Module 3 (Part 1)

Subject CHM012 Chemistry for Engineers

Chapter/Unit 3
Lesson Title Atomic Structure and the Periodic Table
Timeframe Week 3
Lesson Objectives 1. Compare the Bohr model, Rutherford model and the Quantum
mechanical model of the atom
2. Explain the four quantum numbers
3. Elucidate the quantum mechanical model of the atom
4. Give the electronic configuration of an element
5. Compare similarities and contrast differences of element within a
group/period in the periodic table
Overview/Introduction Although the materials in our world vary greatly in their properties,
everything is formed from only about 100 elements and, therefore,
from only about 100 chemically different kinds of atoms. In a sense,
the atoms are like the 26 letters of the English alphabet that join in
different combinations to form the immense number of words in our
language. But what rules govern the ways in which atoms combine?
How do the properties of a substance relate to the kinds of atoms it
contains? Indeed, what is an atom like, and what makes the atoms of
one element different from those of another? In this chapter we
examine the basic structure of atoms and discuss the formation of
molecules and ions, thereby providing a foundation for exploring
chemistry more deeply in later chapters.
Activity Exercises/Assignment
Assessment Problem Set
References Brown, T.L., LeMay Jr., H.E., Bursten, B.E., Murphy, C.J., Woodward, P.M.,
“Chemistry – The Central Science”, (14th edition), Prentice-Hall International,
Inc. (Chapter 1)

THE ATOMIC THEORY OF MATTER

 Greek Philosophers: Can matter be subdivided into fundamental particles?
 Democritus (460–370 BC): All matter can be divided into indivisible atomos.
 Later, however, Plato and Aristotle formulated the notion that there can be no ultimately indivisible
particles, and the “atomic” view of matter faded for many centuries
 The notion of atoms reemerged in Europe as chemists learned to measure the amounts of elements that
reacted with one another to form new substances, the ground was laid for an atomic theory that linked
the idea of elements with the idea of atoms
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Dalton: Proposed atomic theory with the following postulates:

Atoms are the building blocks of matter.

Dalton’s theory explains several laws of chemical combination that were known during his time

Laws of Chemical Combination

 Law of constant composition: (POSTULATE 4) In a given compound, the relative numbers and kinds of
atoms are constant
 Law of conservation of mass: (POSTULATE 3) The total mass of materials present after a chemical
reaction is the same as the total mass present before the reaction
 Law of multiple proportions: (deduce by Dalton) If two elements A and B combine to form more than
one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole
numbers

We can illustrate this law by considering water and hydrogen peroxide, both of which consist of the elements
hydrogen and oxygen. In forming water, 8.0 g of oxygen combine with 1.0 g of hydrogen. In forming hydrogen
peroxide, 16.0 g of oxygen combine with 1.0 g of hydrogen. Thus, the ratio of the mass of oxygen per gram of
hydrogen in the two compounds is 2:1. Using Dalton’s atomic theory, we conclude that hydrogen peroxide
contains twice as many atoms of oxygen per hydrogen atom as does water.

Water: H2O Hydrogen Peroxide: H2O2

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Structure of the Atom

Dalton based his conclusions about atoms on chemical observations made in the laboratory.

The Discovery of Atomic Structure

 By 1850 scientists knew that atoms consisted of charged particles.
 Subatomic particles: those particles that make up the atom.
 Before we summarize the current model, we briefly consider a few of the landmark discoveries that led to
that model.
 As we discuss the development of our current model of the atom, keep in mind this fact:
The law of electrostatic attraction: like charges repel and opposite charges attract.

Cathode Rays and Electrons

 Cathode rays first discovered in mid-1800s from studies of electrical discharge through partially evacuated
tubes (cathode ray tubes or CRTs).
 Cathode rays are radiations produced when high voltage is applied across the tube.
 The voltage causes negative particles to move from the negative electrode (cathode) to the positive
electrode (anode).
 Experiments showed that cathode rays are deflected by electric or magnetic fields in a way consistent with
their being a stream of negative electrical charge

J.J. Thomson (1897)

- demonstrated that cathode rays are composed of tiny, negatively charged subatomic particles called
electrons.

 Consider cathode rays leaving the positive electrode through a small hole.
 If they interact with a magnetic field perpendicular to an applied electric field, then the cathode rays can
be deflected by different amounts.
 The amount of deflection of the cathode rays depends on the applied magnetic and electric fields.
 The path of the electrons can be altered by the presence of a magnetic field.
 In turn, the amount of deflection will depend on the charge-to-mass ratio of the electron.
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 Using this knowledge, J.J. Thomson determined the charge-to-mass ratio of an electron.
o Charge-to-mass ratio: 1.76 x 108 C/g.
o C is a symbol for coulomb.
o SI unit for electric charge.

Robert Millikan (1909)

 measuring the charge of an electron using the Millikan Oil-Drop Experiment.
 Goal: Find the charge on the electron in order to determine its mass.

Millikan Oil-Drop Experiment:

 Oil drops are sprayed above a positively charged plate containing a small hole.
 As the oil drops fall through the hole they acquire a negative charge.
 Gravity forces the drops downward. The applied electric field forces the drops upward.
 When a drop is perfectly balanced, then the weight of the drop is equal to the electrostatic force of
attraction between the drop and the positive plate.
 Millikan carried out the above experiment and determined the charges on the oil drops to be multiples of
1.60 x 10–19 C.
 He concluded the charge on the electron must be 1.60 x 10–19 C.
 Knowing the charge to mass ratio of the electron, we can calculate the mass of the electron:

1.60 10 19 C

Mass   9.10  10 28 g
1.76  10 8 C/g

SET-UP OF MILLIKAN’S OIL DROP EXPERIMENT

Radioactivity
 French Scientist, Henri Bequerel discovered the uranium emits high energy radiation.
 Radioactivity is the spontaneous emission of radiation.
 At Becquerel’s suggestion, Marie Sclodowska Curie and her husband, Pierre, began experiments to
isolate the radioactive components of the compound.
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Ernest Rutherford
Further study of radioactivity, principally by the British scientist Ernest Rutherford revealed three types of
radiation

 Consider the following experiment:

o A radioactive substance is placed in a lead shield containing a small hole so that a beam of radiation
is emitted from the shield.
o The radiation is passed between two electrically charged plates and detected.

 Three spots are observed on the detector:

1. A spot deflected in the direction of the positive plate.
2. A spot that is not affected by the electric field.
3. A spot deflected in the direction of the negative plate.

 A large deflection towards the positive plate corresponds to radiation that is negatively charged and of low
mass. This is called -radiation (consists of electrons).
 No deflection corresponds to neutral radiation. This is called -radiation (similar to X-rays).
 A small deflection toward the negatively charged plate corresponds to high mass, positively charged
radiation. This is called -radiation (positively charged core of a helium atom)
 X-rays and  radiation are true electromagnetic radiation, whereas - and -radiations are actually streams
of particles--helium nuclei and electrons, respectively.

Eugen Goldstein (late 1880s)

– canal rays are composed of positively charged subatomic particle → protons (p+)

And decades later:

James Chadwick (1935), Nobel Prize winner for his discovery → neutron (n) = neutral subatomic particle
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Exercise 1

1. Describe a major contribution to science made by each of the following scientists:

a. Dalton

b. Thompson

c. Millikan

d. Rutherford

2. Hydrogen sulfide is composed of two elements: hydrogen and sulfur. In an experiment, 6.500 g of
hydrogen sulfide is fully decomposed into its elements.

a. If 0.384 g of hydrogen are obtained in this experiment, how many grams of sulfur must be obtained?
b. What fundamental law does this experiment demonstrate?

The Nuclear Model of the Atom

1. JJ Thompson Model of the Atom (1900)

 The plum pudding model: an early picture of the atom. This model was very short-lived
o The Thomson model pictures the atom as a sphere with small electrons embedded in a positively
charged mass called protons.

THOMSON’S PLUM PUDDING MODEL

2. Rutherford Model of the Atom (1910)

 Rutherford with the help of an undergraduate student, Ernest Marsden, carried out the following “gold foil”
experiment:
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 A source of -particles was placed at the mouth of a circular detector.

 The -particles were shot through a piece of gold foil.
o Both the gold nucleus and the -particle are positively charged, so they repel each other.

Outer Region of
Negative Charge

Central Positive
Charge

RUTHERFORD’S GOLD FOIL EXPERIMENT

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Observations:
 Most of the -particles went straight through the foil without deflection.
 Some -particles deflected or bounce back
 If the Thomson model of the atom was correct, then Rutherford’s result was impossible.
 Rutherford modified Thomson’s model as follows:
1. Assume the atom is spherical, but the positive charge must be located at the center with a diffuse
negative charge surrounding it.
2. In order for the majority of -particles that pass through a piece of foil to be undeflected, the
majority of the atom must consist of a low mass, diffuse negative charge - the electron.
3. To account for the small number of large deflections of the -particles, the center or nucleus of the
atom must consist of a dense positive charge – the proton.

THE MODERN VIEW OF ATOMIC STRUCTURE

 The atom consists of positive, negative and neutral entities (protons, electrons and neutrons).
 Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the
atom is due to the nucleus.
 Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.
 The quantity 1.602 x 10–19 C is called the electronic charge. The charge on an electron is –1.602 x 10–19
C; the charge on a proton is +1.602 x 10–19 C; neutrons are uncharged.
o Atoms have an equal number of protons and electrons thus they have no net electrical charge.
o Masses are so small that we define the atomic mass unit, amu.
 1 amu = 1.66054 x 10–24 g.
o The mass of a proton is 1.0073 amu, a neutron is 1.0087 amu, an electron is 5.486 x 10–4
amu.
o The angstrom is a convenient non-SI unit of length used to denote atomic dimensions.
o Since most atoms have radii around 1 x 10–10 m, we define 1 Å = 1 x 10–10 m.

Exercise 2

1. The diameter of a US dime is 17.9 mm, and the diameter of a silver atom is 2.88 Å. How many silver
atoms could be arranged side by side across the diameter of a dime?
2. The diameter of a carbon atom is 1.54 Å. (a) Express this diameter in picometers. (b) How many
carbon atoms could be aligned side by side across the width of a pencil line that is 0.20 mm wide?

Solution:
1. 6.22 x 107 Ag atoms
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Properties of Protons, Neutrons, and Electrons

Subatomic Particle Charge Location Mass (amu)
proton +1 inside nucleus 1.00728
neutron 0 inside nucleus 1.00866
electron -1 outside nucleus 0.00055

Atomic Numbers, Mass Numbers and Isotopes

What makes an atom of one element different from an atom of another element is that the atoms of each
element have a characteristic number of protons.
 Atomic number (Z) = number of protons in the nucleus.
 Because an atom has no net electrical charge, the number of electrons it contains must equal
the number of protons
o Example: All atoms of carbon, for example, have six protons and six electrons, whereas all atoms of
oxygen have eight protons and eight electrons.

 Atoms of a given element can differ in the number of neutrons they contain and, consequently, in mass.
 Mass number (A) = total number of nucleons in the nucleus (i.e. protons and neutrons).
o By convention, for element X, we write:

A
Z
X
Example for (read as “carbon twelve,” carbon-12)

 Because all atoms of a given element have the same atomic number, the subscript is redundant and is
often omitted
 Atoms with identical atomic numbers but different mas numbers are called Isotopes. Isotopes of a
specific element differ in the number of neutrons. Thus isotopes have the same Z but different A.
 There can be a variable number of neutrons for the same number of protons. Isotopes have the same
number of protons but different numbers of neutrons.
o Example: Some isotopes of Carbon
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 The nucleus of an atom (containing protons and neutrons) remains unchanged after ordinary chemical
reactions, but atoms can readily gain or lose electrons.
 Ions – are species formed when an atom either losses or gains electrons.

loss of gain of
Neutral
electron(s) electron(s)
atom

Cation Anion
(+) (-)

Exercise 3
(Use a Periodic Table)
1. How many protons, neutrons, and electrons are in (a) an atom of 197Au (b) an atom of strontium-90?
2. How many protons, neutrons, and electrons are in (a) a 138Ba atom, (b) an atom of phosphorus-31?
3. How many protons, neutrons, and electrons are in the following atoms:
a. 40Ar
b. 55Mn
c. 65Zn
d. 79Se
e. 235U

4. Complete the following table:

Symbol Atomic # Mass # # of Electrons # of Protons # of Neutrons Net Charge

23
11 Na
11 10 12
80 35 1-

Solution:
1. a) 79 protons, 79 electrons, 118 neutrons; b) 38 protons, 38 electrons, 52 neutrons

Assignment
Exercise 1 (#1, #2)
Exercise 2 (#2)
Exercise 3 (#2, #3, #4)