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Atomic Structure (Chapter 4)
4.1
An Atom is the basic unit of an element and is commonly defined as the smallest particle of an
element that retains its identity in a chemical reaction.
Democritus (460 B.C.- 370 B.C.) believed that atoms were indivisible and indestructible, but his ideas
were limited because they didn’t explain chemical behavior and they lacked experimental support.
Atoms of element A are identical. Atoms of element B are identical, but differ from element A,
Atoms of element A and B can physically mix together (left), and atoms of element A and B can combine
chemically to form a compound (right).
How big are atoms? A pure copper penny contains about 2.4 x 1022 atoms of copper. If 1 billion atoms
were lined up side to side they would be about a 1 cm long.
Most atoms have a radius of 5 x 10-11 m to 2 x1 0-10 m. Despite their small size, individual atoms are
observable with instruments such as scanning tunneling microscopes. Ex. Iron atoms are the cones
Atomic models – In order to understand the properties of matter and how substances combine with one
another, we use an atomic model based on very strong experimental evidence.
Models shortcomings: It provides no information about the internal structure of the atom, and thus,
offers no insight as to how and why certain atoms of certain elements combine with other atoms
Section 4.2 Structure of the Nuclear Atom - Subatomic particles - turns out that atoms are divisible and can
be broken down into smaller particles called subatomic particles: protons, neutrons, and electrons
Ultimately, his work, along with the work of Robert Millikan (1909) made it possible to determine the
mass and charge of an electron
m = 9.1094 x 10-31 Kg charge = - 1.6022 x 10-19 Coulombs
From the Cathode ray tube experiments, William Thompson proposed that an atom consisted of a
An illuminating discovery in 1896 by Henri Becquerel eventually lead to the ultimate understanding of the
Experimental set up of the Au foil experiment
Some alpha particles were deflected, suggesting that the gold atoms had a very dense positively
charged center, which Rutherford called the nucleus.
In 1911, the nuclear model (or Rutherford model) of the atom was proposed based on this Nobel Prize
winning work - the discovery of the atomic nucleus
It held that:
1. Most of the volume of the atom is empty space (If a nucleus was 1 cm big, the 1st e- would be 1 Km away)
2. Most of the mass of the atom is concentrated in a dense, positively charged nucleus (very dense - size of a
pea = 250 million tons)
3. Electrons are present in the space surrounding the nucleus
The Modern model of the atom
All atoms are composed of subatomic particles called nucleons (protons and neutrons) and electrons
p+
no
e1. Nucleons are the extremely dense particles that make up the nucleus: 2 types
a. Protons – have a positive electrical charge (+1) equal in magnitude to the charge of an electron (-1)
And have a relative mass of 1 atomic mass unit (amu)
b. Neutrons - virtually have the same mass as a proton, but have no electrical charge
2. Electrons – move around the nucleus in electron orbitals, and have a relative charge of -1, but are
virtually "massless" compared to protons and neutrons (nucleons). e- mass = 1/1840 of an amu
The table below summarizes the properties of the subatomic particles
Due to the extremely small mass, the electron masses are not considered when adding up the mass of an
atom.
Question going forward: If all atoms are composed of the same components (subatomic particles), why do
different atoms have different properties?
**Electrons constitute most of the atomic volume and are the particles that "intermingle" when atoms
combine to form compounds or molecules
Therefore, it is the number of electrons possessed by an atom that determines an atom's properties.
Since neutral atoms have no net charge, the number of electrons must be equal to the number of protons
e- = p+ (in a neutral atom)
Section Review Questions
1. What are the main ideas of Dalton's atomic theory?
2. According to Dalton's theory, is it possible to convert atoms of one element into atoms of another?
3. A sample of copper with a mass of 63.5 g contains 6.02 x 1023 atoms. Calculate the mass of a single
copper atom?
4. What are the three types of subatomic particles? What are their relative charges and relative masses?
5. How does the Rutherford model of the atom describe the structure of atoms.
6. What experimental evidence led Rutherford to conclude that the atom was mostly empty space?
7. What experimental evidence led Rutherford to conclude that the atom contain a dense, positively charged
nucleus?
4.3 Distinguishing Among Atoms
Atoms are represented by the following shorthand standard notation.
Mass number
A
X
Element symbol
Z
Atomic number
Atomic # (Z) – is equal the atom's # of protons, and determines the atom's identity.
Since the proton is the only nucleon with a charge, the atomic number (Z) is also equal to
the nuclear charge ex) 126 C nuclear charge = +6
At this time, it should be noted that the entire periodic table is arranged by increasing (Z) atomic
number (or number of protons)
Mass # (A) – is equal to the sum of protons and neutrons and therefore, neutrons (no) = A-Z
Identify the following atoms, and indicate the # of protons, neutrons, electrons, and nuclear charge
Note: the mass numbers are staggered to the left
and should not be
Ex.
121
51Sb
Ex.
201
80Hg
Ex.
39
19K
Ex.
64
29Cu
Ex.
119
50Sn
Ex.
184
74W
Ex.
23
11Na
Ex.
56
26Fe
Standard notation (not shorthand)
(write the name of the element followed
by the mass number)
Contrary to Dalton’s belief, not all atoms of the same element are identical, often times, they have
different mass numbers (which means they have different numbers of neutrons)
Isotopes - Atoms with the same # of protons, but different numbers of neutrons are called isotopes
Ex. Consider the isotopes of Sodium: Sodium-23 and Sodium-24
Although there are two different types of Sodium, both have identical chemical properties because,
they have the same number of electrons and protons.
Hydrogen has three isotopes, which all have identical chemical properties. Draw the shorthand standard
notation and list how many protons neutrons and electrons each isotope contains
Hydrogen –1
Hydrogen –2
Hydrogen –3
Neon has three as well, which all have identical chemical properties, draw the shorthand standard notation of
each below
From looking at the periodic table, we can surmise which isotope is most abundant in nature. For instance
the atomic mass of hydrogen on the periodic table is 1.0079 and since this number rounds to 1 rather than 2
or 3, we can deduce that hydrogen-1 is the most common isotope.
The most abundant isotope of Sodium is
What about Chlorine?
Carbon?
Measuring Atomic Mass - we use a mass spectrometer, and every element's mass was
determined relative to a neutral atom of carbon-12 or 126C
Detector
plate
03_34
Ion-accelerating
electric field
Accelerated
ion beam
Least massive ions
Positive ions
Sample
Most
massive
ions
Electron beam
Slits
Magnetic field
Heating device
to vaporize sample
From comparing every atom to the Carbon-12 standard, the designation amu (atomic mass unit) was derived
1 amu is equal to one twelfth (1/12) the mass of the Carbon-12 atom, and thus the mass of a proton is
considered to be 1 amu and the mass of a neutron is considered to be 1 amu while electrons are considered
to be "massless".
The entire periodic table is therefore based on the amu of C-12, ultimately, the periodic table was
ordered by atomic number (Z) (not atomic mass due to some discrepancies)
The values for atomic mass on the periodic table represent the average of all isotopes for the element at
naturally occurring percentages.
To calculate the atomic mass of an element, multiply the mass of each isotope by its natural
abundance, expressed as a decimal, and then add the products.
Ex.
Calculate the atomic mass of Carbon
98.89% of Carbon is C-12 (amu of 12.00)
1.11% of Carbon is C-13 (amu of 13.0)
Ex. Determine the average atomic mass of an element that contain 2 isotopes:
Isotope A has an amu of 10.0130 and its natural abundance is 19.9 %
Isotope B has an amu of 11.0093 and its natural abundance is 80.10 %
What is the average atomic mass of the element?
What element is it?
Consider the table below, and check the weighted averages with your calculator
Later in Stoichiometry (the math of chemistry), this weighted average of naturally occurring isotopes for
each element (called atomic mass) becomes extremely useful and forms the base for the practice of
chemistry – It gives us a way to accurately measure quantities of elements and compounds for chemical
reactions.
Periodic Table Preview
The periodic table arranges elements so that they are separated into groups based on a set of repeating
properties
period - each horizontal row in the periodic table. There are 7 periods and the number of elements and
their properties very significantly in each period or row
group - vertical columns in the periodic table (also called families). There are 18 groups and elements in
the same group or family have similar chemical and physical properties. For instance, Na and K are
both explosive metals in water (chemical property) with densities less than 1.0 g/mL (physical property)
Section review questions
1. Express the following in shorthand notation and indicate the number of protons, neutrons, and electrons
in each atom
neon-20
sulfur-32
silver-108
bromine-80
lead-207
uranium-238
2. Three isotopes of oxygen are oxygen-16, oxygen-17, and oxygen-18. Express the following in shorthand
notation and list the number of neutrons contained in each atom and each atoms nuclear charge.
3. The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative
abundance and atomic masses are 69.2% for mass = 62.93 amu, and 30.8% for mass = 64.93 amu. Calculate
the average atomic mass of copper.
4. Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative
abundance of 78.92 amu(50.69%) and 80.92 amu (49.31%).
5. What distinguishes the atoms of one element from atoms of another?
6. What equation tells you how to calculate the number of neutrons in an atom?
7. How do the isotopes of a given element differ from one another?
8. How is atomic mass calculated?
9. The atomic masses of elements are generally not whole numbers. Explain why.
10. Name two elements that have properties similar to those of calcium (Ca).