ATOMIC STRUCTURE
INTRODUCTION
May seem surprising, but questioning what things are made of is as common today as it was among the philosophers of
ancient Greece, even though we approach the question very differently. They believed everything was made of one
or, at most, a few elemental substances (elements). Some believed the elemental substance was water because rivers and
oceans extend everywhere. Others thought it was air, which was ―thinned‖ into fire or ―thickened‖ into clouds, rain, and
rock. Still others believed there were four elements—fire, air, water, and earth—whose properties accounted for hotness,
wetness, sweetness, and all other characteristics of things. Democritus (c. 460–370 BC), the father of atomism, took a
different approach. He focused on the ultimate components of all substances and his reasoning went something like this: if
you cut a piece of, say, copper smaller and smaller, you must eventually reach a particle of copper so small that it can no
longer be cut. Therefore, matter is ultimately composed of indivisible particles, with nothing between them but empty
space.
GENERAL OBJECTIVES
By the end of the chapter, the students shall be able to:
1. Recognize that atomic models are used to explain atoms and help us understand the interaction of elements and
compounds observed on a macroscopic scale.
2. Recognize discoveries from Dalton (atomic theory), Thompson (the electron), Rutherford (the nucleus) and Bohr
(planetary model of atom) and understand how these discoveries lead to the modern atomic theory.
3. understand how periodic properties like atomic radii, ionization energy and electron affinity change with atomic
number and principal quantum number
4. Show understanding of both IUPAC and the ways of grouping elements on the periodic table.
PARTICULATE NATURE OF MATTER
The entire universe is made of two things: matter and energy. Matter is defined as anything which has mass and occupies
space. All matter is particulate in nature. This basically means that between separate bits of matter there are spaces which
contain no matter. Most of the matter in our world is solid, liquid or gas. These are called the three states of matter‘. Most
of the time, we can easily tell which is which: • this book is solid; • sea-water is liquid; • air is a gas
In a solid, the particles are touching each other and they fit together in a regular way. Because there are no big spaces
between particles, the density of the solid is high. Solids in which the ‗packing‘ is very good (the metals) have the highest
densities of all. Because there is very little space between the particles, a solid cannot be compressed by squeezing.
Squeezing or hammering a solid can change its shape (think of hammering a metal or a piece of glass or squeezing some
putty) but its volume doesn‘t change. Because a solid normally keeps its own shape we know that the particles in a solid
must be fixed in place with respect to their neighbours. There are strong bonds between them. Once in position, there the
particles stay, but they are not completely still. Each particle vibrates to and fro. If we heat the solid (add heat energy to it)
the vibrations get stronger, and if we cool a solid (take heat energy out of it) the vibrations get weaker. When we say that
something is ‗hot‘ we mean that its particles are vibrating strongly. A substance in which the particles are only vibrating
weakly is ‗cold‘.
Liquids In a liquid the particles are almost as close together as in a solid, but there is no order in the way they are arranged.
The particles in a liquid are slightly further apart than they are in a solid, so a liquid usually has a slightly lower density
than its solid and the solid sinks in its own liquid. (Water is an exception – ice floats on water.) The spaces between the
particles in a liquid are not very large and so a liquid can only be compressed a very little. If liquids could be compressed a
lot then hydraulic systems such as car brakes could not work. As we all know, liquids do not keep their own shape – they
flow. This means that the particles can move relative to each other and so they can change neighbours freely. But as it
flows, a liquid keeps the same volume. Its particles stay close together. Like the particles in a solid, the particles in a liquid
have strong bonds between them.
Gases Gases have much lower densities than either liquids or solids. You can see this in action when bubbles rise through
fizzy drinks! Gases can be compressed into small fractions of their ordinary volumes (think of pumping up a bicycle tyre).
Gases can be compressed because the particles are far apart from each other. Now imagine opening a bottle of perfume at
one end of a room – it will not be long before you can smell it at the other end of the room. This tells us that the forces
between the gas particles are weak, because if they were strong the particles would stick together. Instead, the particles are
free to move anywhere in the room. If we could watch the movement of any of them, we would see something very like
Brownian motion. The gas particles would move about randomly, in straight lines, only changing direction when they
bumped into one another, into air particles or bounced off the walls of the room.
Atoms, molecules and ions are the building blocks of everything you see around : the papers you are looking at, your study
table, your books, etc. Such is the amazing power of nature and fundamental nature of these particles. Despite the discovery
of sub-particles like electrons, protons and neutrons, an atom continues to remain the fundamental particle because of the
fact that it is the smallest unit that exhibits the chemical properties of an element.
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Chemistry is essentially the study of matter and the changes it undergoes in everyday activities like cooking to more
complex processes such as photosynthesis. In essence, the heart of chemistry involves studying changes around our world.
Let us discuss some basic ideas of atoms, molecules and the matter they make--elements and compounds.
Atoms
As stated earlier, an atom is the smallest constituent particle of an element which exhibits the chemical properties of an
element and also can take part in a chemical reaction. Atoms are extremely small and their sizes are about an angstrom
(10-10m). Although visualizations are nowadays made by very sophisticated microscopes, no one has actually seen an atom.
Initial studies showed that it is spherical in shape and indivisible, but later it was found that atoms have been subdivided
into simpler parts. These simpler parts are called the subatomic particles. The subatomic particles include the quarks and
the leptons: Atoms join together to form molecules and when atoms gain or lose electrons ions are formed.
Matter can be classified into three types based on its composition—elements, compounds, and mixtures.
An element is a substance that cannot be broken down into simpler substances by normal chemical means. An element is
composed of atoms that have the same atomic number, that is, each atom has the same number of protons in its nucleus as
all other atoms of that element such as silicon, oxygen or copper. A sample of silicon contains only silicon atoms. A key
point to remember is that the macroscopic properties of a piece of silicon, such as colour, density, and combustibility, are
different from those of a piece of copper because silicon atoms are different from copper atoms; in other words, each
element is unique because the properties of its atoms are unique. Most elements exist in nature as populations of atoms.
However, several elements occur naturally as molecules:
Molecule is the smallest physical unit of an element or compound, consisting of two or more similar atoms in an element
and two or more different atoms in a compound that can participate in a chemical reaction. Elemental oxygen, for
example, occurs in air as diatomic (two-atom) molecules.
A compound is a type of matter composed of two or more different elements that are chemically bound together. Be sure
you understand that the elements in a compound are not just mixed together; rather, their atoms have joined chemically.
Ammonia, water, and carbon dioxide are some common compounds.
Atomicity is the total number of atoms present in one molecule of an element or compound. For example in O2 there are
two atoms of oxygen so the atomicity of O 2 is 2(diatomic).
Monoatomic elements are single atoms existing as a unit and not bound to one another. Usually the noble gases such as
helium and argon are said to be monoatomic.
Homonuclear molecules are molecules composed of only one type of element. Homonuclear molecules may consist of
various numbers of atoms, depending on the element's properties. Some elements form molecules of more than one size.
Noble gases rarely form bonds, so they only have one atom. The most familiar homonuclear molecules are diatomic,
meaning they consist of two atoms, though not all diatomic molecules are homonuclear. Homonuclear diatomic molecules
include hydrogen (H2), oxygen (O2), nitrogen (N2) and all of the halogens. Ozone (O3) is a common triatomic homonuclear
molecule. Homonuclear tetratomic molecules include arsenic (As4) and phosphorus (P4).
Heteronuclear molecules consist of atoms different elements two different elements. There is an abundance of
heteronuclear molecules. Two types of bonding can occur in a heteronuclear diatomic molecule: ionic and covalent.
Ionic bonding occurs when a metal bonds with a nonmetal. Covalent bonding occurs when two nonmetals bond together.
The following are examples of heteronuclear diatomic molecules with ionic bonding and they are normally called ionic
compounds.
 NaCl - Sodium Chloride
 MgO - Magnesium Oxide
 KBr - Potassium Bromide
 CaO - Calcium Oxide
The following are examples of heteronuclear diatomic molecules with covalent bonding.
 HCl - Hydrochloric Acid
 CO - Carbon Monoxide
 ClF - Chlorine Monofluoride
 HBr - Hydrobromic Acid
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Evidence Of Particulate Nature Of Matter
Robert Brown (1773–1858), a Scottish botanist, found good evidence for particulate nature of matter when he was using a
microscope to look at pollen grains in water. He saw that the grains were always moving – but in random, zig-zag paths.
Brown suggested that the water was made of particles which were bombarding the pollen grains and knocking them about.
We can see the same zig-zag motion if we look at smoke particles in the air through a microscope. This phenomenon is
called Brownian motion and its one of the evidence of particulate nature of matter.
The following provide evidence to support the theory that all matter is made of particles:
1. Crystallization (formation of Crystals)
2. Dissolution of crystals
3. Diffusion
4. Brownian motion
5. Osmosis
Diffusion
When you walk past a cosmetics counter in a department store you can usually smell the perfumes. For this to happen gas
particles must be leaving open perfume bottles and be spreading out through the air in the store. This spreading out of a gas
is called diffusion and it takes place in a haphazard and random way. Diffusion also takes place in liquids but it is a much
slower process than in gases. This is because the particles of a liquid move much more slowly.
Brownian motion
Brownian motion is the random movement of microscopic particles in a fluid, as a result of continuous bombardment from
molecules of the surrounding medium. Robert Brown (1773–1858), a Scottish botanist, found good evidence for this when
he was using a microscope to look at pollen grains in water. He saw that the grains were always moving – but in random,
zig-zag paths. Brown suggested that the water was made of particles which were bombarding the pollen grains and
knocking them about. We can see the same zig-zag motion if we look at smoke particles in the air through a microscope. It
is called Brownian motion.
Crystallization
Crystallization is the process whereby the atoms or molecules of highly organized solid structure known as crystal is
formed out of solution. Crystals can form through precipitating from a solution, melting or more rarely deposition directly
from a gas. The crystals formed from the solution were tiny particles within the solution and this supports particulate nature
of matter. Crystallization is also a chemical solid–liquid separation technique, in which mass transfer of a solute from the
liquid solution to a pure solid crystalline phase occurs. A typical laboratory technique for crystal formation is to dissolve
the solid in a solution in which it is partially soluble, usually at high temperatures to obtain super saturation. The hot
mixture is then filtered to remove any insoluble impurities. The filtrate is allowed to slowly cool. Crystals that form are
then filtered and washed with a solvent in which they are not soluble, but is miscible with the mother liquor. The process is
then repeated to increase the purity in a technique known as recrystallization.
Questions
1. Define the following terms: Atom, Molecule, element, atomicity
2. State two evidence of particulate nature of matter
3. Give one example each of diatomic, triatomic and tetra-atomic molecules
DALTON'S ATOMIC THEORY
Dalton expressed his theory in a series of postulates. Like most great thinkers, Dalton incorporated the ideas of others into
his own to create the new theory. As we go through the postulates, which are presented here in modern terms, let‘s see
which were original and which came from others. (Later, we can examine the key differences between Dalton‘s postulates
and our present understanding.)
Postulates of Dalton’s Atomic Theory
The main points of Dalton‘s atomic theory, an explanation of the structure of matter in terms of different combinations of
very small particles, are given by the following postulates:
1. Elements are made up of small particles called atoms
2. Atoms can neither be created nor destroyed.
3. Atoms of the same elements are identical –they have the same mass and size
4. Atoms of different elements have different mass and size
5. Atoms combine to form compounds and they do so in simple whole numbers.
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DISCOVERY OF THE ELECTRON
Fortunately, we can acquire a qualitative understanding of atomic structure without having to retrace all the
discoveries that preceded atomic physics. We do, however, need a few key ideas about the interrelated phenomena of
electricity and magnetism, which we briefly discuss here. Electricity and magnetism were used in the experiments that led
to the current theory of atomic structure. Certain objects display a property called electric charge, which can be either
positive or negative. Positive and negative charges attract each other, while two positive or two negative charges repel each
other. All objects of matter are made up of charged particles. An object having equal numbers of positively and negatively
charged particles carries no net charge and is electrically neutral. If the number of positive charges exceeds the number of
negative charges, the object has a net positive charge. If negative charges exceed positive charges, the object has a net
negative charge.
Sometimes when one substance is rubbed against another, as in combing hair, net electric charges build up on the
objects, implying that rubbing separate some positive and negative charges. Moreover, when a stationary (static) positive
charge builds up in one place, a negative charge of equal size appears somewhere else; charge is balanced. They are
deflected from their straight-line path into a curved path in a plane perpendicular to the field. Think of the field or region of
influence of the magnet as represented by a series of invisible lines of force running from the North Pole to the south pole
of the magnet.
CRT, the abbreviation for cathode-ray tube, was once a familiar acronym. Before liquid crystal display (LCD) was
available, the CRT was the heart of computer monitors and TV sets. The first cathode-ray tube was made by Michael
Faraday (1791-1867) about 150 years ago. When he passed electricity through glass tubes from which most of the air had
been evacuated, Faraday discovered cathode rays, a type of radiation emitted by the negative terminal, or cathode. The
radiation crossed the evacuated tube to the positive terminal, or anode. Later scientists found that cathode rays travel in
straight lines and have properties that are independent of the cathode material (that is, whether it is iron, platinum, and so
on). The cathode rays produced in the CRT are invisible, and they can be detected only by the light emitted by materials
that they strike. These materials, called phosphors example zinc sulphide, are painted on the end of the CRT so that the path
of the cathode rays can be revealed. (Fluorescence is the term used to describe the emission of light by a phosphor when it
is struck by energetic radiation.) Another significant observation about cathode rays is that they are deflected by electric
and magnetic fields in the manner expected for negatively charged particles.
The cathode ray tube experiment (J.J. Thompson’s experiment)
In 1897 the British physicist J. J. Thompson conducted a series of experiments that showed that atoms were not indivisible
particles by using cathode ray tube. In this apparatus, two electrodes from a high-voltage source are sealed into a glass tube
from which the air has been evacuated. The negative electrode is called the cathode; the positive one, the anode. When the
high-voltage current is turned on, cathode rays leave the negative electrode, they move toward the anode, where some rays
pass through a hole to form a beam. This beam bends away from the negatively charged plate and toward the positively
charged plate. (Cathode rays are not directly visible, but do cause certain materials such as zinc sulphide to glow so you can
observe them). In a similar experiment, in which cathode rays are seen to bend when a magnet is brought toward them.
Thompson showed that the characteristics of cathode rays are independent of the material making up the cathode. From
such evidence, he concluded that a cathode ray consists of a beam of negatively charged particles (or electrons) and that
electrons are constituents of all matter
Cathode rays and their properties
(a) Deflection of cathode rays in an electric field. The beam of cathode rays is deflected as it travels from left to right in the
field of the electrically charged condenser plates. The deflection corresponds to that expected of negatively charged
particles.
(b) Deflection of cathode rays in a magnetic field. The beam of cathode rays is deflected as it travels from left to right in
the field of the magnet. The deflection corresponds to that expected of negatively charged particles.
(c) Determining the mass-to-charge ratio, for cathode rays. The cathode-ray beam strikes the end screen undeflected if the
forces exerted on it by the electric and magnetic fields are counterbalanced. By knowing the strengths of the electric and
magnetic fields, together with other data, a value of can be obtained. Precise measurements yield
a value of per coulomb. (Because cathode rays carry a negative charge, the sign of the mass-to-charge ratio is also negative.
In 1897, by the method outlined, J. J. Thompson (1856 1940) established the ratio of mass (m) to electric charge
(e) for cathode rays. Also, Thompson concluded that cathode rays are negatively charged fundamental particles of matter
found in all atoms. (The properties of cathode rays are independent of the composition of the cathode.) Cathode rays
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subsequently became known as electrons, a term first proposed by George Stoney in 1874.
Robert Millikan (1868-1953) determined the electronic charge e through a series of oil-drop experiments (1906-1914). The
currently accepted value of the electronic charge e, expressed in coulombs to five significant figures, is -1.6022 x 10-19 C.
By combining this value with an accurate value of the mass-to-charge ratio for an electron, we find that the mass of an
electron is 9 .1094 x 10-28 g.
Once the electron was seen to be a fundamental particle of matter found in all atoms, atomic physicists began to
speculate on how these particles were incorporated into atoms. The commonly accepted model was that proposed by J. J.
Thompson. Thompson thought that the positive charge necessary to counterbalance the negative charges of electrons in a
neutral atom was in the form of nebulous cloud. Electrons, he suggested, floated in a diffuse cloud of positive charge
(rather like a lump of gelatin with electron fruit embedded in it). This model became known as the plum-pudding model
because of its similarity to a popular English dessert.
The Nuclear Atom
In 1909, Rutherford, with his assistant Hans Geiger, began a line of research using particles as probes to study the inner
structure of atoms. Based on Thompson‘s plum-pudding model, Rutherford expected that most particles in
a beam of particles would pass through thin sections of matter largely undeflected, but that some particles would be slightly
scattered or deflected as they encountered electrons. By studying these scattering patterns, he hoped to
deduce something about the distribution of electrons in atoms. The apparatus used for these studies is pictured in the figure
above. Alpha particles were detected by the flashes of light they produced when they struck a
zinc sulphide screen mounted on the end of a telescope. When Geiger and Ernst Marsden, a student, bombarded very thin
foils of gold with particles, they observed the following:
--The majority of particles penetrated the foil undeflected.
-- Some particles experienced slight deflections.
-- A few (about 1 in every 20,000) suffered rather serious deflections as they penetrated the foil.
A similar number did not pass through the foil at all, but bounced back in the direction from which they had come. The
large-angle scattering greatly puzzled Rutherford. As he commented some years later, this observation was about as
credible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you. By 1911, though,
Rutherford had an explanation. He based his explanation on a model of the atom known as the nuclear atom and having
these features:
1. Most of the mass and all of the positive charge of an atom are centered in a very small region called the nucleus.
The remainder of the atom is mostly empty space.
2. The magnitude of the positive charge is different for different atoms and is approximately one-half the
atomic weight of the element.
3. There are as many electrons outside the nucleus as there are units of positive charge on the nucleus.
The atom as a whole is electrically neutral.
Discovery of Protons and Neutrons
Rutherford‘s nuclear atom suggested the existence of positively charged fundamental particles of matter in the nuclei of
atoms. Rutherford himself discovered these particles, called protons, in 1919 in studies involving the scattering
of particles by nitrogen atoms in air. The protons were freed as a result of collisions between particles and the nuclei of
nitrogen atoms. At about this same time, Rutherford predicted the existence in the nucleus of electrically neutral
fundamental particles. In 1932, James Chadwick showed that a newly discovered penetrating radiation consisted of beams
of neutral particles. These particles, called neutrons, originated from the nuclei of atoms. Thus, it has been
only for about the past 100 years that we have had the atomic model suggested.
Properties of the three key subatomic particles
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Name
Proton
Neutron
Electron
Relative charge
+1
0
-1
Relative mass
1
1
1/1840
Location in atom
Inside nucleus
Inside nucleus
Outside nucleus
Properties of Cathode Rays
 They travel in straight lines from the negative pole (cathode).
 They produce fluorescence in the glass walls of the discharge tube.
 They cast shadows if some target is placed in their path-travel in straight line.
 They can produce mechanical motion, e.g., they cause a light pedal wheel placed in their path to rotate.
Questions
1. Describe the structure of the atom according to
a.
J. J. Thompson
b.
Rutherford
2. Identify the scientist associated with the following discoveries
a.
Proton
b.
Neutron
c.
Electron
3. State the charges associated with the following particles
a.
Proton
b.
Neutron
c.
Electron
QUANTUM MECHANICS
The early twentieth century revolutionized how we think about physical reality- before that time, all descriptions of matter
had been deterministic—the present completely determining the future. Quantum mechanics changed that. This new theory
suggested that for subatomic particles—electrons, neutrons, and protons—the present does NOT completely determine the
future. For example, if you shoot one electron down a path and measure where it lands, a second electron shot down the
same path under the same conditions will most likely land in a different place! Several gifted scientists, such as Albert
Einstein, Niels Bohr, Louis de Broglie, Max Planck, Werner Heisenberg, P. A. M. Dirac, and Erwin Schrödinger,
developed quantum-mechanical theory. Their new theory, however, made even some of them uncomfortable. Bohr said,
―Anyone who is not shocked by quantum mechanics has not understood it.‖ Schrödinger wrote, ―I don‘t like it, and I‘m
sorry I ever had anything to do with it.‖ Albert Einstein disbelieved it stating, ―God does not play dice with the universe.‖
In fact, Einstein attempted to disprove quantum mechanics—without success—until he died. Today, quantum mechanics
forms the foundation of chemistry—explaining the periodic table and the behaviour of the elements in chemical bonding.
In this section, we examine the quantum-mechanical model of the atom, a model that explains the strange behavior
of electrons. In particular, we focus on how the model describes electrons as they exist within atoms, and later we shall see
how those electrons determine the chemical and physical properties of elements.
THE NATURE OF LIGHT
Before we explore electrons and their behavior within the atom, we must understand some of the properties of light. As
quantum-mechanical theory was developed, light was (surprisingly) found to have many characteristics in common with
electrons. Chief among these characteristics is the wave–particle duality of light. Certain properties of light are best
described by thinking of it as a wave, while other properties are best described by thinking of it as a particle. In this section,
we first explore the wave behavior of light, and then its particle behavior. We then turn to electrons to see how they display
the same wave–particle duality.
The Wave Nature of Light
Light is electromagnetic radiation, a type of energy embodied in oscillating electric and magnetic fields. A magnetic field is
a region of space where a magnetic particle experiences a force (think of the space around a magnet). An electric field is a
region of space where an electrically charged particle experiences a force. A proton, for example, has an electric field
around it. If you bring another charged particle into that field, that particle experiences a force.
The Particle Nature of Light
Prior to the early 1900s, and especially after the discovery of the diffraction of light, light was thought to be purely a wave
phenomenon. Its behaviour was described adequately by classical electromagnetic theory, which treated the electric and
magnetic fields that constitute light as waves propagating through space. However, a number of discoveries brought the
classical view into question. Chief among these was the photoelectric effect. The photoelectric effect is the observation that
many metals emit electrons when light shines upon them, In 1905, Albert Einstein proposed a bold explanation for the
photoelectric effect: Light energy must come in packets. According to Einstein, the amount of energy (E) in a light packet
depends on its frequency. Einstein‘s idea that light is quantized elegantly explains the photoelectric effect. The emission of
electrons from the metal depends on whether or not a single photon has sufficient energy to dislodge a single electron.
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Atomic Spectroscopy and the Bohr Model
The discovery of the particle nature of light began to break down the division that existed in nineteenth-century physics
between electromagnetic radiation, which was thought of as a wave phenomenon, and the small particles (protons,
neutrons, and electrons) that compose atoms, which were thought to follow Newton‘s laws of motion. Just as the
photoelectric effect suggested the particle nature of light, so certain observations about atoms began to suggest a wave
nature for particles. The most important of these observations came from atomic spectroscopy, the study of the
electromagnetic radiation absorbed and emitted by atoms.
Atomic Spectra
When an atom absorbs energy—in the form of heat, light, or electricity—it often re-emits that energy as light. For
example, a neon sign is composed of one or more glass tubes filled with neon gas. When an electric current is
passed through the tube, the neon atoms absorb some of the electrical energy and re-emit it as the familiar red
light of a neon sign. If the atoms in the tube are different (that is, not neon), they emit light of a different colour. In
other words, atoms of each element emit light of a characteristic colour. Mercury atoms, for example, emit light
that appears blue, helium atoms emit light that appears violet, and hydrogen atoms emit light that appears reddish.
Closer inspection of the light emitted by various atoms reveals that it contains several distinct wavelengths. We
can separate the light emitted by a single element in a glass tube into its constituent wavelengths by passing it through a
prism (just like we separate the white light from a light bulb. The result is a series of bright lines called an emission
spectrum. Emission spectrum is the emission of electromagnetic radiation by atoms or other species resulting from
electronic transitions from higher to lower energy states. The emission spectrum of a particular element is always the
same—it consists of the same bright lines at the same characteristic wavelengths and hence called line spectrum. We can
use it to identify the element. Line spectrum is a spectrum that consists of narrow, brightly coloured parallel lines on a
dark background. For example, light arriving from a distant star contains the emission spectra of the elements that compose
the star. Analysis of the light allows us to identify the elements present in the star. There is difference between the white
light spectrum and the emission spectra of specific elements such as hydrogen, helium and barium. The white light
spectrum is continuous, meaning that there are no sudden interruptions in the intensity of the light as a function of
wavelength—the spectrum consists of light of all wavelengths. The emission spectra of hydrogen, helium, and barium,
however, are not continuous—they consist of bright lines at specific wavelengths, with complete darkness in between. That
is, only certain discrete wavelengths of light are present. Classical physics could not explain why these spectra consisted of
discrete lines. In fact, according to classical physics, an atom composed of an electron orbiting a nucleus should emit a
continuous white light spectrum. Even more problematic, the electron should lose energy as it emits the light and spiral into
the nucleus. According to classical physics, an atom should not even be stable.
THE BOHR MODEL
The Danish physicist Niels Bohr (1885–1962) attempted to develop a model for the atom that explained atomic spectra. In
his model, electrons travel around the nucleus in circular orbits (analogous to those of the planets around the sun).
However, in contrast to planetary orbits—which can theoretically exist at any distance from the sun—Bohr‘s orbits exist
only at specific, fixed distances from the nucleus. The energy of each Bohr orbit is also fixed, or quantized. Bohr called
these orbits stationary states and suggested that, although they obey the laws of classical mechanics, they also possess ―a
peculiar, mechanically unexplainable, stability.‖ We now know that the stationary states were really manifestations of the
wave nature of the electron, which we will expand upon shortly. Bohr further proposed that, in contradiction to classical
electromagnetic theory, no radiation is emitted by an electron orbiting the nucleus in a stationary state. It is only when an
electron jumps, or makes a transition, from one stationary state to another that radiation is emitted or absorbed.
The transitions between stationary states in a hydrogen atom are quite unlike any transitions that we might be
familiar with in the macroscopic world. The electron is never observed between states; it is observed only in one state or
another. The emission spectrum of an atom consists of discrete lines because the stationary states exist only at specific,
fixed energies. The energy of the photon emitted when an electron makes a transition from one stationary state to another is
the energy difference between the two stationary states. Transitions between stationary states that are closer together,
therefore, produce light of lower energy (longer wavelength) than transitions between stationary states that are farther apart.
In spite of its initial success in explaining the line spectrum of hydrogen, the Bohr model left many unanswered
questions. It did, however, serve as an intermediate model between a classical view of the
electron and a fully quantum-mechanical view, and therefore has great historical and conceptual importance. Nonetheless,
it was ultimately replaced by a more complete quantum-mechanical theory that
fully incorporated the wave nature of the electron.
ATOMIC SPECTROSCOPY AND THE IDENTIFICATION OF ELEMENTS
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When you check out of the grocery store, a laser scanner reads the barcode on the items that you buy. Each item has a
unique code that identifies the item and its price. Similarly, each element in the periodic table has a spectrum unlike that of
any other element. Each spectrum is unique and, as such, can be used to identify the substance. The presence of intense
lines in the spectra of a number of metals is the basis for flame tests, simple tests used to identify elements in ionic
compounds in the absence of a precise analysis of a compound‘s spectrum. For example, the emission spectrum of sodium
features two closely spaced, bright yellow lines. When a crystal of a sodium salt (or a drop of a solution containing a
sodium salt) is put into a flame, the flame glows bright yellow. Other metals exhibit similarly characteristic colours in
flame tests. Each colour represents an especially bright spectral emission line (or a combination of two or more such lines).
Similar emissions form the basis of the colours seen in fireworks.
Although the emission of light from elements is easier to detect, the absorption of light by elements is even more
commonly used for purposes of identification. Whereas an emission spectrum consists of bright lines on a dark
background, an absorption spectrum consists of dark lines on a bright background. An absorption spectrum is measured by
passing white light through a sample and observing what wavelengths are missing due to absorption by the sample. Note
that, in the spectra of mercury, the absorption lines are at the same wavelengths as the emission lines. This is because the
processes that produce them are mirror images. In emission, an electron makes a transition from a higher energy level to a
lower energy one. In absorption, the transition is between the same two energy levels, but from the lower level to the higher
one. Absorption spectrometers, found in most chemistry laboratories, typically plot the intensity of absorption as a function
of wavelength. Such plots are useful both for identifying substances (qualitative analysis) and for determining the
concentration of substances (quantitative analysis). Quantitative analysis is possible because the amount of light absorbed
by a sample depends on the concentration of the absorbing substance within the sample. For example, the concentration of
Ca2+ in a hard water sample can be determined by measuring the quantity of light absorbed by the calcium ion at its
characteristic wavelength.
The Wave Nature of Matter: The de BroglieWavelength, the uncertainty Principle, and Indeterminacy
The heart of the quantum-mechanical theory that replaced Bohr‘s model is the wave nature of the electron, first proposed
by Louis de Broglie (1892–1987) in 1924 and later confirmed by experiments in 1927. It seemed incredible at the time, but
electrons—which were then thought of only as particles and known to have mass—also have a wave nature. The wave
nature of the electron is seen most clearly in its diffraction.
 The de Broglie Wavelength
Louis de Broglie proposed that electron has the characteristics of both matter and wave. He concluded that single electron
traveling through space has a wave nature; its wavelength is related to its kinetic energy (the energy associated with its
motion). The faster the electron is moving, the higher its kinetic energy and the shorter its wavelength. The wavelength ( )
of an electron of mass m moving at velocity 𝒱 is given by the de Broglie relation:
𝒱
where h is Planck‘s constant. Notice that the velocity of a moving electron is related to its wavelength— knowing one is
equivalent to knowing the other.
•The uncertainty Principle
As we just saw in the de Broglie relation, the velocity of an electron is related to its wave nature. The position of an
electron, however, is related to its particle nature. (Particles have well-defined positions, but waves do not.) Consequently,
our inability to observe the electron simultaneously as both a particle and a wave means that we cannot simultaneously
measure its position and its velocity with infinite precision. Werner Heisenberg formalized this idea with the equation:
𝒱
where
is the uncertainty in the position, 𝒱 is the uncertainty in the velocity, m is the mass of the particle, and h is
Planck‘s constant.
In other words, Heisenberg uncertainty principle states that, the more accurately you know the position of an
electron, the less accurately you can know its velocity and vice versa. The complementarity of the wave nature and particle
nature of the electron results in the complementarity of velocity and position. Although Heisenberg‘s uncertainty principle
may seem puzzling, it actually solves a great puzzle. Without the uncertainty principle, we are left with a paradox: How can
something be both a particle and a wave? Saying that an object is both a particle and a wave is like saying that an object is
both a circle and a square—a contradiction. Heisenberg solved the contradiction by introducing complementarity—an
electron is observed as either a particle or a wave, but never both at once.
Quantum Mechanics and the Atom
8
As we have seen, the position and velocity of the electron are complementary properties—if we know one accurately, the
other becomes indeterminate. Since velocity is directly related to energy (recall that kinetic energy equals 1/2 mv2),
position and energy are also complementary properties—the more we know about one, the less we know about the other.
Many of the properties of an element, however, depend on the energies of its electrons.
In the following paragraphs, we describe the probability distribution maps for electron states in which the electron
has well-defined energy, but not well-defined position. In other words, for each of these states, we can specify the energy of
the electron precisely, but not its location at a given instant. Instead, the electron‘s position is described in terms of an
orbital, a probability distribution map showing where the electron is likely to be found. An orbital is therefore said to be
the volume of space in an atom where there is high probability of finding an electron. Since chemical bonding often
involves the sharing of electrons between atoms, the spatial distribution of atomic electrons is important to bonding.
A total of four quantum numbers are used to describe completely the movement and trajectories of each electron
within an atom. Each electron in an atom has a unique set of quantum numbers; according to the Pauli Exclusion
Principle, no two electrons can share the same combination of four quantum numbers. Quantum numbers are important
because they can be used to determine the electron configuration of an atom and the probable location of the atom's
electrons. Quantum numbers are also used to determine other characteristics of atoms, such as ionization energy and the
atomic radius.
The Four Electronic Quantum Numbers
Quantum numbers designate specific shells, subshells, orbitals, and spins of electrons. This means that they describe
completely the characteristics of an electron in an atom. There are a total of four quantum numbers: the principal quantum
number (n), the angular momentum quantum number (l), the magnetic quantum number (m l), and the electron spin
quantum number (ms).
 The Principal Quantum Number ( n)
The principal quantum number, n , designates the principal electron shell. Because n describes the most probable distance
of the electrons from the nucleus, the larger the number n is, the farther the electron is from the nucleus, the larger the size
of the orbital, and the larger the atom is. n can be any positive integer starting at 1, as n=1 designates the first principal
shell (the innermost shell). The first principal shell is also called the ground state, or lowest energy state. This explains why
n cannot be 0 or any negative integer, because there exists no atoms with zero or a negative amount of energy
levels/principal shells. When an electron is in an excited state or it gains energy, it may jump to the second principle shell,
where n=2. This is called absorption because the electron is "absorbing" photons or energy. Known as emission, electrons
can also "emit" energy as they jump to lower principle shells, where n decreases by whole numbers. As the energy of the
electron increases, so does the principal quantum number, e.g., n = 3 indicates the third principal shell, n = 4 indicates the
fourth principal shell, and so on.
 The Angular Momentum Quantum Number (l)
The angular momentum quantum number (l) tells us the ―shape‖ of the orbitals. The values of l depend on the value of the
principal quantum number, n. For a given value of n, l has possible integral values from 0 to (n -1). If n =1, there
is only one possible value of l; that is, l=0. If n=2, there are two values of l, given by 0 and 1. If n=3, there are three values
of /, given by 0, 1and 2. The value of l is generally designated by the letters s, p, d, . . . as follows:
l
0 1 2 3 4 5
s p d f g h
Name of orbital
Thus, if l=0, we have an s orbital; if l=1, we have a p orbital; and so on. The unusual sequence of letters (s, p, and d) has a
historical origin. Physicists who studied atomic emission spectra tried to correlate the observed spectral lines with the
particular energy states involved in the transitions. They noted that some of the lines were sharp; some were rather spread
out, or diffuse; and some were very strong and hence referred to as principal lines. Subsequently, the initial letters of each
adjective were assigned to those energy states. However, after the letter d and starting with the letter f (for fundamental),
the orbital designations follow alphabetical order. A collection of orbitals with the same value of n is frequently called a
shell. One or more orbitals with the same n and l values are referred to as a subshell. For example, the shell with n=2 is
composed of two subshells, l=0 and 1 (the allowed values for n=2). These subshells are called the 2s and 2p subshells
where 2 denotes the value of n, and s and p denote the values of l.
Question
If n = 7, what are the possible values of l?
Answer: Since l can be zero or a positive integer less than (n-1), it can have a value of 0, 1, 2, 3, 4, 5 or 6.
 The Magnetic Quantum Number (ml)
The magnetic quantum number (ml) describes the orientation of the orbital in space. Within a subshell, the value of ml
depends on the value of the angular momentum quantum number, l. For a certain value of l, there are (2l+1) integral values
of ml as follows: -l, (-l+1), . . ., 0, . . . (l-1), l
If l = 0, then ml = 0. If l =1, then there are three values of ml, namely -1, 0, and 1. If l=2, there are five values of ml, namely
-2, -1, 0, 1, and 2. The number of ml values indicates the number of orbitals in a subshell with a particular l value. To
conclude our discussion of these three quantum numbers, let us consider a situation in which n=2 and l=1. The values of n
and l indicate that we have a 2p subshell, and in this subshell we have three 2p orbitals (because there are three values of
ml, given by -1, 0, and 1)
9
Question
If n=3, and l=2, then what are the possible values of ml ?
Answer: Since ml must range from –l to +l, then ml can be: -2, -1, 0, 1, or 2.
 The Electron Spin Quantum Number ( ms )
Unlike n, l , and ml , the electron spin quantum number, ms‘ does not depend on another quantum number. It designates
the direction of the electron spin and may have a spin of +1/2, represented by ↑, or –1/2, represented by ↓. This means that
when ms is positive the electron has an upward spin, which can be referred to as "spin up." When it is negative, the
electron has a downward spin, so it is "spin down." The significance of the electron spin quantum number is its
determination of an atom's ability to generate a magnetic field or not.
2
2
Note: For a principal quantum number, n, the number of orbitals in it is n and the number of electrons in the orbitals is 2n
Question
List the possible combinations of all four quantum numbers when n=2, l=1, and ml=0.
Answer:
The fourth quantum number is independent of the first three, allowing the first three quantum numbers of two electrons to
be the same. Since the spin can be +1/2 or =1/2, there are two combinations: n = 2, l=1, ml = 0, ms = +1/2 and n=2, l=1, ml
= 0 ms = -1/2
Questions
1. What values of the angular momentum (l ) and magnetic quantum numbers(ml) are allowed for a principal quantum
number(n) of 3? How many orbitals exist for n = 3?
2. Specify the l and ml values for n = 4.
3. List the quantum numbers associated with all of the 5d orbitals. How many 5d orbitals exist?
SHAPE OF ORBITALS
An atomic orbital can also be represented by a geometrical shape that encompasses the volume where the electron is likely
to be found most frequently—typically, 90% of the time. For example, the 1s orbital can be represented as the threedimensional sphere. Each principal level with n=2 or greater contains three p orbitals (ml = -1, 0, +1). The p orbitals are not
spherically symmetric like the s orbitals, but have two lobes of electron density on either side of the nucleus and a node
located at the nucleus(dumpbell. The three p orbitals differ only in their orientation and are orthogonal (mutually
perpendicular) to one another. It is convenient to define an x-, y-, and z-axis system and then label each p orbital as px, py,
and pz. The 3p, 4p, 5p, and higher p orbitals are all similar in shape to the 2p orbitals, but they contain additional nodes
(like the higher s orbitals) and are progressively larger in size.
ELECTRON CONFIGURATIONS
Electron confguration is the distribution of electrons among the various atomic orbitals. Three basic principles are used to
govern the distribution of electrons and these are Aufbau rule, Pauli Exclusion Principle and Hund‘s rule of maximum
multiplicity.
 Aufbau Rule
Aufbau rule states that electrons are placed in atomic orbitals in increasing order of energy level. First Lower energy
orbitals occupy with electrons then higher energy ones. If an empty lower energy orbital is present, electron cannot move to
higher energy orbital. The energy of various orbitals is given below.
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p < 8s <.
 Pauli Exclusion Principle
The Pauli exclusion principle, which summarizes experimental observations, states that no two electrons in an atom can
have the same four quantum numbers. If one electron in an atom has the quantum numbers n=1, l =0, ml=0,
and ms= +1/2, no other electron can have these same quantum numbers. In other words, you cannot place two electrons
with the same value of ms in a 1s orbital.
To write an electron configuration for an element, we first find its atomic number from the periodic table—this number
equals the number of electrons. Then we use the order of filling to distribute the electrons in the appropriate orbitals.
Remember that each orbital can hold a maximum of two electrons.
10
Consequently,
•The s sublevel has only one orbital and can therefore hold only 2 electrons.
•The p sublevel has three orbitals and can hold 6 electrons.
•The d sublevel has five orbitals and can hold 10 electrons.
•The f sublevel has seven orbitals and can hold 14 electrons.
Questions
Write an electron configuration for each element. (a) Mg (b) P (c) Br (d) Al
Ans:
2 2
6 2
or
[Ne] 3s2
12Mg: 1s 2s 2p 3s
2 2
6 2
3
1s 2s 2p 3s 3p
or
[Ne] 3s23p3
15P:
2 2
6 2
6 2
10
5
or
[Ar] 4s23d 104p5
35Br: 1s 2s 2p 3s 3p 4s 3d 4p
2 2
6 2
1
or
[Ne] 3s23p1
13Al: 1s 2s 2p 3s 3p
Questions
Write electron configurations for each element. (a) Cl (b) Si (c) Sr (d) O
 Hund’s rule of maximum multiplicity
Hund’s rule states that when filling degenerate orbitals, electrons fill them singly first, with parallel spins before pairing
with opposite spins.
Note: degenerate orbitals are orbitals with the same energy level and shape.
Example 1: Nitrogen Atoms
Consider the correct electron configuration of
the nitrogen(Z=7)
atom: 1s2 2s2 2p3
Example 2: Oxygen Atoms
Next, consider oxygen (Z = 8) atom, the element after
nitrogen in the same period; its electron configuration
is: 1s2 2s2 2p4
Questions
Write the orbital diagram for sulphur and determine its number of unpaired electrons.
Ans:
Two unpaired electrons
Questions
Write the orbital diagram for each of the following and determine its number of unpaired electrons in it
a. Ar
c. S
b. C
d. K
ELECTRON CONFIGURATIONS, VALENCE ELECTRONS AND THE PERIODIC TABLE
Mendeleev arranged the periodic table so that elements with similar chemical properties lie in the same column. We can
begin to make the connection between an element‘s properties and its electron configuration by superimposing the electron
configurations of the first 18 elements onto a partial periodic table.
As we move to the right across a row (which is also called a period), the orbitals fill in the correct order. With each
subsequent row, the highest principal quantum number increases by one. Notice that as we move down a column,
the number of electrons in the outermost principal energy level (highest n value) remains the same. The key connection
between the macroscopic world (an element‘s chemical properties) and the particulate world (an atom‘s electronic
structure) lies in these outermost electrons. An atom‘s valence electrons are the most important in chemical bonding. For
main-group elements, the valence electrons are those in the outermost principal energy level.
11
For transition elements, we also count the outermost d electrons among the valence electrons (even though they are not in
an outermost principal energy level). The chemical properties of an element depend on its valence electrons, which are
instrumental in bonding because they are held most loosely (and are therefore the easiest to lose or share). We can now see
why the elements in a column of the periodic table have similar chemical properties: They have the same number of
valence electrons.
We distinguish valence electrons from all the other electrons in an atom, which we call core electrons. The core
electrons are those in complete principal energy levels and those in complete d and f sublevels. For example, silicon, with
the electron configuration 1s22s22p63s23p2 has four valence electrons (those in the n = 3 principal level) and ten core
electrons.
Questions
1. Write an electron configuration for phosphorus, carbon and Calcium.
2. Identify the valence electrons and core electrons in each atom.
 How the Electron Configuration of an Element Relates to Its Properties
The chemical properties of elements are largely determined by the number of valence electrons they contain. The
properties of elements are periodic because the number of valence electrons is periodic. Mendeleev grouped elements into
families (or columns) based on observations about their properties. We now know that elements in a family have the same
number of valence electrons. In other words, elements in a family have similar properties because they have the same
number of valence electrons.
Perhaps the most striking family in the periodic table is the column labeled 8A, known as the noble gases. The
noble gases are generally inert—they are the most unreactive elements in the entire period table. Why? Notice that each
noble gas has eight valence electrons (or two in the case of helium), and they all have full outer quantum levels. We do not
cover the quantitative (or numerical) aspects of the quantum-mechanical model in this book, but calculations of the overall
energy of the electrons within atoms with eight valence electrons (or two for helium) show that these atoms are particularly
stable. In other words, when a quantum level is completely full, the overall potential energy of the electrons that occupy
that level is particularly low. Recall that systems with high potential energy tend to change in ways that lower their
potential energy. Systems with low potential energy, on the other hand, tend not to
change—they are stable. Because atoms with eight electrons (or two for helium) have particularly low potential energy, the
noble gases are stable—they cannot lower their energy by reacting with other atoms or molecules.
We can explain a great deal of chemical behavior with the simple idea that elements without a
noble gas electron configuration react to attain a noble gas configuration. This idea works particularly well for main-group
elements. In this section, we first apply this idea to help differentiate between metals and nonmetals. We then apply the
idea to understand the properties of several individual families of elements. Lastly, we apply the idea to the formation of
ions.
Metals and Nonmetals
We can understand the broad chemical behaviour of the elements by superimposing one of the most general properties of
an element—whether it is a metal or nonmetal—with its outer electron configuration in the form of a periodic table. Metals
are opaque materials with good electrical and thermal conductivity and normally form ions by losing electrons. Metals lie
on the lower left side and middle of the periodic table and share some common properties: They are good conductors of
heat and electricity; they can be pounded into flat sheets (malleability); they can be drawn into wires (ductility); they are
often shiny; and most importantly they tend to lose electrons when they undergo chemical changes. For example, sodium is
among the most reactive metals. Its electron configuration is 1s22s22p63s1. Notice that its electron configuration is one
electron beyond the configuration of neon, a noble gas. Sodium can attain a noble gas electron configuration by losing that
one valence electron—and that is exactly what it does. When we find sodium in nature, we most often find it as Na +, which
has the electron configuration of neon (1s22s22p6). The other main-group metals in the periodic table behave similarly:
They tend to lose their valence electrons in chemical changes to attain noble gas electron configurations. The transition
metals also tend to lose electrons in their chemical changes, but they do not generally attain noble gas electron
configurations.
Non-metals are any of a number of chemical elements that form negative ions, have acidic oxides, and are generally poor
conductors of heat and electricity. Non-metals lie on the upper right side of the periodic table. The division between metals
and nonmetals is the zigzag diagonal line running from boron to astatine. Nonmetals have varied properties— some are
solids at room temperature, others are liquids or gases—but as a whole they tend to be poor conductors of heat and
electricity, and most importantly they all tend to gain electrons when they undergo chemical changes. Chlorine is among
the most reactive nonmetals. Its electron configuration is 1s 22s22p63s23p5. Notice that its electron configuration is one
electron short of the configuration of argon, a noble gas.
12
Chlorine can attain a noble gas electron configuration by gaining one electron—and that is exactly what it does. When we
find chlorine in nature, we often find it as Cl-, which has the electron configuration of argon (1s22s22p63s23p6). The other
nonmetals in the periodic table behave similarly: They
tend to gain electrons in chemical changes to attain noble gas electron configurations.
Many of the elements that lie along the zigzag diagonal line that divides metals and nonmetals are
metalloids and exhibit mixed properties. Several metalloids are classified as semiconductors because of their intermediate
(and highly temperature-dependent) electrical conductivity. A metalloid is any chemical element which has properties in
between those of metals and nonmetals, or that has a mixture of them. Our ability to change and control the conductivity of
semiconductors makes them useful in the manufacture of the electronic chips and circuits central to computers, cellular
telephones, and many other devices. The six commonly recognised metalloids are boron, silicon, germanium, arsenic,
antimony, and tellurium.
Question
Classify the following elements as metals, non-metals or metalloids: Ar, N, B, Be, F, Mg, P, C, Si, I
Families of Elements
We have already seen that the group 8A elements, called the noble gases, are mostly unreactive. The most familiar noble
gas is probably helium, used to fill buoyant balloons. Helium is chemically stable—it does not combine with other elements
to form compounds—and is therefore safe to put into balloons. Other noble gases are neon (often used in electronic signs),
argon (a small component of our atmosphere), krypton, and xenon.
The group 1A elements, called the alkali metals, all have an outer electron configuration of ns1.
Like sodium, a member of this family, the alkali metals have electron configurations that are one electron beyond a noble
gas electron configuration. In their reactions, they readily, and sometimes violently, lose the ns1 electron to form ions with a
1+ charge. A marble-sized piece of sodium, for example, explodes violently when dropped into water. Lithium, potassium,
and rubidium are also alkali metals and as such behave in a similar way.
The group 2A elements, called the alkaline earth metals, all have an outer electron configuration of ns2. They have
electron configurations that are two electrons beyond a noble gas configuration. In their reactions, they tend to lose the two
ns2 electrons—though not quite as violently as the alkali metals—to form ions with a 2+ charge. Calcium, for example,
reacts fairly vigorously when dropped into water but does not explode as dramatically as sodium. Magnesium (a common
low-density structural metal), strontium, and barium are other alkaline earth metals.
The group 7A elements, the halogens, all have an outer electron configuration of ns 2np5. Like chlorine, a member
of this family, their electron configurations are one electron short of a noble gas configuration. Consequently, in their
reactions with metals, they tend to gain one electron to form ions with a 1- charge. One of the most familiar halogens is
chlorine, a greenish-yellow gas with a pungent odour. Because of its reactivity, chlorine is used as a sterilizing and
disinfecting agent. Other halogens include bromine, a red-brown liquid that easily evaporates into a gas; iodine, a purple
solid; and fluorine, a pale-yellow gas.
The Formation of Ions
We have just seen that metals tend to form positively charged ions (cations) and nonmetals tend to form negatively charged
ions (anions). A number of main-group elements in the periodic table always form ions with a noble
gas electron configuration. Consequently, we can reliably predict their charges. As we have already seen, the alkali metals
tend to form cations with a 1+ charge, the alkaline earth metals tend to form ions with a 2+ charge, and the halogens tend to
form ions with a 1- charge. In each of these cases, the ions have noble gas electron configurations. This is true of the rest of
the ions. Nitrogen for example, has an electron configuration of 1s 22s22p3. The N3- ion has three additional electrons and an
electron configuration of 1s22s22p6, which is the same as the configuration
of neon, the nearest noble gas. Notice that, for the main-group elements that form cations with predictable charge, the
charge is equal to the group number. For main-group elements that form anions with predictable charge, the
charge is equal to the group number minus eight. Transition elements may form various ions with different charges.
Questions
Predict the charges of the monoatomic (single atom) ions formed by these main-group elements.
Al
S
O
Rb
F
N
Mg
Li
P
Electron Configurations and Magnetic Properties
An unpaired electron generates a magnetic field due to its spin. Consequently, an atom or ion that contains unpaired
electrons is attracted to an external magnetic field, and we say that the atom or ion is paramagnetic.
An atom or ion in which all electrons are paired is not attracted to an external magnetic field—it is instead slightly
repelled—and we say that the atom or ion is diamagnetic.
Questions
Write the electron configuration and orbital diagram for each ion and determine whether each is diamagnetic or
paramagnetic.
(a) Al3+
(b) S2(c) Fe3+
(d) Co2+
(e) N3(f) Ca2+
ISOTOPES: WHEN THE NUMBER OF NEUTRONS VARIES
13
All atoms of a given element have the same number of protons; however, they do not necessarily have the same number of
neutrons. Since neutrons have nearly the same mass as protons (1 amu), this means that—contrary to what John Dalton
originally proposed in his atomic theory—all atoms of a given element do not have the same mass. For example, all neon
atoms contain 10 protons, but they may contain 10, 11, or 12 neutrons. All three types of neon atoms exist, and each has a
slightly different mass.
Atoms with the same number of protons but different numbers of neutrons are called isotopes. Some elements, such as
beryllium (Be) and aluminium (Al), have only one naturally occurring isotope, while other elements, such as neon (Ne) and
chlorine (Cl), have two or more. The relative amount of each different isotope in a naturally occurring sample of a given
element is roughly constant. For example, in any natural sample of neon atoms, 90.48% of them are the isotope with 10
neutrons, 0.27% are the isotope with 11 neutrons, and 9.25% are the isotope with 12 neutrons.
These percentages are called the natural abundance of the isotopes. Each element has its own characteristic natural
abundance of isotopes. However, advances in mass spectrometry have allowed accurate measurements that reveal small but
significant variations in the natural abundance of isotopes for many elements.
The atomic number (Z) is the number of protons in the nucleus of each atom of an element.
The sum of the number of neutrons and protons in an atom is its mass number and is represented by the symbol A:
A = number of protons (p) + number of neutrons (n)
For neon, with 10 protons, the mass numbers of the three different naturally occurring isotopes are 20, 21, and 22,
corresponding to 10, 11, and 12 neutrons, respectively.
We symbolize isotopes using this notation:
where X is the chemical symbol, A is the mass number, and Z is the atomic number. Therefore, the symbols for the neon
isotopes are
,
and
. Notice that the chemical symbol, Ne, and the atomic number, 10, are redundant: If the
atomic number is 10, the symbol must be Ne. The mass numbers, however, are different for the different isotopes,
reflecting the different number of neutrons in each one.
A second common notation for isotopes is the chemical symbol (or chemical name) followed by a dash and the mass
number of the isotope.
In this notation, the neon isotopes are: Ne-20 Ne-21 Ne-22 or neon-20 neon-21 neon-22
Question
a. What are the atomic number (Z), mass number (A), and symbol of the chlorine isotope with 18 neutrons?
b. How many protons, electrons, and neutrons are present in an atom of
?
c. What are the atomic number, mass number, and symbol for the carbon isotope with 7 neutrons?
d. How many protons and neutrons are present in an atom of
?
•
Mass Spectrometry: Measuring the Mass of Atoms and Molecules
14
Mass spectrometry is an instrumental method for identifying the chemical constitution of a substance by means of the
separation of gaseous ions according to their differing mass and charge — called also mass spectroscopy
The masses of atoms and the percent abundances of isotopes of elements are measured using mass spectrometry, a
technique that separates particles according to their mass. In a mass spectrometer, such, the sample (containing the atoms
whose mass is to be measured) is injected into the instrument and vapourized. The vapourized atoms are ionized by an
electron beam—the electrons in the beam collide with the atoms, removing electrons and creating positively charged ions.
The ions are then accelerated into a magnetic field. When ions drift through a magnetic field, they
experience a force that bends their trajectory. The amount of bending depends on the mass (and charge) of the ions—the
trajectories of lighter ions are bent more than those of heavier ones (of the same charge). Finally, the ions strike a detector
and produce an electrical signal that is recorded. The result is the separation of the ions according to their mass, producing
a mass spectrum. A graph of the intensity of the detector signal versus particle atomic mass is called a mass spectrum.The
position of each peak on the x-axis indicates the mass of the isotope that was ionized, and the intensity (indicated by the
height of the peak) indicates the relative abundance of that isotope. The number of isotopes, masses of various isotopes and
the relative abundance of each isotope are some of the information obtained from the mass spectrum.
Information from the mass spectrum of an element is used to determine the relative atomic mass of the element
Mass spectrometry can be used to identify elements from the mass numbers of the isotopes e.g. in space probes
on another planet, such as Mars, the number of isotopes and their mass numbers will be the same as on Earth. However, the
relative abundances may well be different so the relative atomic mass will be different as well.
Mass spectrometry can also be used to determine the relative molecular mass of a molecule.
Atomic Mass
Mass of atom is called atomic mass. Since, atoms are very small consequently actual mass of an atom is very small. For
example the actual mass of one atom of hydrogen is equal to 1.673 x 10 -24 g. This is equal to
0.000000000000000000000001673 gram. To deal with such small number is very difficult. Thus for convenience, relative
atomic mass is used.
Carbon-12 is considered as unit to calculate atomic mass. Carbon-12 is an isotope of carbon. The relative mass of all atoms
are found with respect to C-12.
One atomic mass = 1/12 of the mass of one atom of C-12.
This means atomic mass unit = 1/12th of carbon-12
Thus atomic mass is the relative atomic mass of an atom with respect to 1/12 th of the mass of carbon-12 atom. ‗amu‘ is the
abbreviation of Atomic mass unit, but now it is denoted just by ‗u‘.
The atomic mass of hydrogen atom = 1u.
This means one hydrogen atom is 1 times heavier than 1/12th of the carbon atom.
The atomic mass of oxygen is 16 u, this means one atom of oxygen is 16 times heavier than 1/12 th of carbon atom.
Absolute mass or Actual atomic mass:
It is found that the actual atomic mass of a carbon-12 atom is equal to 1.9926×10−23g.
∴ 1u = 1.9926×10-23/12g
⇒ 1u = 1.6605×10−24 g
Thus by multiplying the relative atomic mass with 1.6605 × 10 -24 g we can get the absolute or actual mass of an atom.
Example:
Find the absolute mass oxygen (O). [O = 16u]
Solution:
The atomic mass of oxygen is 16u
We know that, 1u = 1.6605×10−24 g
Therefore, Absolute mass of oxygen
= 1.6605×10−24×16 g
= 2.6568×10-23 g
Question
15
1.
2.
The mass of an element is 3.8191×10-23 g, if the the mass of 1/12 of carbon (1 amu) is 1.660539040×10−24g determine
the relative atomic mass of the element.
An element is found to be 40 times the mass of carbon 1/12 of carbon. Determine the actual mass of the element.
Relative Atomic Mass(Ar): The Average Mass Of An Element’s Atoms
An important part of Dalton‘s atomic theory is that all atoms of a given element have the same mass. We learned that
because of isotopes, the atoms of a given element often have different masses, so Dalton was not completely correct. We
can, however, calculate an average mass—called the relative atomic mass—for each element. The atomic mass of each
element is listed directly beneath the element‘s symbol in the periodic table and represents the average mass of the isotopes
that compose that element, weighted according to the natural abundance of each isotope. For example, the periodic table
lists the atomic mass of chlorine as 35.45 amu. Naturally occurring chlorine consists of 75.77% chlorine-35 atoms (mass
34.97 amu) and 24.23% chlorine-37 atoms (mass 36.97 amu). We can calculate its atomic mass: Relative Atomic mass =
0.7577(34.97 amu) + 0.2423(36.97 amu) = 35.45 amu
Relative atomic mass of an element can be defined as the ratio of the average mass of an element to one twelfth of the mass
of an atom of carbon-12.
From the calculation, the relative atomic mass is also said to be the weighted average mass of an atom of an element
taking the mixture of isotopes into account. To calculate relative atomic mass, add together (mass number x
abundance/total abundance) for each isotope
Abundances can be given as percentages, fractions or may have to be worked out from the line heights on the spectrum.
where m- mass of isotope and h-natural abundance of isotope (percentage or fraction).
Question
Calculate the atomic mass of copper if copper-63 is 69.17% abundant and copper-65 is 30.83% abundant.
Ans:
Ar = 63.546
Questions
1. Strontium consists of four isotopes with masses of 84 (abundance 0.50%), 86 (abundance of 9.9%), 87 (abundance of
7.0%), and 88 (abundance of 82.6%). Calculate the relative atomic mass of strontium.
2. Titanium has five common isotopes: 46Ti (8.0%), 47Ti (7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%). What is the
average atomic mass of titanium?
3. Boron exists in two isotopes, boron-10 and boron-11. If the relative atomic mass of boron is 10.81, which isotope
should be more abundant?
4. Lithium-6 is 4% abundant and lithium-7 is 96% abundant. What is the average mass of lithium?
5. Iodine is 80% 127I, 17% 126I, and 3% 128I. Calculate the relative atomic mass of iodine.
6. The natural abundance for boron isotopes is 19.9 % 10B and 80.1 % 11B . Calculate boron‘s atomic mass.
7. Hydrogen is 99% 1H, 0.8 % 2H, and 0.2 % 3H. Calculate its relative atomic mass.
8. Magnesium has three naturally occurring isotopes with masses of 23.99 amu, 24.99 amu, and 25.98 amu and natural
abundances of 78.99%, 10.00%, and 11.01%, respectively. Calculate the atomic mass of magnesium.
9. Gallium has two naturally occurring isotopes: Ga-69 with a mass of 68.9256 amu and a natural abundance of 60.11%,
and Ga-71. Use the atomic mass of gallium from the periodic table to find the mass of Ga-71.
10. Chlorine exists as two isotopes, 35Cl and 37Cl. The relative atomic mass of chlorine is 35.45. Calculate the percentage
abundance of each isotope.
PEAKS OF MOLECULES
For elements you get a series of signals or ion peaks for each isotope present and the ratio of peak heights gives you the
relative proportion of each isotope in the element so that you can calculate the relative atomic mass of an element. This
'simple' spectra of mononuclear ions like [Na]+ is only true for non-molecular elements like metals or noble gases, but for
molecular elements like nitrogen or the halogens things are not so simple.
For larger e.g. organic molecules, things can be very complex indeed, as molecules fragment and many different ions can
be formed but you can get the relative molecular mass of a molecule by identifying what is called the molecular ion peak,
that is, when one electron is knocked of the molecule but the molecule retains its full molecular structure.
e.g. bbenzoic acid (Mr = 122) gives a molecular ion peak of m/z = 122, due to [C6H5COOH]+
Chlorine Example
16
The mass spectrum of chlorine is a good example of a molecular element whose mass spectra can be a bit tricky when first
encountered. Chlorine consists of two principal stable isotopes, chlorine-37 (~25% is 37Cl) and chlorine-35 (~75% is 35Cl).
Chlorine consists of Cl2 diatomic molecules, which may or may not split on ionisation, so how can we explain the presence
of five peaks and not just two for the two isotopes?
The result of the ionisation process and subsequent fragmentation of chlorine molecules is a series of 5 different mass
peaks from the various isotopic monatomic or molecular ion possibilities.
[37Cl37Cl]+ or [37Cl2]+ m/z = 74 (molecular ion)
[37Cl35Cl]+ m/z
= 72 (note that you must show the two isotopes separately in the molecular ion)
[35Cl35Cl]+ or [35Cl2]+ m/z = 70 (molecular ion)
[37Cl]+ m/z
= 37 (mononuclear ion, monoatomic fragment)
[35Cl]+ m/z
= 35 (mononuclear ion, monoatomic fragment)
The Mass Spectrum Of Bromine Br2
You get five peaks in the spectra of bromine molecules. For molecules completely atomised you get two peaks (m/z) of
almost equal height from [79Br]+ and [81Br]+ mononuclear ions.
Because its ~50% of each isotope, the relative atomic mass of bromine is ~80 and hence the equality of peaks 1 [79Br]+ and
2 [81Br]+ from the monoatomic ions from the fragmentation and ionisation of bromine molecules.
However, as with chlorine, molecular bromine is also ionised without fragmentation, giving rise to three more ion
permutations (3 more m/z peaks). [79Br79Br]+ (158), [79Br81]Br+ (160) and [81Br81Br]+ (162)
So the presence of all five peaks is explained in the mass spectrum of bromine, and, because you are dealing with millions
of randomised ionised atoms, the ratio of the two monoatomic peaks can be used to accurately determine the relative
atomic mass of bromine.
The data book quotes for the stable isotopes: 79Br (50.69%) and 81Br (49.31)
The ratio of the heights for the monatomic ions in the mass spectrum of bromine would 50.69 : 49.31 ~ 1 : 1 as observed.
Ar(Br) = (50.69 x 79) + (49.31 x 81) / 100 = 79.90
m/z 79Br 81Br The ratio of the 2nd set of peaks (3 to 5) can be readily explained with a simple probability
79
Br 158 160 table, and a bit simpler than the chlorine example!
81
Br 160 162 This assumes (for simplicity) that we have exactly 50% bromine-79 and 50% bromine-81
isotopes and how they might be combined in the molecular ions on a random basis.
The ratio of peak heights expected for m/z values of 158 : 160 : 162 would be 1 : 2 : 1 and this is what you observe in the
mass spectrum of bromine.
Questions
Predict the number of peaks for the following molecules in a mass spectrometer
1. Hydrogen molecules, H2
stable isotopes 1H, 2H
2. Nitrogen molecules, N2,
stable isotopes of 14N, 15N
3. Oxygen molecules, O2
stable isotopes 16O, 17O, 18O
End Of Chapter Review Questions
1. The space between proton and electron in hydrogen atom is
A. Absolutely empty
B. Full of electromagnetic radiation
C. Full of air D. Full of Ether
2. According to classical theory, the proposed circular path of an electron in Rutherford atomic model will be:
A. Circular
B. Straight line
C. Parabolic
D. Spiral
3. Alpha-particle that come closer to nuclei:
A. Are deflected more
B. Are deflected less
C. Make more collisions
D. None
4. When alpha particle are sent through a thin metal foil, most of them go straight through the foil because
A. a-particles are much heavier than electrons
B. a- particles are positively charged
C. most part of the atom is empty space
D. a-particle move with high velocity
5. Rutherford‘s ‗alpha (α) particles scattering experiment‘ resulted in to discovery of
A. Electron B. Proton
C. Nucleus in the atom D. Atomic mass
6. Elements with valency 1 are
A. always metals B. always metalloids C. either metals or non-metals D. always non-metals
7. The first model of an atom was given by
A. Bohr
B. E. Goldstein
C. Rutherford
D. J.J. Thomson
8. An atom with 3 protons and 4 neutrons will have a valency of
A. 3
B. 7
C. 1
D. 4
9. What Greek philosopher was the first person to propose the idea that matter is made of tiny individual particles called
atoms?
A. Democritus
B. Dalton
C. Bohr
D. Rutherford
10. The development of the cathode ray tube led to the discovery of what subatomic particle?
A. proton B. electron
C. positron
D. neutron
11. Experimental evidence indicates that die nucleus of an atom
17
A. has a negative charge
C. has no charge
B. contains most of die mass of die atom
D. contains a small percentage of die mass of the atom
12. Which species is diamagnetic?
a) Cr2+
b) Zn
c) Mn
d) C
13. What is the electron confguration for Fe2+?
a) [Ar]4s23d6
b) [Ar]4s23d4
c) [Ar]4s03d6
d) [Ar]4s23d8
14. Name the person who first proposed the atomic theory of matter on scientific basis.
15. Mention the sub atomic particles.
16. What is the basic rule, regarding the behaviour of charged particles?
17. What are cathode rays(cathode ray particles)?
18. Name the quantum number that specifies the shape of an atomic orbital.
19. Name the quantum number that specifies the size of an atomic orbital.
20. Name the quantum number that designates the orientation of the atomic orbital.
21. What is the shape of: (a) s-orbital: (b) p- orbital?
22. On the basis of orientation, how the p-orbitals are designated?
23. What is the maximum number of electrons that can be accommodated in:
(a) sorbital: (b) p- orbitals; (c) d-orbitals;
(d) f-orbitals?
24. How many number of orbital are possible for f-sub shell?
25. What values of the angular momentum (l ) and magnetic (ml) quantum numbers are allowed for a principal quantum
number (n) = 3? How many orbitals exist for n = 3?
26. What is an electron configuration? Provide an example
27. What are degenerate orbitals? According to Hund‘s rule, how are degenerate orbitals occupied?
28. List all orbitals from 1s through 5s according to increasing energy for multi-electron atoms
29. What are valence electrons? Why are they important?
30. Explain why the s block in the periodic table has only two columns while the p block has six.
31. Explain why the rows in the periodic table become progressively longer as we move down the table. For example,
the first row contains 2 elements, the second and third rows each contain 8 elements, and the fourth and fifth rows
each contain 18 elements.
32. List the number of valence electrons for each family in the periodic table, and explain the relationship between
the number of valence electrons and the resulting chemistry of the elements in the family.
a. alkali metals
b. alkaline earth metals
c. halogens
d. oxygen family
33. Magnesium occurs in three fairly common isotopes,
,
and
, which have percent abundances of 78.9%,
10.0 %, and 11.1 %, respectively. Calculate the average atomic mass of magnesium.
34. Chlorine occurs in two common isotopes. It appears as
its average atomic mass?
75.8% of the time and as
18
24.2% of the time. What is