charge = to number lost
VII.
Isotopes of an element
Atoms of the same element must have the same number of protons and electrons, but they can have
different numbers of neutrons
Isotopes are atoms of an element that differ by the number of neutrons in the nucleus
Isotopes have different mass numbers
Isotopes of an atom are chemically the same since the electrons and protons, not the neutrons,
determine the atom’s chemical properties – Radioactive isotopes – C-12 and C-14
Most elements have at least one naturally occurring isotope
There are about 1000 naturally occurring isotopes
Tin Sn has the most isotopes at 10
Page 96 example 4,
Two of the isotopes of carbon are carbon-12 and carbon-13. Give the chemical symbol for each.
The mass numbers are given in the names of the isotopes. Carbon is atomic number 6. All atoms of carbon
have six protons.
concept practice 13, practice problems 14 and 15
13. How are isotopes of the same element alike? How are they different?
14. Three isotopes of oxygen are oxygen-16, oxygen-17, and oxygen-18. Write the chemical symbol, including the
atomic number and mass number, for each.
15. Use Table 4.3 to determine the number of protons, electrons, and neutrons in each of the five isotopes of zinc.
VIII.
Atomic Masses
-
The mass of the largest atom is about 10-24g, which is too small of a unit to work with, so
when using atomic mass relative mass of atoms are used
How mass of atoms were first calculated
http://antoine.frostburg.edu/chem/senese/101/atoms/faq/atomic-masses-withoutmasspec.shtml
Mass spectrometer now used to measure mass of atoms
http://www.asms.org/whatisms/p1.html
The atomic mass unit (amu) is defined as 1/12 the mass of a carbon – 12 atom (element is named
by its name and atomic mass number)
C – 12 has 6 protons and 6 neutrons so the mass of a single proton or neutron is 1 amu
Most mass numbers in the periodic table are not whole numbers because in nature most elements
occur as a mixture of two or more isotopes – and Mass defect
The atomic mass of an element is the weighted average mass of the atoms in a naturally occurring
sample of the element
A weighted average mass reflects both the mass and the relative abundance of the isotopes in nature
Page 100 example 5 Copper has two isotopes: copper-63 and copper-65. Given that the atomic mass
of copper from the periodic table is 63.546 amu, which of the isotopes of copper is most abundant?
concept practice 18 and 19
18. What data must you have about the isotopes of an element to be able
to calculate the atomic mass of the element?
4
19. There are three isotopes of silicon with mass numbers of 28, 29, and
30. The atomic mass of silicon is 28.086 amu. Comment on the
relative abundance of these three isotopes.
IX.
Calculate the average atomic mass of an element from isotope data
To calculate average atomic mass of an element three things must be known
Number of stable isotopes of that element
Mass of each element
Natural percent of abundance of each isotope
Mass Defect – problem 20 – Cu – 63 actual amu 62.93 – it should 63 because it is all one isotope small amount of mass lost to energy when protons and neutrons combined in nucleus
Page 101 Example 6,
Element X has two natural isotopes. The isotope with mass 10.012 amu has a relative abundance of 19.91%. The
isotope with mass 11.009 has a relative abundance of 80.09%. Calculate the atomic mass of this element and
name it.
Solution
Find the mass that each isotope contributes to the weighted average by multiplying the mass by its relative
abundance. Then add the products.
10X
10.012 amu x 0.1991 = 1.993 amu UX
11.009 amu x 0.8009 = 8.817 amu Total 10.810 amu
Element X is boron, atomic number 5.
Practice problem 20
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