The Periodic Table The most fun you can have with a bunch of squares. Take a Moment to make a list...
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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?
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
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?
Slide 2
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.
Slide 3
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!!!
Slide 4
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.
Slide 5
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.
Slide 6
A.E. Beguyer de Chancourtois
Slide 7
Slide 8
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.
Slide 9
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.
Slide 10
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).
Slide 11
Slide 12
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
Slide 13
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!!
Slide 14
Lithium Reacting with Water
Slide 15
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.
Slide 16
Strontium and Barium are more reactive and are stored in
oil.
Slide 17
Magnesium burns and is used for survival and camping type fire
starting.
Slide 18
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.
Slide 19
Transition metals form colored compounds and solutions!
Slide 20
Other Metals
Slide 21
Metalloids
Slide 22
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
Slide 23
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 wont even bond with
each other.
Slide 24
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 non- reactive, they
were difficult to discover. Noble Gas Greek Meaning Vocabulary/SAT
Word Helium Helios Sun Heliocentric Neon Neos New Neophyte,
Neoclassical Krypton Hidden Cryptography, Cryptic Xenon Stranger
Xenophobic Argon Lazy, inactive
Slide 25
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.
Slide 26
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 its charge. You will be
given main charges, label them on your table.
Slide 27
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?
Slide 28
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!!!!
Slide 29
Elemental Properties Metals Non-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 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
Slide 30
Properties and Reactivity Francium 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 Fluorine represents the 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.
Slide 31
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.
Slide 32
Fluorine and Chlorine are both pale yellow/green gases Bromine
is a liquid, however it will begin to evaporate and form dangerous
vapor immediately, see the orange cloud? Iodine forms an almost
metallic crystal, which is also very volatile, and sublimes easily
into a purple vapor.
Slide 33
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.
Slide 34
Carbon 3 Forms of Phosphorus
Slide 35
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.
Slide 36
Characteristics of Metalloids As you cross from the metals to
the non-metals you will encounter the metalloids, also referred to
as the semi- metals. 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.
Slide 37
Navigating By Landmarks
Slide 38
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?
Slide 39
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.
Slide 40
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.
Slide 41
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.
Slide 42
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!!!
Slide 43
What is Electronegativity? Have a working definition ready
Whats 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.
Slide 44
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.
Slide 45
Can you predict the electronegativity trend across the periods
of the table?
Slide 46
Navigating By Landmarks
Slide 47
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.
Slide 48
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.
Slide 49
Can you predict the trend for Ionization Energy across a
period?
Slide 50
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.
Slide 51
Size VS Radius Before we talk about radius, its 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?
Slide 52
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 ATOMS ELECTRONS ARE GOING TO
SPREAD OUT!!!!
Slide 53
Radius
Slide 54
Slide 55
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.
Slide 56
Regents Reference Tables Take out your Regents Reference
Tables, Lets review the helpful information found here!!!
Slide 57
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.