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The Table of Elements

An element is a pure substance that cannot be broken down into a simpler chemical substance
by any physical or chemical means.

For Example) Silver cannot be broken down into a simpler chemical substance.

Elements are the building blocks of all substances.

An element

For Example) Silver is represented by the symbol Ag (which stands for Argentum – the Latin
symbol is an abbreviation for a chemical element.
word for Silver).

A compound is a pure substance composed of two or more different elements that are
chemically joined together.

For Example) Water (H2O); Methane (CH4)

Metals:

A metal is an element that is lustrous, malleable, ductile, and conducts heat and
electricity.

Metals are always solids.

In general, metals are located on the left and central parts of the periodic table.

Metals become more reactive (easily able to react during a chemical reaction) as you move
from right to left, and from top to bottom. Caesium (Cs) is the most reactive metal (even
though it is above Francium).

There is a reason why pots and pans are made of metal – metal is easily shaped and an
excellent conductor of thermal energy (heat).

Copper is very flexible and ductile – which allows it to be pulled into wires.

The lustre and malleability of gold and silver (and their resistance to corrosion) makes
them ideal for jewelry.

Non-metals:

A non-metal is an element, usually a gas or a dull-powdery solid, that does not conduct
heat or electricity.

Non-metals are generally found in the upper right portion of the periodic table.

The only liquid non-metal is bromine.

Non-metals are usually gases at room temperature (for example: nitrogen, oxygen, and
hydrogen)

Non-metals are that are solids are not shiny, not ductile, and not malleable.

For example: Coal is a solid non-metal (mainly composed of carbon) – it is dull and brittle.

Since non-metals are generally poor conductors of thermal energy and electricity, we can
take advantage of these properties when we insulate our houses in the winter. For
example, we use argon gas in between two panes of glass to reduce energy loss through our
windows.

Metalloids:

A metalloid is an element that has properties of both metals and non-metals.

A bold line that resembles a downward staircase, starting from boron, separates the
metals and the non-metals on the periodic table. Elements located along the staircase line
are called metalloids.

For example) boron, silicon, and arsenic
Patterns on the Periodic Table

The periodic table is arranged in a particular way.

All the elements in the same column have similar physical and chemical properties. For example:
since lithium, sodium and potassium are all in the same column, they would all have similar physical
and chemical properties.

Chemical Families:

A Chemical

The chemical families are numbered from left to right on the periodic table.

The far left column on the periodic table is Group #1 (and is called Alkali Metals)

Family is a column of elements with similar properties on the periodic table.
Alkali Metals (lithium, sodium, and potassium) are all shiny, silvery, and soft,
but highly reactive. They all have relatively low densities and they can float on
water. We see them in many everyday things, such as: table salt (NaCl), baking
soda (NaHCO3); potassium is an important mineral found in bananas.

The 2nd column to the left on the periodic table is Group #2 (and is called Alkaline Earth
Metals)

Alkaline Earth Metals (beryllium, magnesium, calcium, strontium, barium, and
radium) are shiny and silvery, but are not as soft or reactive as Alkali Metals. We
see them in everyday things such as getting calcium in milk and dairy products to
help build strong bones and teeth; strontium is used to build a strong shell in coral;
magnesium is used in fireworks since it burns with bright, colourful flames.

The 2nd column to the far right of the periodic table is Group #17 (and is called Halogens)

All Halogens are very reactive therefore, they are rarely found in elemental
form. They often form compounds with alkali metals. Many halogens are poisonous
in large amounts. However, chlorine is poisonous in small amounts (that is why they
used chlorine gas as a chemical weapon during World War 1). We use chlorine in
small, diluted amounts to kill bacteria in swimming pools. Iodine is usually dissolved
in alcohol and is used to disinfect scrapes and cuts. Bromine can be added to
lightbulbs to make them burn brighter.

The halogens fluorine, and chlorine are gases at room temperature; bromine is a
liquid; and iodine and astatine are solids.

The far right column on the periodic table is Group #18 (and is called Noble Gases)

Noble Gases are elements that are very stable in nature and are unreactive.
Noble gases will not react with other elements and therefore, you will never see
them used in a chemical reaction. Noble gases are colourless, odourless, and
tasteless. However, they glow brightly when an electrical current is passed
through them. For example) Argon glows blue, Krypton glows pink, and Xenon
glows purple. Also, because they are so stable, they are safe to use in various
situations: Helium is used in balloons.

Periods:

A period is a row on the periodic table.

Elements in the same horizontal row on the periodic table show some trends of increasing
or decreasing reactivity.
The Evolution of the Atomic Theory
A)
An Indivisible Particle – The Atom

Around 400 BCE, the Greek philosopher Democritus proposed that all matter can be divided into
smaller and smaller pieces until a single indivisible particle is reached. He named this particle the
atom (the smallest unit of an element).

Democritus proposed that atoms are: of different sizes, in constant motion, and are separated
by empty spaces (voids).
B)
Earth, Water, Air, and Fire: Aristotle (around 450 BCE)

Another famous Greek philosopher, Aristotle, rejected the idea of the atom.

Aristotle believed that all matter is made up of four basic substances: earth, water, air, and fire
– these substances were thought to have four specific properties: dry, wet, cold, and hot.
C)
The Billiard Ball Model

In 1807, John Dalton revived Democritus’s theory of the indivisible atom.

Dalton proposed that: all atoms of an element are identical; atoms of different elements are
different; and atoms are rearranged to form new substances in chemical reactions, but that they
are never created or destroyed.
D)
Thomson’s Experiment – The Electron

In 1897, J.J. Thomson discovered that extremely small negatively charged particles could be
emitted by very hot materials. Those particles were attracted to the positive ends of a circuit.

Thomson theorized that: atoms contain negatively charged electrons; since atoms are neutral,
the rest of the atom is a positively charged sphere; and, negatively charged electrons are evenly
distributes thoughout the atoms.

Thomson’s model was called the “plum pudding” model because the electrons embedded in the
atom resembled the raisins in a plum pudding.
E)
The Gold Foil Experiment – The Nucleus and the Proton

In 1909, Ernest Rutherford predicted that if positive and negative charges were uniformly
distributed throughout atoms, then tiny positively charged particles shot at a thin piece of gold
foil would pass through the foil.

When the experiment was performed, most of the particles passed through the foil unaffected.
Also, a small number of particles were deflected at very large angles, as though something was
repelling them.

Rutherford reasoned that these large angles of deflection were caused by a collision with a small,
concentrated, positively charged central mass inside the atom.

In Rutherford’s revised model, he theorized that: The centre of the atom has a positive charge.
This centre is called the nucleus – it contains the atom’s mass but occupies a very small space; the
nucleus is surrounded by a cloud of negatively charged electrons; and, most of the atom is empty
space.

In 1920, Rutherford discovered the proton (a positively charged particle in the atom’s nucleus).
Measurements of the atomic mass showed that protons alone could not account for the total mass
of a nucleus. So, Rutherford predicted that there must be a third particle in the nucleus that
has about the same mass as the proton, but that was neutral in charge.
F)
Chadwick’s Experiments – The Neutron

In 1932, James Chadwick (Rutherford’s student) proposed that: An atom must be an empty
sphere with a tine dense central nucleus; this nucleus contains positively charged protons and
neutral particles called neutrons; the mass of a neutron is about the same as that of a proton;
negatively charged electrons circle rapidly through the empty space that surrounds the nucleus;
and a neutral atom has the same number of protons as electrons.

This model of the atom is called the planetary model.
Charge
Location
Relative Mass
Symbol
G)
Proton
Neutron
Electron
Positive (+)
In nucleus
1
P+
Neutral (O)
In nucleus
1
nO
Negative (-)
Orbitting nucleus
1/2000
e-
Electron Orbits

Niels Bohr, a Danish scientist, studied the hydrogen atom and the light that it produces when it is
excited by thermal energy or electricity.

When white light is shone through a prism, a full rainbow of colours is seen. When light produced
by hydrogen is examined in the same way, only a few lines of colour are seen. Most colours are
missing.

In 1913, Bohr used this evidence to propose the following theory: The farther apart the electron
is from the nucleus, the greater its energy; electrons cannot be between orbits, but they can
jump to and from different orbits; electrons release energy as light when they jump from higher
to lower orbits; each orbit can hold a certain maximum number of electrons – the maximum
number of electrons in the first, second, and third shells are 2, 8, and 8 respectively.

This model of the atom is known at the Bohr-Rutherford model, and is useful at explaining the
properties of the first 20 elements.
Explaining the Periodic Table

Atomic Number:

The Atomic
Number is the number of protons in the atom’s nucleus.

For example) Hydrogen has 1 proton in its nucleus, so it has an atomic number of 1.

Every element’s atomic number is listed on the periodic table. Elements are arranged
according to increasing atomic number on the periodic table.

Mass Number:

The Mass
Number is the number of protons and neutrons in an atom’s nucleus.

For example) Lithium has an atomic number of 3 – so all lithium atoms contain 3 protons.
However, lithium’s mass number if 7 – so there must be 4 neutrons along with the 3
protons within lithium’s nucleus.

Sometimes, an element can have variations of how many neutrons it has within its nucleus.
When an atom has the same number of protons, but a different number of neutrons within
its nucleus, it is called an
isotope.
For example) The lithium isotope with a mass number
of 6 is called lithium-6. The lithium isotope with a mass number of 7 is called lithium-7.

Atomic Mass:

The Atomic
Mass is the mass of an atom in atomic mass units (u).

The atomic mass of each element is given below the element symbol on the periodic table.

The atomic masses given on the periodic table are not whole numbers. For example)
Lithium has an atomic mass of 6.94 u. This is determined by taking a weighted average of
the masses of its isotopes. You can usually determine the most common isotope for an
element by rounding the atomic mass to the nearest whole number.

Bohr-Rutherford Diagrams of an Atom:

A Bohr-Rutherford Diagram shows the numbers and locations of protons, neutrons, and
electrons in an atom.

The number of protons equals the atomic number

The number of neutrons equals the difference between the mass number and the atomic
number

The number of electrons equals the number of protons in a neutral atom.

For example) Sodium has an atomic number of 11 and its mass number is 22.990. We can
round this mass number to 23. Therefore, sodium will have 11 protons (atomic number), 12
neutrons (mass number minus atomic number), and 11 electrons (same as the number of
protons).

The element’s group number will represent the number of electrons in its last (valence)
shell. For example) All elements in Group #1 will have 1 electron in the last shell; and all
elements in group 5 will have 5 electrons in their last shell.

The element’s period number will represent how many shells (orbits) it will have around its
nucleus. For example) All elements in period #2 will have 2 shells surrounding its nucleus;
and all elements in period 4 will have 4 shells surroundings its nucleus.

The electrons that are furthest from the nucleus have the weakest attraction to the
nucleus, so it is easily able to jump to another energy shell.
 Practice Problems for Bohr-Rutherford Diagrams:
a) Sodium
b) Nitrogen
c) Helium
d) Calcium
From Charcoal to Diamonds

Charcoal, pencil, and diamonds are all made up of carbon atoms. They only differ in how the
carbon atoms are arranged.

Charcoal is created when carbon atoms form an endless structure that is shapeless and
disorganized. It appears as a soft black solid, and we use it in our barbeque.

Graphite is an organized structure in comparison to charcoal.
Each carbon atom attaches to
three other carbon atoms to form a sheet of interconnected hexagons. This form of carbon is
what pencil “lead” is actually made from – under slight pressure, the carbon sheets slide across
each other, leaving behind the top layer of carbon atoms on the surface of the writing paper.

Diamonds are created when carbon atoms are arranged under conditions of extremely high
temperatures and pressure. The carbon atoms are interconnected in three dimensions – similar to
a playground climbing structure. This strongly reinforced framework is what gives diamond its
remarkable hardness. The closeness of the atoms makes diamonds very dense and also allows
light to bend, producing its much admired sparkle when diamond is cut.