The Periodic Table
The most fun you can have with a bunch of squares.
 Take a Moment to make a list of two or
three elements you like, or are familiar
with.
 What do you know about that element?
Why do you remember it?
Why do we organize the elements?
 A quick history:
Many elements were familiar since
ancient times. Au, Ag, Hg, Sn, C, Pb, Cu, Fe
and S, had been discovered by various
cultures.
The practice of alchemy led to the
discovery of Phosphorus, by Hennig Brand
around 1669. He hides his discovery.
What is an element?
 Early definition Earth, Air, Fire, Water, Ether….
 Later definitions included the discovery that one
element could not be transformed into another by
ordinary means.
 The discovery of radioactive decay alters this idea, but
mainly, this is true.
 Our definition? An element will always have the same
atomic number!!!
Why do we organize the elements?
 Between 1669 and 1869, the work of discovering and
documenting new elements becomes one of the
focuses of practitioners of science everywhere.
 Practitioners of science everywhere are beginning to
look for patterns.
 Why?
1) They wanted to discover more.
2) Understanding relationships between elements
would further their understanding of reactions, and
compounds.
Attempts at Organization
 Johann Dobereiner – begins to work with element
properties and starts associating them in groups of
three, called triads.
 A.E. Beguyer de Chancourtois – makes a cylindrical
table, first geometric representation of patterns found
in elemental behavior.
 John Newlands – proposes a musical analogy and
organizes them into octets, groups of eight. Also shows
the patterns, but is too limiting.
A.E. Beguyer de Chancourtois
Dmitri Mendeleev
 Publishes a table in 1869 that gives rise to the modern
table layout
 He organizes the elements by mass, which is the
information they have at the time
 Some of the masses are incorrectly calculated so the
table is not entirely in order
 His table is considered the most influential because it
leaves room for expansion and anticipates the
existence of elements of other masses, leaving spaces
for them to be placed there in the future.
Atomic Number
 Later, work by Henry Moseley enables chemists and
physicists to accurately find the number of protons in
the nucleus of an atom.
 Using this information the table is placed in the order
we use today.
 THE ELEMENTS ARE IN ORDER BY THEIR ATOMIC
NUMBERS.
 This allows patterns to emerge.
Elements are in groups and periods
 Groups are also known as families, these are elements
in columns, that share similar chemical behavior.
 Periods are rows across the table.
 While a regents chemistry table will show electron
configurations, your “naked” tables will not, so you
should know 2 things:
1) All elements in a group have the same # of valence
electrons
2) All elements in a period have the same number of
energy levels (shells).
Important Groups
 Definition: Alkali – derived from the Arabic word, al-
qali, which referred to ashes, which when mixed with
water form a basic solution used throughout history to
make soap.
 Groups known as alkali will make salts, (you will
recognize later as ionic compounds), that form basic
solutions when placed in water.
 Group 1 = Alkali metals
Alkaline Metals
 Group 1: Alkali metals:
- Form ions with 1+ charge by losing one electron.
- Do not occur in their metallic form in nature,
they are too reactive. They are always bound in a
compound.
- These are the most reactive metals on the table!!
Lithium Reacting with
Water
Earth Alkaline Metals
 Group 2: Known as the Earth Alkaline Metals.
- Will acquire 2+ charges due to the loss of two
valence electrons
- Also very reactive, these will also not be found
in nature in their metallic forms.
- The elements Ra – Radium is a naturally
occurring radioactive element, may not behave as
others in the group.
Strontium and Barium are more reactive
and are stored in oil.
Magnesium burns and is
used for survival and
camping type fire starting.
Transition Metals
 Known as transition metals because they
make the transition to using the d orbitals
for bonding.
 Transition metals behave differently than
group 1 and two metals, and are often able to
acquire a variety of charges, depending on
which element they are reacting with.
Transition metals form colored
compounds and solutions!
Other Metals
Metalloids
The Halogens
 Group 17 are elements known as the Halogens
 When they form compounds, they are often
referred to as halide compounds
 Halogens tend to gain one electron from other
elements while bonding, so they acquire 1- charges
THESE ARE THE MOST REACTIVE
NON-METALS
The Noble Gases
 Group 18 on the table.
 Often called INERT gases, under normal
circumstances they will not react with anything.
 Have a complete valence, with eight electrons,
accounts for their reluctance to react.
 As atoms of a gas they are always MONATOMIC,
they are never diatomic. Their atoms won’t even
bond with each other.
The Noble Gases
 Adventures in literacy: during science classes you will
learn many Latin and Greek prefixes that can help with
word decoding. Because the noble gases are so nonreactive, they were difficult to discover.
Noble Gas
Helium
Neon
Krypton
Xenon
Argon
Greek Meaning
Helios – Sun
Neos – New
Hidden
Stranger
Lazy, inactive
Vocabulary/SAT Word
Heliocentric
Neophyte, Neoclassical
Cryptography, Cryptic
Xenophobic
The Diatomic Gases
 Sometimes called the magic seven
 Br, I, N, Cl, H, O and F
 Form the shape of a seven on the table, except for
Hydrogen
 These gases all form diatomic molecules in order to
become more stable. Several are very reactive,
Nitrogen, N2, is the least reactive and most stable.
 In most cases you can assume that elemental diatomic
gases should be written as Cl2, etc. Particularly if you
have been told in words that it is its gaseous form.
Understanding Patterns in
Compound Formation
 We noted earlier that elements in groups 1, 2, 13, 16 and
17 tend to pick up certain charges.
 These elements will pair with one another when
reacting so that their charges balance/cancel out to
zero.
 You will sometimes be asked to identify an element
based on it’s charge.
 You will be given main charges, label them on your
table.
Balancing Charge in Compound
Formation
What element could form a
compound with the formula
XCl?
What element could form a
compound with the formula
MgX2 ? How about X2O?
Strength Does not lie in what
you have. It lies in what you can
give.
Quote I found on the tag of a herbal tea bag. Of course it
reminded me of the metals.
Remember Metals react by giving
up electrons!!!!
Elemental Properties
Metals
 Shiny..have metallic luster
 Lose electrons to acquire




positive charges
Conduct heat and electricity
Are malleable and ductile,
(can be made into sheets and
wires)
Have high melting and
boiling points
React with acids to form
Hydrogen gas
Non-metals
 No luster, may be glassy
 Gain electrons to acquire
negative charges, (exception
– Hydrogen).
 DO NOT conduct, are
insulators because they
prevent heat and electricity
from passing through
 Tend to have low melting and
boiling points
 Form molecules and are
found in molecular
substances
Properties and Reactivity
 Francium represents the
 Fluorine represents the
most reactive metallic
element, so it has the
greatest metallic character
in its group
 Francium is the most
reactive metal in this group
for the same reason, it is
the element that will most
readily lose its valence
electron
most reactive non-metallic
element, so it has the least
metallic character of the
elements in its group.
 Fluorine is the most
reactive non-metal for the
same reason. Fluorine will
literally tear an electron off
of pretty much any
element.
So metals are reacting by giving up an electron and non-metals are reacting
by taking an electron.
States of Matter
 You will need to know the following about what




elements are solids, liquids, and gases at room
temperature.
Most metals are solids, except Mercury, (Hg), which is
a liquid.
Many non-metals are gases, however NOT ALL OF
THEM. C, P, S, Se, Te, Po and, At are all solids.
At standard conditions the Halogens are not all gases.
Bromine is a liquid, and Iodine is a purple crystalline
solid.
The Noble gases are of course, gases.
Bromine is a liquid, however
it will begin to evaporate
and form dangerous vapor
immediately, see the orange
cloud?
Fluorine and Chlorine are
both pale yellow/green
gases
Iodine forms an almost metallic crystal, which is
also very volatile, and sublimes easily into a purple
vapor.
What are Allotropes?
 Remember: the state and the physical properties of a
substance are dependent on the arrangement of its
atoms or molecules.
 Allotropes are forms of elements which have adopted
different molecular/crystalline structure. Despite still
being pure samples of the element, the differences in
structure result in differences in their PHYSICAL
properties.
Carbon
3 Forms of Phosphorus
What is Metallic Character?
 We sometimes refer to the elements as
having more or less metallic character as you
move around the table.
 Essentially the closer you are getting to the
metals, the more metallic character the
elements will have. Therefore they will give
up electrons more easily, etc.
Characteristics of Metalloids
 As you cross from the metals to the non-metals you will
encounter the metalloids, also referred to as the semimetals.
 Metalloids have hybrid characteristics. In other words, a
metalloid may be brittle, but might conduct electricity.
They are literally in between.
 The properties of metalloids allows elements like silicon, to
be used in electrical applications where metals could not
function, for example, in microchips. This is why the birth
place of many computer companies in California is referred
to as Silicon Valley. They are also used to produce solar
cells, solid state data storage, transistors etc.etc.
Navigating By Landmarks
Effective Nuclear Charge
 What type of charge is found on the nucleus
of an atom?
 How does this effect its electrons?
 How does the number of electrons, or the
number of energy levels affect the atoms
structure?
 How does it affect the way the atom will
behave?
Nuclear Charge VS Energy Levels
 The amount of nuclear charge, which is basically the
number of protons, affects how many electrons an
atom can hold onto.
 More importantly, it affects HOW WELL the atom is
able to hold on and the AMOUNT OF ENERGY it
takes to do so.
 You may recall that energy levels that lie closer to the
nucleus tend to be lower in energy, while electrons
located further away from the nucleus are in higher
energy levels.
Nuclear Charge VS Energy Levels
 The larger an atom becomes, the more energy levels it
will have to have to house all of its electrons.
 The more electrons there are the more difficult it
becomes to pull them close to the nucleus.
 HOWEVER… as we said earlier, all the elements
located in a group will have the SAME number of
energy levels. SO…
Despite having more electrons, they are not
moving into higher energy levels and are no more
difficult to hold onto.
Nuclear Charge VS Energy Levels
 But… as you move down a group, you will see that atoms
begin to pick up more energy levels. This makes it
increasingly difficult to hang onto electrons, or to pull
them in close to the nucleus.
 Additionally, the effect of more electrons is that they are
interfering with the electromagnetic attraction of the
nucleus reaching the outermost electrons. This effect is
known as SHEILDING. Think of it as the electromagnetic
force being partially used up before it can reach the outer
electrons at full force. This means the outermost electrons
of bigger atoms will be more vulnerable, and will bond, or
be lost, more readily to achieve stability.
HOMEWORK
 1) Re-label your blank periodic tables again!!!
Take new ones if needed. You must be able to
label all the important information: charge,
group names, increasing and decreasing
metallic character.
 2) Make a table that can be used as a study
guide. Do it on the computer or design by
hand. The table should go over the properties
of the metals, non-metals and metalloids,
including information about major groups
found in those areas!!! Be ready for tomorrow I
will be checking it and collecting it!!!
What is Electronegativity?
 Have a working definition ready…
What’s a working definition?
- A working definition is an informal way of
considering a concept that is more useful for
you to use to both navigate questions and
apply the concept while working.
What is Electronegativity?
 For Electronegativity you need to have in
your working definition:
- It is a tendency to attract electrons.
- It varies across the table according to a
pattern or trend.
- It is impacted by shielding and
effective nuclear charge.
Can you predict the
electronegativity trend
across the periods of the
table?
Navigating By Landmarks
Remember your landmarks!!
Fluorine Just happens to be the
MOST ELECTRONEGATIVE
element on the table
Whenever you are headed towards fluorine,
the elements are increasing in
electronegativity.
Ionization Energy
 What is your working definition of
Ionization Energy?
- It should include energy and the
loss of electrons
- It should include the pattern you find
moving from metals to non-metals on the
table.
Can you predict the trend for
Ionization Energy across a
period?
Remember your landmarks!!
If Fluorine is the MOST ELECTRONEGATIVE
element,
it will take the MOST ENERGY to remove an
electron from a fluorine atom.
So as you are moving towards
Fluorine, Ionization energy is also
increasing.
Size VS Radius
 Before we talk about radius, it’s
important to remember that we are
talking about size –
IN TERMS OF VOLUME…
 In other words, how much space is this
atom taking up? How far out are its
electrons spreading?
Size VS Radius
 During a test, when you are nervous, if you
forget this trend, you may be tempted to
look at the MASS of the elements as you go
across a period, and assume that more
mass means the atom is getting bigger.
 THIS IS NOT TRUE, remember most of an
atom is empty space, (thanks Rutherford ).
We are not talking about MASS we are
talking about HOW FAR THE ATOM’S
ELECTRONS ARE GOING TO SPREAD
OUT!!!!
Radius
Radius
Ionic Radius
 Recall that ions are formed by loss or gain of
electrons.
 Losing electrons: Cation formation,
Acquiring a positive charge. Reduces the
radius of an atom.
 Gaining electrons : Anion formation,
Acquiring a negative charge. Increases the
radius of an atom.
Regents Reference Tables
Take out your Regents Reference
Tables, Let’s review the helpful
information found here!!!
Label the Table
 Now that you have graphed, and know the
major trends found in the groups and
periods, you should add arrows that indicate
these trends to your periodic table.
 Do this tonight!!! Include Electronegativity,
ionization energy, and atomic and ionic
radius. By tomorrow you should have
produced a table showing ALL the
important things you need to know about
the periodic table.