Lesson 1: The Periodic Table
Learning Objectives
Explain how elements are organized into the periodic table.
Describe how some characteristics of elements relate to their positions on the periodic table.
In the 19th century, many previously unknown elements were discovered, and scientists noted that certain sets of elements
had similar chemical properties. For example, chlorine, bromine, and iodine react with other elements (such as sodium) to
make similar compounds. Likewise, lithium, sodium, and potassium react with other elements (such as oxygen) to make
similar compounds. Why is this so?
In 1864, Julius Lothar Meyer, a German chemist, organized the elements by atomic mass and grouped them according to
their chemical properties. Later that decade, Dmitri Mendeleev, a Russian chemist, organized all the known elements
according to similar properties. He left gaps in his table for what he thought were undiscovered elements, and he made
some bold predictions regarding the properties of those undiscovered elements. When elements were later discovered
whose properties closely matched Mendeleev’s predictions, his version of the table gained favor in the scientific community.
Because certain properties of the elements repeat on a regular basis throughout the table (that is, they are periodic), it
became known as the periodic table.
Mendeleev had to list some elements out of the order of their atomic masses to group them with other elements that had
similar properties.
The periodic table is one of the cornerstones of chemistry because it organizes all the known elements on the basis of their
chemical properties. A modern version is shown in Figure 4.1.1
. Most periodic tables provide additional data (such as atomic mass) in a box that contains each element’s symbol. The
elements are listed in order of atomic number.
Figure 4.1.1
: Modern Periodic Table. (Public Domain; PubChem modified Leticia Colmenares).
Features of the Periodic Table
Elements that have similar chemical properties are grouped in columns called groups (or families). As well as being
numbered, some of these groups have names—for example, alkali metals (the first column of elements), alkaline earth
metals (the second column of elements), halogens (the next-to-last column of elements), and noble gases (the last column
of elements).
Each row of elements on the periodic table is called a period. Periods have different lengths; the first period has only 2
elements (hydrogen and helium), while the second and third periods have 8 elements each. The fourth and fifth periods
have 18 elements each, and later periods are so long that a segment from each is removed and placed beneath the main
body of the table.
Metals, Nonmetals and Metalloids
Certain elemental properties become apparent in a survey of the periodic table as a whole. Every element can be classified
as either a metal, a nonmetal, or a semimetal, as shown in Figure 4.1.2. A metal is a substance that is shiny, typically (but
not always) silvery in color, and an excellent conductor of electricity and heat. Metals are also malleable (they can be beaten
into thin sheets) and ductile (they can be drawn into thin wires). A nonmetal is typically dull and a poor conductor of
electricity and heat. Solid nonmetals are also very brittle. As shown in Figure 4.1.2, metals occupy the left three-fourths of
the periodic table, while nonmetals (except for hydrogen) are clustered in the upper right-hand corner of the periodic table.
The elements with properties intermediate between those of metals and nonmetals are called semimetals (or metalloids).
Elements adjacent to the bold zigzag line in the right-hand portion of the periodic table have semimetal properties.
Figure 4.1.2
: Types of Elements. Elements are either metals, nonmetals, or semimetals. Each group is located in a different part of the
periodic table.
Representative, Transition and Inner-transition
Another way to categorize the elements of the periodic table is shown in Figure 4.1.3
. The first two columns on the left and the last six columns on the right are called the main group or representative
elements. The ten-column block between these columns contains the transition metals. The two rows beneath the main
body of the periodic table contain the inner transition metals. The elements in these two rows are also referred to as,
respectively, the lanthanide metals and the actinide metals.
Figure 4.1.3
: Special Names for Sections of the Periodic Table. Some sections of the periodic table have special names. The elements
lithium, sodium, potassium, rubidium, cesium, and francium are collectively known as alkali metals. Alkali metals are the first
column. Alkaline earth metals are the second. Halogens are the second to last column. Noble gases are the last column.
To Your Health: Transition Metals in the Body
Most of the elemental composition of the human body consists of main group elements. The first element appearing on the
list that is not a main group element is iron, at 0.006 percentage by mass. Because iron has relatively massive atoms, it
would appear even lower on a list organized in terms of percent by atoms rather than percent by mass.
Iron is a transition metal. Transition metals have interesting chemical properties, partially because some of their electrons
are in d subshells. The chemistry of iron makes it a key component in the proper functioning of red blood cells.
Red blood cells are cells that transport oxygen from the lungs to cells of the body and then transport carbon dioxide from the
cells to the lungs. Without red blood cells, animal respiration as we know it would not exist. The critical part of the red blood
cell is a protein called hemoglobin. Hemoglobin combines with oxygen and carbon dioxide, transporting these gases from
one location to another in the body. Hemoglobin is a relatively large molecule, with a mass of about 65,000 u.
The crucial atom in the hemoglobin protein is iron. Each hemoglobin molecule has four iron atoms, which act as binding
sites for oxygen. It is the presence of this particular transition metal in your red blood cells that allows you to use the oxygen
you inhale.
Other transition metals have important functions in the body, despite being present in low amounts. Zinc is needed for the
body’s immune system to function properly, as well as for protein synthesis and tissue and cell growth. Copper is also
needed for several proteins to function properly in the body. Manganese is needed for the body to metabolize oxygen
properly. Cobalt is a necessary component of vitamin B-12, a vital nutrient. These last three metals are not listed explicitly in
Table 2.1.2, so they are present in the body in very small quantities. However, even these small quantities are required for
the body to function properly.
As previously noted, the periodic table is arranged so that elements with similar chemical behaviors are in the same group.
Chemists often make general statements about the properties of the elements in a group using descriptive names with
historical origins. For example, the elements of Group 1 are known as the alkali metals, Group 2 are the alkaline earth
metals, Group 17 are the halogens, and Group 18 are the noble gases.
Group 1: The Alkali Metals
The alkali metals are lithium, sodium, potassium, rubidium, cesium, and francium. Hydrogen is unique in that it is generally
placed in Group 1, but it is not a metal. The compounds of the alkali metals are common in nature and daily life. One
example is table salt (sodium chloride); lithium compounds are used in greases, in batteries, and as drugs to treat patients
who exhibit manic-depressive, or bipolar, behavior. Although lithium, rubidium, and cesium are relatively rare in nature, and
francium is so unstable and highly radioactive that it exists in only trace amounts, sodium and potassium are the seventh
and eighth most abundant elements in Earth’s crust, respectively.
Group 2: The Alkaline Earth Metals
The alkaline earth metals are beryllium, magnesium, calcium, strontium, barium, and radium. Beryllium, strontium, and
barium are rare, and radium is unstable and highly radioactive. In contrast, calcium and magnesium are the fifth and sixth
most abundant elements on Earth, respectively; they are found in huge deposits of limestone and other minerals.
Group 17: The Halogens
The halogens are fluorine, chlorine, bromine, iodine, and astatine. The name halogen is derived from the Greek words for
“salt forming,” which reflects that all the halogens react readily with metals to form compounds, such as sodium chloride and
calcium chloride (used in some areas as road salt).
Compounds that contain the fluoride ion are added to toothpaste and the water supply to prevent dental cavities. Fluorine is
also found in Teflon coatings on kitchen utensils. Although chlorofluorocarbon propellants and refrigerants are believed to
lead to the depletion of Earth’s ozone layer and contain both fluorine and chlorine, the latter is responsible for the adverse
effect on the ozone layer. Bromine and iodine are less abundant than chlorine, and astatine is so radioactive that it exists in
only negligible amounts in nature.
Group 18: The Noble Gases
The noble gases are helium, neon, argon, krypton, xenon, and radon. Because the noble gases are composed of only
single atoms, they are called monatomic. At room temperature and pressure, they are unreactive gases. Because of their
lack of reactivity, for many years they were called inert gases or rare gases. However, the first chemical compounds
containing the noble gases were prepared in 1962. Although the noble gases are relatively minor constituents of the
atmosphere, natural gas contains substantial amounts of helium. Because of its low reactivity, argon is often used as an
unreactive (inert) atmosphere for welding and in light bulbs. The red light emitted by neon in a gas discharge tube is used in
neon lights.
To Your Health: Radon
Radon is an invisible, odorless noble gas that is slowly released from the ground, particularly from rocks and soils whose
uranium content is high. Because it is a noble gas, radon is not chemically reactive. Unfortunately, it is radioactive, and
increased exposure to it has been correlated with an increased lung cancer risk.
Because radon comes from the ground, we cannot avoid it entirely. Moreover, because it is denser than air, radon tends to
accumulate in basements, which if improperly ventilated can be hazardous to a building’s inhabitants. Fortunately,
specialized ventilation minimizes the amount of radon that might collect. Special fan-and-vent systems are available that
draw air from below the basement floor, before it can enter the living space, and vent it above the roof of a house.
After smoking, radon is thought to be the second-biggest preventable cause of lung cancer in the United States. The
American Cancer Society estimates that 10% of all lung cancers are related to radon exposure. There is uncertainty
regarding what levels of exposure cause cancer, as well as what the exact causal agent might be (either radon or one of its
breakdown products, many of which are also radioactive and, unlike radon, not gases). The US Environmental Protection
Agency recommends testing every floor below the third floor for radon levels to guard against long-term health effects.
Why do elements in a given group have similar properties?
The periodic table is organized on the basis of similarities in elemental properties, but what explains these similarities? It
turns out that the shape of the periodic table reflects the filling of subshells with electrons, as shown in Figure 4.1.4
. Starting with the first period and going from left to right, the table reproduces the order of filling of the electron subshells in
atoms. Furthermore, elements in the same group share the same valence shell electron configuration. For example, all
elements in the first column have a single s electron in their valence shells, so their electron configurations can be described
as ns1 (where n represents the shell number). This last observation is crucial. Chemistry is largely the result of interactions
between the valence electrons of different atoms. Thus, atoms that have the same valence shell electron configuration will
have similar chemistry.
Figure 4.1.4
The Shape of the Periodic Table. The shape of the periodic table reflects the order in which electron shells and subshells
fill with electrons.
Valence Electrons and Group Number
The number of valence electrons of an element can be determined by the periodic table group (vertical column) in which the
element is categorized. With the exception of groups 3–12 (the transition metals), the units digit of the group number
identifies how many valence electrons are associated with a neutral atom of an element listed under that particular column.
Table 4.1.1
. The Group number and the number of valence electrons.
Periodic table group
Valence electrons
Group 1 (I) (alkali metals)
1
Group 2 (II) (alkaline earth metals)
2
Groups 3-12 (transition metals)
2*
Group 13 (III) (boron group)
3
Group 14 (IV) (carbon group)
4
Group 15 (V) (pnictogens)
5
Group 16 (VI) (chalcogens)
6
Group 17 (VII) (halogens)
7
Group 18 (VIII or 0) (noble gases)
8**
* The general method for counting valence electrons is generally not useful for transition metals.
** Except for helium, which has only two valence electrons.
Atomic Radius
The periodic table is useful for understanding atomic properties that show periodic trends. One such property is the atomic
radius (Figure 4.1.5 ). The atomic radius is defined as one-half the distance between the nuclei of identical atoms that are
bonded together. The units for atomic radii are picometers, equal to 10−12 meters. As an example, the internuclear
distance between the two hydrogen atoms in an H2 molecule is measured to be 74pm . Therefore, the atomic radius of a
hydrogen atom is 742=37pm .
As mentioned earlier, the higher the shell number, the farther from the nucleus the electrons in that shell are likely to be. In
other words, the size of an atom is generally determined by the number of the valence electron shell. Therefore, as we go
down a column on the periodic table, the atomic radius increases. As we go across a period on the periodic table, however,
electrons are being added to the same valence shell; meanwhile, more protons are being added to the nucleus, so the
positive charge of the nucleus is increasing. The increasing positive charge attracts the electrons more strongly, pulling
them closer to the nucleus. Consequently, as we go across a period, from left to right, the atomic radius decreases. These
trends are seen clearly in Figure 4.1.5
Figure 4.1.5
Trends on the Periodic Table. Atomic radii of the representative elements
measured in picometers. The relative sizes of the atoms show several
trends with regard to the structure of the periodic table. Atoms become
larger going down a group and going from right to left across a period.
What is the Reactivity Series?
The reactivity series of metals, also known as the activity series, refers to the arrangement of metals in the descending
order of their reactivities.
The data provided by the reactivity series can be used to predict whether a metal can displace another in a single
displacement reaction. It can also be used to obtain information on the reactivity of metals towards water and acids.
3,093
A chart of the reactivity series of common metals is provided below.
Activity Series of Metals
Metals tend to readily lose electrons and form cations. Most of them react with atmospheric oxygen to form metal oxides.
However, different metals have different reactivities towards oxygen (unreactive metals such as gold and platinum do not
readily form oxides when exposed to air).
Salient Features
The metals at the top of the reactivity series are powerful reducing agents since they are easily oxidized. These metals
tarnish/corrode very easily.
The reducing ability of the metals grows weaker while traversing down the series.
The electro positivity of the elements also reduces while moving down the reactivity series of metals.
All metals that are found above hydrogen in the activity series liberate H2 gas upon reacting with dilute HCl or dilute
H2SO4.
Metals that are placed higher on the reactivity series have the ability to displace metals that are placed lower from their salt
solutions.
Higher ranking metals require greater amounts of energy for their isolation from ores and other compounds.
Another important feature of the activity series is that while travelling down the series, the electron-donating ability of the
metals reduces.
Long Tabular Form of the Reactivity Series
The reactivities of metals are tabulated below (in the descending order) along with their corresponding ions. Note that the
metals in Red react with cold water, those in Orange cannot react with cold water but can react with acids, and those in Blue
only react with some strong oxidizing acids.
Reactivity Series of Metals
Caesium
Francium
Rubidium
Potassium
Sodium
Lithium
Barium
Radium
Strontium
Calcium
Magnesium
Beryllium
Aluminium
Titanium
Manganese
Zinc
Ions Formed
Cs+
Fr+
Rb+
K+
Na+
Li+
Ba2+
Ra2+
Sr2+
Ca2+
Mg2+
Be2+
Al3+
Ti4+
Mn2+
Zn2+
Reactivity Series of Metals
Ions Formed
Chromium
Iron
Cadmium
Cobalt
Nickel
Tin
Lead
Hydrogen
Antimony
Bismuth
Copper
Tungsten
Mercury
Silver
Platinum
Gold
Cr3+
Fe3+
Cd2+
Co2+
Ni2+
Sn2+
Pb2+
H+ (Non-Metal, Reference for Comparison)
Sb3+
Bi3+
Cu2+
W3+
Hg2+
Ag+
Pt4+
Au3+
Despite being a non-metal, hydrogen is often included in the reactivity series since it helps compare the reactivities of the
metals. The metals placed above hydrogen in the series can displace it from acids such as HCl and H2SO4 (since they are
more reactive).
Important uses of Reactivity Series
Apart from providing insight into the properties and reactivities of the metals, the reactivity series has several other
important applications. For example, the outcome of the reactions between metals and water, metals and acids, and single
displacement reactions between metals can be predicted with the help of the activity series.
Reaction Between Metals and Water
Calcium and the metals that are more reactive than calcium in the reactivity series can react with cold water to form the
corresponding hydroxide while liberating hydrogen gas. For example, the reaction between potassium and water yields
potassium hydroxide and H2 gas, as described by the chemical equation provided below.
2K + 2H2O → 2KOH + H2
Therefore, the reactivity series of metals can be used to predict the reactions between metals and water.
Reaction Between Metals and Acids
Lead and the metals ranking above lead on the activity series form salts when reacted with hydrochloric acid or sulphuric
acid. These reactions also involve the liberation of hydrogen gas. The reaction between zinc and sulphuric acid is an
example of such a reaction. Here, zinc sulfate and H2 gas are formed as products. The chemical equation is:
Zn + H2SO4 → ZnSO4 + H2
Thus, the reactions between metals and some acids can be predicted with the help of the reactivity series.
Single Displacement Reactions Between Metals
The ions of low ranking metals are readily reduced by high ranking metals on the reactivity series. Therefore, low ranking
metals are easily displaced by high ranking metals in the single displacement reactions between them.
A great example of such a reaction is the displacement of copper from copper sulphate by zinc. The chemical equation for
this reaction is given by:
Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)
This concept has several practical applications in the extraction of metals. For example, titanium is extracted from titanium
tetrachloride via a single displacement reaction with magnesium. Thus, the reactivity series of metals can also be used to
predict the outcome of single displacement reactions.
How to find the Atomic Number
The atomic number of an element is simply the number of protons in its nucleus. The easiest way to find the atomic number,
is to look on a periodic table, the atomic number is in the upper left corner, or is the largest number on the square.
Finding the Number of Protons
The number of protons in an atom is equal to the atomic number of the element. For example, let’s use oxygen. According
to the periodic table, oxygen has the atomic number eight. The atomic number is located above the element’s symbol. Since
oxygen has an atomic number of eight, there must be eight protons total. Moreover, the number of protons never changes
for an element.
Finding the Number of Neutrons
The number of neutrons in an atom can be calculated by subtracting the atomic number from the atomic mass. Both of
these numbers can be found on the periodic table. The atomic number is listed above the symbol of the element whereas
the mass number is placed below. Let’s keep using oxygen as our example. Its atomic mass is 15.999 atomic mass units
(amu) and its atomic number is 8. When we subtract 8 from 15.999, we will get 8. Also, it should be noted that the number of
neutrons for an element may vary. Some elements have isotopes, which have different masses and therefore different
numbers of neutrons.
Finding the Number of Electrons
The number of electrons in an atom is equal to the atomic number of an element, for neutrally charged species. This means
the number of electrons and the number of protons in an element are equal. Therefore, the number of electrons in oxygen is
8. Moreover, since these two subatomic particles, electrons and protons, have opposite charges, they cancel out and keep
the atom neutral.
Protons, Neutrons, and Electrons Summary Table
{Number of Protons} = {Atomic Number}
{Number of Neutrons} = {Atomic Mass-{Atomic Number}
{Number of Electrons} = {Atomic Number}
Figure 1. Atoms are the building blocks of
molecules found in the universe—air, soil,
water, rocks—and also the cells of all
living organisms. In this model of an
organic molecule, the atoms of carbon
(black), hydrogen (white), nitrogen (blue),
oxygen (red), and sulfur (yellow) are
shown in proportional atomic size. The
silver rods represent chemical bonds.
(credit: modification of work by Christian
Guthier)
Lesson 2: The Structure of an atom
An atom is the smallest unit of matter that retains all of the chemical properties of an element. Elements are forms of matter
with specific chemical and physical properties that cannot be broken down into smaller substances by ordinary chemical
reactions.
The chemistry discussed in BIS2A requires us to use a model for an atom. While there are more sophisticated models, the
atomic model used in this course makes the simplifying assumption that the standard atom is composed of three subatomic
particles, the proton, the neutron, and the electron. Protons and neutrons have a mass of approximately one atomic mass
unit (a.m.u.). One atomic mass unit is approximately 1.660538921 x 10-27kg—roughly 1/12 of the mass of a carbon atom
(see table below for more precise value). The mass of an electron is 0.000548597 a.m.u. or 9.1 x 10-31kg. Neutrons and
protons reside at the center of the atom in a region call the nucleus while the electrons orbit around the nucleus in zones
called orbitals, as illustrated below. The only exception to this description is the hydrogen (H) atom, which is composed of
one proton and one electron with no neutrons. An atom is assigned an atomic number based on the number of protons in
the nucleus. Neutral carbon (C), for instance has six neutrons, six protons, and six electrons. It has an atomic number of six
and a mass of slightly more than 12 a.m.u.
Table 1. Charge, mass, and location of subatomic particles
Protons, neutrons, and electrons
Charge
Mass (a.m.u.)
Proton
+1
~1
Neutron
0
~1
Electron
–1
~0
Mass (kg)
1.6726 x 10-27
1.6749 x 10-27
9.1094 x 10-31
Location
nucleus
nucleus
orbitals
Table 1 reports the charge and location of three subatomic particles—the neutron, proton, and electron. Atomic mass unit =
a.m.u. (a.k.a. dalton [Da])—this is defined as approximately one twelfth of the mass of a neutral carbon atom or
1.660538921 x 10−27 kg. This is roughly the mass of a proton or neutron.
Figure 2. Elements, such as helium depicted here, are made up of atoms.
Atoms are made up of protons and neutrons located within the nucleus and
electrons surrounding the nucleus in regions called orbitals. (Note: This
figure depicts a Bohr model for an atom—we could use a new open source
figure that depicts a more modern model for orbitals. If anyone finds one
please forward it.)
Relative sizes and distribution of elements
The typical atom has a radius of one to two angstroms (Å). 1Å = 1 x 10-10m. The typical nucleus has a radius of 1 x 10-5Å
or 10,000 smaller than the radius of the whole atom. By analogy, a typical large exercise ball has a radius of 0.85m. If this
were an atom, the nucleus would have a radius about 1/2 to 1/10 of your thinnest hair. All of that extra volume is occupied
by the electrons in regions called orbitals. For an ideal atom, orbitals are probabilistically defined regions in space around
the nucleus in which an electron can be expected to be found.
The properties of living and nonliving materials are determined to a large degree by the composition and organization of
their constituent elements. Five elements are common to all living organisms: Oxygen (O), Carbon (C), Hydrogen (H),
Phosphorous (P), and Nitrogen (N). Other elements like Sulfur (S), Calcium (Ca), Chloride (Cl), Sodium (Na), Iron (Fe),
Cobalt (Co), Magnesium, Potassium (K), and several other trace elements are also necessary for life, but are typically found
in far less abundance than the "top five" noted above. As a consequence, life's chemistry—and by extension the chemistry
of relevance in BIS2A—largely focuses on common arrangements of and reactions between the "top five" core atoms of
biology.
Figure 3. A table illustrating the abundance of elements in the
human body. A pie chart illustrating the relationships in
abundance between the four most common elements.
J.J. Thomson Biographical Data
Tomson was born December 18, 1856, Cheetham Hill, near Manchester, England. He died August 30, 1940, Cambridge,
Cambridgeshire, England. Thomson is buried in Westminster Abbey, near Sir Isaac Newton. J.J. Thomson is credited with
the discovery of the electron, the negatively charged particle in the atom. He is known for the Thomson atomic theory.
Many scientists studied the electric discharge of a cathode ray tube. It was Thomson's interpretation that was important. He
took the deflection of the rays by the magnets and charged plates as evidence of "bodies much smaller than atoms."
Thomson calculated these bodies had a large charge-to-mass ratio and he estimated the value of the charge itself. In 1904,
Thomson proposed a model of the atom as a sphere of positive matter with electrons positioned based on electrostatic
forces. So, he not only discovered the electron but determined it was a fundamental part of an atom.
Thomson Atomic Theory
Thomson's discovery of the electron completely changed the way people viewed atoms. Up until the end of the 19th
century, atoms were thought to be tiny solid spheres. In 1903, Thomson proposed a model of the atom consisting of positive
and negative charges, present in equal amounts so that an atom would be electrically neutral. He proposed the atom was a
sphere, but the positive and negative charges were embedded within it. Thomson's model came to be called the "plum
pudding model" or "chocolate chip cookie model". Modern scientists understand atoms consist of a nucleus of positivelycharged protons and neutral neutrons, with negatively-charged electrons orbiting the nucleus. Yet, Thomson's model is
important because it introduced the notion that an atom consisted of charged particles.
Interesting Facts About J.J. Thomson
Prior to Thomson's discovery of electrons, scientists believed the atom was the smallest fundamental unit of matter.
Thomson called the particle he discovered 'corpuscles' rather than electrons.
Thomson's master's work, Treatise on the motion of vortex rings, provides a mathematical description of William Thomson's
vortex theory of atoms. He was awarded the Adams Prize in 1884.
Thomson discovered the natural radioactivity of potassium in 1905.
In 1906, Thomson demonstrated a hydrogen atom had only a single electron.
Thomson's father intended for J.J. to be an engineer, but the family did not have the funds to support the apprenticeship.
So, Joseph John attended Owens College in Manchester, and then Trinity College in Cambridge, where he became a
mathematical physicist.
In 1890, Thomson married one of his students, Rose Elisabeth Paget. They had a son and a daughter. The son, Sir George
Paget Thomson, received the Nobel Prize in Physics in 1937.
Thomson also investigated the nature of positively-charged particles. These experiments led to the development of the
mass spectrograph.
Thomson was closely aligned with chemists of the time. His atomic theory helped explain atomic bonding and the structure
of molecules. Thomson published an important monograph in 1913 urging the use of the mass spectrograph in chemical
analysis.
Many consider J.J. Thomson's greatest contribution to science to be his role as a teacher. Seven of his research assistants,
as well as his own son, went on to win the Nobel Prize in Physics. One of his best-known students was Ernest Rutherford,
who succeeded Thomson as Cavendish Professor of Physics.
3.4: Rutherford's Experiment- The Nuclear Model of the Atom
Learning Objectives
Describe Thomson's "plum pudding" model of the atom and the evidence for it.
Describe Rutherford's gold foil experiment and explain how this experiment altered the "plum pudding" model.
The electron was discovered by J.J. Thomson in 1897. The existence of protons was also known, as was the fact that
atoms were neutral in charge. Since the intact atom had no net charge and the electron and proton had opposite charges,
the next step after the discovery of subatomic particles was to figure out how these particles were arranged in the atom.
This is a difficult task because of the incredibly small size of the atom. Therefore, scientists set out to design a model of
what they believed the atom could look like. The goal of each atomic model was to accurately represent all of the
experimental evidence about atoms in the simplest way possible.
Following the discovery of the electron, J.J. Thomson developed what became known as the "plum pudding" model (Figure
\(\PageIndex{1}\)) in 1904. Plum pudding is an English dessert similar to a blueberry muffin. In Thomson's plum pudding
model of the atom, the electrons were embedded in a uniform sphere of positive charge like blueberries stuck into a muffin.
The positive matter was thought to be jelly-like or a thick soup. The electrons were somewhat mobile. As they got closer to
the outer portion of the atom, the positive charge in the region was greater than the neighboring negative charges and the
electron would be pulled back more toward the center region of the atom.
Figure \(\PageIndex{1}\) The "plum pudding" model.
However, this model of the atom soon gave way to a new model developed by New
Zealander Ernest Rutherford (1871 - 1937) about five years later. Thomson did still
receive many honors during his lifetime, including being awarded the Nobel Prize in
Physics in 1906 and a knighthood in 1908.
Atoms and Gold
In 1911, Rutherford and coworkers Hans Geiger and Ernest Marsden initiated a series of
groundbreaking experiments that would completely change the accepted model of the
atom. They bombarded very thin sheets of gold foil with fast moving alpha particles.
According to the accepted atomic model, in which an atom's mass and charge are uniformly distributed throughout the
atom, the scientists expected that all of the alpha particles would pass through the gold foil with only a slight deflection or
none at all. Surprisingly, as shown in Figure \(\PageIndex{2}\) (while most of the alpha particles were indeed undeflected, a
very small percentage (about 1 in 8000 particles) bounced off the gold foil at very large angles. Some were even redirected
back toward the source. No prior knowledge had prepared them for this discovery. In a famous quote, Rutherford exclaimed
that it was "as if you had fired a 15-inch [artillery] shell at a piece of tissue and it came back and hit you."
Rutherford needed to come up with an entirely new model of the atom in order to explain his results. Because the vast
majority of the alpha particles had passed through the gold, he reasoned that most of the atom was empty space. In
contrast, the particles that were highly deflected must have experienced a tremendously powerful force within the atom. He
concluded that all of the positive charge and the majority of the mass of the atom must be concentrated in a very small
space in the atom's interior, which he called the nucleus. The nucleus is the tiny, dense, central core of the atom and is
composed of protons and neutrons.
Rutherford's atomic model became known as the nuclear model. In the nuclear atom, the protons and neutrons, which
comprise nearly all of the mass of the atom, are located in the nucleus at the center of the atom. The electrons are
distributed around the nucleus and occupy most of the volume of the atom. It is worth emphasizing just how small the
nucleus is compared to the rest of the atom. If we could blow up an atom to be the size of a large professional football
stadium, the nucleus would be about the size of a marble.
Rutherford's model proved to be an important step towards a full understanding of the atom. However, it did not completely
address the nature of the electrons and the way in which they occupied the vast space around the nucleus. For this and
other insights, Rutherford was awarded the Nobel Prize in Chemistry in 1908. Unfortunately, Rutherford would have
preferred to receive the Nobel Prize in Physics because he considered physics superior to chemistry. In his opinion, “All
science is either physics or stamp collecting.”
Henry Moseley
Henry Moseley (born November 23, 1887, Weymouth, Dorset, England—died August 10, 1915, Gallipoli, Turkey) English
physicist who experimentally demonstrated that the major properties of an element are determined by the atomic number,
not by the atomic weight, and firmly established the relationship between atomic number and the charge of the atomic
nucleus.
Educated at Trinity College, Oxford, Moseley in 1910 was appointed lecturer in physics at Ernest (later Lord) Rutherford’s
laboratory at the University of Manchester, where he worked until the outbreak of World War I, when he entered the army.
His first researches were concerned with radioactivity and beta radiation in radium. He then turned to the study of the X-ray
spectra of the elements. In a brilliant series of experiments he found a relationship between the frequencies of
corresponding lines in the X-ray spectra. In a paper published in 1913, he reported that the frequencies are proportional to
the squares of whole numbers that are equal to the atomic number plus a constant.
Periodic Table of the elements concept image (chemistry)
Britannica Quiz
Facts You Should Know: The Periodic Table Quiz
Known as Moseley’s law, this fundamental discovery concerning atomic numbers was a milestone in advancing the
knowledge of the atom. In 1914 Moseley published a paper in which he concluded that the atomic number is the number of
positive charges in the atomic nucleus. He also stated that there were three unknown elements, with atomic numbers 43,
61, and 75, between aluminum and gold. (There are, in fact, four. Moseley identified gaps in the periodic table for
technetium [43], promethium [61], and rhenium [75], but he missed hafnium [atomic number 72] because its discovery had
been erroneously claimed.)
Moseley enlisted in the army when World War I broke out in 1914. He was shot in the head by a Turkish sniper at the Battle
of Suvla Bay (in Turkey). His death at the age of 27 deprived the world of one of its most promising experimental physicists.
Electron Shells and the Bohr Model
Orbitals in the Bohr model: The Bohr model was developed by Niels Bohr in 1913. In this model, electrons exist within
principal shells. An electron normally exists in the lowest energy shell available, which is the one closest to the nucleus.
Energy from a photon of light can bump it up to a higher energy shell, but this situation is unstable and the electron quickly
decays back to the ground state. In the process, a photon of light is released.
As previously discussed, there is a connection between the number of protons in an element, the atomic number that
distinguishes one element from another, and the number of electrons it has. In all electrically-neutral atoms, the number of
electrons is the same as the number of protons. Each element, when electrically neutral, has a number of electrons equal to
its atomic number.
An early model of the atom was developed in 1913 by Danish scientist Niels Bohr (1885–1962). The Bohr model shows the
atom as a central nucleus containing protons and neutrons with the electrons in circular orbitals at specific distances from
the nucleus. These orbits form electron shells or energy levels, which are a way of visualizing the number of electrons in the
various shells. These energy levels are designated by a number and the symbol “n.” For example, 1n represents the first
energy level located closest to the nucleus.
Electrons fill orbit shells in a consistent order. Under standard conditions, atoms fill the inner shells (closer to the nucleus)
first, often resulting in a variable number of electrons in the outermost shell. The innermost shell has a maximum of two
electrons, but the next two electron shells can each have a maximum of eight electrons. This is known as the octet rule
which states that, with the exception of the innermost shell, atoms are more stable energetically when they have eight
electrons in their valence shell, the outermost electron shell. Examples of some neutral atoms and their electron
configurations are shown in. As shown, helium has a complete outer electron shell, with two electrons filling its first and only
shell. Similarly, neon has a complete outer 2n shell containing eight electrons. In contrast, chlorine and sodium have seven
and one electrons in their outer shells, respectively. Theoretically, they would be more energetically stable if they followed
the octet rule and had eight.
Figure 2.5.1
: Bohr diagrams: Bohr diagrams indicate how many
electrons fill each principal shell. Group 18 elements
(helium, neon, and argon are shown) have a full outer,
or valence, shell. A full valence shell is the most stable
electron configuration. Elements in other groups have
partially-filled valence shells and gain or lose electrons
to achieve a stable electron configuration.
An atom may gain or lose electrons to achieve a full
valence shell, the most stable electron configuration.
The periodic table is arranged in columns and rows
based on the number of electrons and where these
electrons are located, providing a tool to understand
how electrons are distributed in the outer shell of an
atom. As shown in, the group 18 atoms helium (He), neon (Ne), and argon (Ar) all have filled outer electron shells, making it
unnecessary for them to gain or lose electrons to attain stability; they are highly stable as single atoms. Their non-reactivity
has resulted in their being named the inert gases (or noble gases). In comparison, the group 1 elements, including hydrogen
(H), lithium (Li), and sodium (Na), all have one electron in their outermost shells. This means that they can achieve a stable
configuration and a filled outer shell by donating or losing an electron. As a result of losing a negatively-charged electron,
they become positively-charged ions. When an atom loses an electron to become a positively-charged ion, this is indicated
by a plus sign after the element symbol; for example, Na+. Group 17 elements, including fluorine and chlorine, have seven
electrons in their outermost shells; they tend to fill this shell by gaining an electron from other atoms, making them
negatively-charged ions. When an atom gains an electron to become a negatively-charged ion this is indicated by a minus
sign after the element symbol; for example, F-. Thus, the columns of the periodic table represent the potential shared state
of these elements’ outer electron shells that is responsible for their similar chemical characteristics.
Key Points
In the Bohr model of the atom, the nucleus contains the majority of the mass of the atom in its protons and neutrons.
Orbiting the positively-charged core are the negatively charged electrons, which contribute little in terms of mass, but are
electrically equivalent to the protons in the nucleus.
In most cases, electrons fill the lower- energy orbitals first, followed by the next higher energy orbital until it is full, and so on
until all electrons have been placed.
Atoms tend to be most stable with a full outer shell (one which, after the first, contains 8 electrons), leading to what is
commonly called the ” octet rule “.
The properties of an element are determined by its outermost electrons, or those in the highest energy orbital.
Atoms that do not have full outer shells will tend to gain or lose electrons, resulting in a full outer shell and, therefore,
stability.