PERIODIC TABLE NOTES
(from chapters 5 and 6)
I. History of the Periodic Table
As the number of elements began to grow, chemists needed a way to
______________ all of these elements.
[In the 1700’s there were _______ known elements. As new technologies
were developed many new elements were discovered. In less than _______
years the number of known elements had __________.
A. J.W. Dobereiner:
- Early _________
- Grouped elements into sets of three’s called _______.
- He grouped elements like _____________, Cl, Br, & I; Ca, Sr, & Ba.
- [If you notice these groups on the current periodic table, they are all in
the same row.]
- The elements in the triad had similar ____________ properties.
- The element placed in the middle has properties which are an
_____________ of the other two elements.
B. J.A.R. Newlands:
- __________
- There were _____ known elements at this time.
- He noticed that when the elements were arranged by increasing
__________________, the properties of the _____ element were like
those of the _____, the _____ like the ______, 10th like 3rd, etc.
- He called the pattern he saw the ____________________ because the
pattern repeated every ________ element. Unfortunately his important
recognition of the _____________ (or repetitive) pattern of element
properties was not accepted because of its use of a _________ term
(octave-eight notes which make up a scale).
C. Dmitri Mendeleev:
D. Lothar Meyer:
- _________
- both of these scientists published nearly _______ ways to classify the
elements
**________________ was given credit because he could demonstrate
and ______________ the classification system.
- Mendeleev eventually produced the ___________ periodic table,
basically arranging the elements in order of
_____________________________. It was arranged so that elements in
the same column had ____________ properties. His periodic table is very
similar to the one we use today and because of this he is called the
__________________________________.
- Mendeleev sometimes ______ the pattern of increasing atomic mass to
keep the elements with __________ properties in the same
____________.
- He switched _______ sets of elements (one example is Co and Ni).
[Note: these are out of order according to atomic mass on the periodic
table.] He said that the atomic masses had been measured ___________
and when correctly measured, would match his table.
- Mendeleev was also able to predict the __________ and
_____________ of missing elements. He correctly predicted the
properties of ____________(he called it Ekasilicon) which was later
_____________.
E. The Periodic Law:
Mendeleev was incorrect about the atomic masses being wrong.
In 1913, _____________ did experiments from which he discovered the
_____________________ (the number of protons in the nucleus). He
recognized that the correct way to arrange the elements was by ____________
atomic number (not ___________________ as Mendeleev had done).
Moseley then developed the _______________ which states that when elements
are arranged in order of increasing __________________, their chemical and
physical properties show a ______________ pattern.
II. Organizing the Periodic Table:
[Because of the way that the periodic table is arranged, you can predict an
elements properties by knowing its position on the table and what that position
signifies.]
The modern table has approximately ____________ squares. Each square
includes certain information, like the _____________, element ___________,
atomic ____________, outer electron configuration, element name and other
information depending on the table.
The shape of the perio dic table comes from the _____________. The two rows
at the bottom actually fit into the table, but are placed at the bottom so the table
can fit on one page and is not so ______. See Fig 5-8, pg 164 for a picture of
the table with these two rows included.
A. Groups/Families –
- ___________ columns on the periodic table;
- contain elements with similar _____________
- labeled with _______ designation, or numbered 1-18.
- We will use the A/B designation.
[Please label your periodic table with the groups 1A – 8A (or 0) and 1B – 8B as it
is done on the periodic table on the wall or on page 166-167 (red letters).]
- A groups are referred to as the _________________________
- B groups are called the ______________
- For the A groups, the _____________ number indicates the number of
____________ electron for that group.
B. Periods –
- _________________ rows on the periodic table.
- There are _____ periods on the periodic table.
-The two periods at the bottom are a part of periods ____ (La-Vb) and
____ (Ac-No).
[Please label the periods on your periodic table as in done on the wall or pg. 166167 in your book.]
[There are a number of ways to group elements on the periodic table.]
C. Metals/Nonmetals/Semimetals:
[There is a division line on the periodic table called the stairstep. It is drawn in
red on the wall table and starts on the box which contains B. Draw this on your
own periodic table. This is what we use to identify metals and nonmetals.]
Metals:
- located to the ______ of the stairstep
- are good conductors of ___________ and electricity
- most are ____________ ( can be hammered into thin sheets)
- most are _____________ ( can be drawn into fine wires)
- have _________ ( or shine)
- metals are mostly ___________ at room temperature (except
_____)
Nonmetals:
-
are located to the _______ of the stairstep.
are poor _______________ of heat and electricity
tend to be _______________
-
many are _____________ at room temperature
do not have luster ( or shine)
Semimetals:
- are located _________________ to (or touching) the stairstep.
- have properties of both metals and non metals.
Note: Al is actually a metal by its behavior, but we will call it a
semimetal since it touches the stairstep.
D. Group Names:
Group 1A: Alkali Metals ( does not include H)
- are ______ metals, soft enough to be cut with a knife.
- _______________ rapidly when exposed to air.
- very reactive with _________, even water found in air; become
more reactive as you move down the group.
- stored under ______l to prevent reaction with moisture and
oxygen
- react with ____________ to form salts
- never found _______ in nature
- These metals all have _______ valence electron (outer energy
level electron)
- Readily lose one electron to form an ion with a _____ charge
- Outer electron configuration is __________ followed by
__________.
Group 2A: Alkaline Earth Metals
- Have all of the properties of ____________. Are ____________
as similar as group 1A elements are.
- Have ________ valence electrons
- Readily lose _____ electrons to form an ion with a _____ charge
- Outer electron configuration is the ___________ followed by
_____.
- These elements are never found _________ or uncombined in
nature.
Group 3A: Boron Group
- Contains semimetals and metals
- Have ____ valence electrons
- Readily lose ____ electrons to form an ion with a ____ charge
- Outer electron configuration is the ________ followed by
_______________
- Most important element in this group is Al.
Group 4A: Carbon Group
- Contains nonmetals, semimetals and metals
- Have ____ valence electrons
- Will either _______ valence electrons to become an ion with a
____ charge, or ______ valence electrons to become an ion with
a ____ charge.
- Outer electron configuration is the _______ followed by
___________
- Most important element in this group is C.
Group 5A: Nitrogen Group
- Contains nonmetals, semimetals and metals
- Have ____ valence electrons
- Readily gain ____ electrons to form an ion with a ____ charge
- Outer electron configuration is the _____ followed by
______________
- Most important element in this group is N.
Group 6A: Oxygen Group
- Contains nonmetals, semimetals and metals
- Have _____ valence electrons
- Readily gain ____ electrons to form an ion with a ____ charge
- Outer electron configuration is the _________ followed by
_____________
- Most important element in this group is O.
Group 7A: Halogens
- Contains metals and semimetals
- Have ____ valence electrons
- Readily gain ____ electron to form an ion with a ____ charge
- Outer electron configuration is the ________ followed by
___________________
- All of these elements form as ________________ molecules (F2,
Cl2, Br2, I2, etc).
- These elements are very __________ and therefore do not exist
as __________ elements in nature.
- Some are very dangerous gases, like F and Cl.
Group 8A ( or 0): Noble Gases
- Are very ____________, sometimes called ________ gases.
- Have ______ valence electrons
- Will not __________ or __________ electrons, because they
already have a full set of 8 valence electrons.
- Outer electron configuration is the ___________ followed by
___________________
Group B elements: Transition Metals
- There is a great variety in this group. This group contains many of
the common ____________ metals and metals used in
_____________.
- They will lose a various number of electrons.
- The outer configuration is basically _________________. We will
discuss this more later.
Bottom Two Rows: Inner Transition Metals
- Lanthanides: La-Yb
- Actinides: Ac-No
- There is a great variety in this group.
- They will lose a various number of electrons.
- The outer configuration is basically _______________. We will
discuss this more later.
III. Outer Electron Configurations
Valence electrons are electrons which occupy the ____________ principle
_____________________. These electrons are basically responsible for an
atoms ____________________ behavior.
As we have just discussed, atoms in the same ___________ have the same
number of ___________ electrons. Because of this, all atoms within a certain
group behave in a ___________ way. Since the valence electrons are the ones
responsible for an atoms behavior, these are the electrons which are important
for us to be aware of. In the last chapter we learned how to write the electron
configuration for various atoms. This configuration showed us where each
electron in the atom was located.
Since we are really interested in the valence electrons, we can now write a
_________________ electron configuration which will only indicate these
valence electrons. This configuration is called the outer electron configuration.
This outer electron configuration has 4 parts:
1. The ____________________________________ which is determined
by finding the noble gas (group 8A) in the period ____________ the
period the atom you are working with is in. The noble gas is placed in
______________ brackets. See example below.
2. The ____________________ is determined by the
________________ number if the outer electrons are in the s and p
orbitals. If the electron is in the d orbital, the energy level is the
_________________. If the electron is in the f orbital, the energy is
the ___________________.
3. The ______________________ and 4. The __________________ in
that orbital is determined from the group number.
Recall:
Group 1A: s1
Group 2A: s2
Group 3A: s2 p1
Group 4A: s2 p2
Group 5A: s2 p3
Group 6A: s2 p4
Group 7A: s2 p5
Group 8A: s2 p6
Transition Metals: s2 d1-10. To determine the number for the d
orbital, you count over starting on the left of the B section.
Inner Transition Metals: s2 f1-14 .To determine the number for the f
orbital, you count over starting on the left of the inner transition
metal section.
Example: Write the outer electron configuration for Na.
1. noble gas:
2. outer energy level: comes from the period #.
3 and 4. orbital and number of electrons: comes from group #.
Now lets put it all together:
Example: Write the outer electron configuration for I.
1. noble gas:
2. outer energy level: comes from the period #.
3. 3 and 4. orbital and number of electrons: comes from group #.
Now lets put it all together:
**Note: you must put the energy level # in front of the
s and p.
Example: Write the outer electron configuration for Ni.
1. noble gas:
2. outer energy level: comes from the period #.
3 and 4. orbital and number of electrons: comes from group
( To determine the electron # for the d, we need to count over
starting with Sc to Ni. Ni is the 8th element in the B section, so it will
be d8. Remember the energy level number is also different for d
(period # -1), so we will write 3 d8 in our outer electron
configuration.)
Now lets put it all together:
**Note: you must put the energy level # in front of the
s and d.
Example: Write the outer electron configuration for U.
1. noble gas:
2. outer energy level: comes from the period #.
3 and 4. orbital and number of electrons: comes from group.
(To determine the electron # for the f, we need to count over
starting with Th to U. U is the 3rd element in the inner transition
metal section, so it will be f3. Remember the energy level number
is also different for f (period # -2), so we will write 5 f3 in our outer
electron configuration.)
Now lets put it all together:
**Note: you must put the energy level # in front of the
s and d.
You need to be able to write these outer electron configurations. To check your
work, see the period table on pg.166-167 in your book.
IV. Periodic Trends
Many properties of elements will change in a predictable way as you move
across and down the periodic table. These are called periodic trends.
1. Atomic Radius: is the distance from the center of an atoms nucleus to its
outermost electron.
As you move across a period from L to R, the atomic radius decreases.
As you move down a group from top to bottom, the atomic radius
increases.
Remember that the largest atom is Fr and the smallest atom is F.
Why: moving down a group, atoms have more electrons and more
energy levels, so are bigger.
Moving across a period, electrons are more strongly attached to the
more positive nucleus, so atom is smaller.
2. Ionic Radius: is the distance from the center of an ions nucleus to its
outermost electron.
No convenient trend, but you should remember the following.
When an atom loses electrons (to become a positive ion), the ion is
smaller than the original atom.
When an atom gains electrons ( to become a negative ion), the ion is
larger than the original atom.
3. Ionization Energy: is the energy needed to remove one of an atom’s
electrons, or how strongly an atom holds its outermost electrons.
As you move across a period from L to R, the ionization energy increases.
As you move down a group from top to bottom, the ionization energy
decreases.
Remember that the atom with the highest ionization energy is F and the
smallest ionization energy is Fr.
Why: Larger atoms(Fr) hold electrons less tightly than small atoms
(F), so the electrons in the larger atoms are easier to remove.
4. Electronegativity: reflects an atoms ability to attract electrons, or how much
does an atom want .
As you move across a period from L to R, the electronegativity increases.
As you move down a group from top to bottom, the electronegativity
decreases.
Remember that the atom with the highest electronegativity is F and the
smallest electronegativity is Fr.
Why: Atoms with valence #’s closer to eight want electrons more
and therefore have a higher electronegativity.