Download Trends on the Periodic Table

The Periodic Table
Chapter 6
Chemistry 112
6.1 Development of the Modern
Periodic Table
 The properties of the elements in the
table repeat in a “periodic” way which is
why the term periodic is used to describe
the table.
 Example: periods of the moon
Antoine Lavoisier was credited with:
Compiling a list of 23 elements including
silver, gold, carbon and oxygen
 In 1864, Newlands noticed if he arranged the
elements by atomic mass, their properties
repeated every eight element. He named this
relationship the Law of Octaves, after the musical
notes that repeat every eight tone. The
acceptance of this law was hampered because the
law did not work for all elements and “octave”
was criticized because many thought the musical
analogy was unscientific.
Lothar Meyer and Dmitri Mendeleev were
the two scientists that demonstrated a
relationship between atomic mass and
elemental properties in 1869. Mendeleev
received the most credit because he
published first.
Mendeleev is known as the “father of the
periodic table”
Mendeleev’s table was widely accepted
because he was able to predict the
existence and properties of undiscovered
elements
Henry Mosley is credited with arranging
the Periodic Table by atomic number,
instead of by atomic mass.
Periodic Law is a periodic repetition of
chemical and physical properties of the
elements when they are arranged by
increasing atomic number.
Group – elements in the same column (also
known as family)
Periods – elements in the same horizontal
row
A = Representative elements
B = transition elements
The three main classifications of elements
are Metals, Non-metals, and metalloids.
Some common properties of metals,
nonmetals, and metalloids are:
Metals = hard, shiny, conduct electricity
and heat, malleable, ductile.
Nonmetals = brittle, dull, poor conductors
Metalloids = properties of metals and
nonmetals
Metals are to the left of the stairstep (not
including the metalloids), except
Hydrogen.
Non-metals are to the right.
All elements touching the line are
metalloids EXCEPT for Aluminum
These definitions will be discussed more
completely later in your notes:
– Atomic radii - Atomic radius measures size of
atoms and is defined by how closely an atom
lies to a neighboring atom.
– Ionization energy - Energy required to remove
an electron from an atom which is used to
overcome the attraction between the positive
nucleus and the negative electron.
– Electronegativity - The relative ability of an
atom to attract electrons in a chemical bond.
Advancement of Chemistry
There are three factors that led to the
advancement of chemistry in the 1800’s:
– The advent of electricity
– Development of a spectrometer
– The industrial revolution
Oxidation Numbers
The positive or negative charge of an ion
Equal to the number of electrons
transferred from an atom to form an ion
Sodium = loses one electron = +1
Chlorine = gains one electron = -1
Why do they act like this????? Think
about how many valence electrons all
atoms want.
Oxidation Numbers of Groups
Alkali metals have 1 valence electron that
they want to LOSE. They all form a +1
charge.
Alkaline earth metals have 2 valence
electrons that they want to LOSE. They
all form a +2 charge.
Transition metals for ions in more than
one way. We will skip the “d-block” for
this chapter and revisit it next unit.
Boron Group has three valence electrons
they want to LOSE. They all form a +3
charge.
Carbon Group has four valence electrons.
It is just as easy to lose four as it is to gain
four electrons. Carbon group forms a +4
or -4 charge. (there are some exceptions
that we will talk about next unit).
 Nitrogen group has five valence electrons.
They want to GAIN three more to make a total
of eight. They all form a -3 charge.
 Oxygen group has six valence electrons and
they want to GAIN two more. They all form a
-2 charge.
 Halogens all have seven valence electrons and
want to GAIN one more. They all form a -1
charge.
 Nobel gases have a full valence shell already
and do not form ions.
Practice Problems: What are the oxidation
numbers of the following ions?
K
 1+
P
 3-
I
 1-
 Ar
0
 Al
 3+
S
 2-
Write the ion with the proper
oxidation number or charge for
each of the following:
Mg
Cs
Al
N
O
Br
Kr
Trends on the Periodic Table
You will be responsible for five trends.
1. Atomic radius
2. Metal reactivity
3. Non-metal reactivity
4. Ionization Energy
5. Electronegativity
Atomic Radius
All trends on the periodic table relate to
atomic radius! They all have to do with
how easy it is to lose or gain electrons.
Atomic radius measures size of atoms
Defined by how closely an atom lies to a
neighboring atom.
Trend for within a period and within a
group….graph to find out!!
Trends for Atomic Radius
 Atomic Radius INCREASES as you go down
the P.T. because there are increasing number of
energy levels (think of it as layers of
fat…getting bigger).
 Atomic Radius DECREASES across the P.T.
because there is more of a positive pull on the
electrons. More of an attraction pulls the
electrons a bit closer. (think about having four
people pull on a rope versus five. Five people
are able to pull harder and will therefore bring
the rope closer)
Hands on a clock:
A way to remember the trends is to relate
them to hands on a clock.
You point an arrow in the direction where
the trend INCREASES.
For atomic Radius, you would have a hand
pointing down and a hand pointing left.
This looks like 9:30 on a clock.
Practice Problems: Which atomic
radius is bigger? (Hint, write he symbols how they appear
on the periodic table…which ever one the arrow points to is bigger)
 Na or S
 Na
 Ba or Ca
 Ba
 K or Ra
 Ra (period matters
more than group)
 Cl or Au
 Au
 As or P
 As
 Fr or Cs
 Fr
Ion Radius
 Atoms that lose electrons to become cations,
always become smaller. (when you lose weight,
you become smaller!) Since there are less
electrons that are repelling each other, they can
settle in a little closer to each other.
 Atoms that gain electrons to become anions,
always become larger. (when you gain weight,
you become bigger). More electrons mean
more repulsion. They have to spread out even
further!
Practice Problems: Which is
SMALLER! (pay attention to the questions)
 Ca atom or Ca ion
(hint…what oxidation number
does Ca form?)
 F atom or F ion
 Fe+2 or Fe+3
 Ca+2 ion is smaller
because it loses 2 e F atom is smaller
because the ion gains
1e Fe+3 is smaller
because it loses 3 ecompared to Fe+2 that
loses 2 e-
Metal Reactivity
 When Metals react, they lose electrons due to the
fact they have a low amount of valence electrons.
 Would it be easier for an electron to be lost if it
was close to the nucleus or far away?
 Electrons are lost easier when they are farther
away from the nucleus since there is less of a
positive pull on them.
 Therefore, the greater the Atomic Radius, the
more reactive the metal is. (easier to lose = more
reactive)
Metal Reactivity Trend
http://video.google.com/videoplay?docid=-2134266654801392897&q=brainiac&hl=en
Blocked by Cobb County but you can watch this at home!
INCREASES DOWN a Group due to the
fact that the electrons are further from the
positive pull of the nucleus and they are
easier to remove.
DECREASES ACROSS a period due to
the fact that there is more of a positive pull
holding onto the electrons and they are
harder to remove.
Which metal is more reactive?
 Na or Cs
 Cs
 Ni or Fe
 Fe
 Au or Cu
 Au
 Ag or K
 K (group always matters
more…alkali metals are more
reactive than transition metals!
Good thing or we wouldn’t be
able to wear jewerly!!)
Non Metal Reactivity
 When Non-metals react, they gain electrons due
to the fact that they have a high number of
valence electrons.
 Would it be easier to gain electrons if they were
closer to the nucleus or farther away?
 Electrons are gained easier if they are closer to
the nucleus due to the positive attraction.
 The trend is exactly opposite of Metal
Reactivity and Atomic Radius.
 Noble Gases are excluded…they hardly ever
react!
Non-Metal Reactivity Trend
DECREASES DOWN a group due to the
fact that the atomic radius is getting larger
and it is harder for the positive nucleus to
attract electrons.
INCREASES ACROSS a period due to the
fact that there is more of a positive pull and
is easier to attract the electrons.
What time would this be on a clock?
Which non-metals is more reactive?
 F or Br
F
 S or Cl
 Cl
 N or He
 N (remember Nobel
gases don’t react!)
Ionization Energy
Energy required to remove an electron
from an atom.
Energy is needed to overcome the
attraction between the positive charge in
the nucleus and the negative charge of the
electron.
What type of atoms want to lose electrons?
Metals…therefore, they have lower I.E.
1st I.E. = Energy to remove the 1st electron
2nd I.E. = Energy to remove the 2nd electron
3rd I.E. = etc.
Trends for I.E.
I.E. DECREASES down a group. It
becomes easier to remove an electron
because there are more energy levels
blocking the positive pull.
I.E. INCREASES across a period. There is
more of a positive pull holding on to the
electrons, therefore, it is harder to remove
them. You are also approaching the Nonmetals, which gain electrons!
When will you see a drastic
increase in I.E.???
 After the element has lost its valence electrons.
 Example: Potassium’s 1st I.E. is low. Its 2nd is
very high.
 Calcium’s 1st and 2nd I.E. is low. Its 3rd is high.
 See page 168 Table 6-2
Practice problems: Which atom
has a lower 1st I.E.?
 Ca or Br
 Cl or I
 Au or Cu
 Ag or Rb
 Mg or I
 Ca
I
 Au
 Rb
 Mg (metals always
have a lower I.E.
than
nonmetals….metals
want to lose,
nonmetals want to
gain)
Electronegativity
 The relative ability of its atoms to attract
electrons in a chemical bond.
 What type of elements want to ATTRACT
electrons??
 Nonmetals! Therefore, nonmetals have a high
E.N. value.
 Numerical value from 0.0 (least) to 4.0 (most)
 Unit is a Pauling
 Noble Gases are excluded!!
Trends for E.N.
Down a group, E.N. DECREASES. More
energy levels makes it harder to attract
electrons
Across a period, E.N. INCREASES. More
and more valence electrons, and less for
the atoms to have a full octet. More
positive pull also. Also, pointing towards
the non-metals.
Which element is more
electronegative?
 F or Be
 F (highest E.N. value
on P.T.)
 S or Po
S
 Sr or Br
 Br (Nonmetals always
higher than metals)
 Kr or As
 As (Nobel gases are
are excluded)
Using E.N. values to calculate
bond types:
Ionic Bond – electrostatic force holds
oppositely charged particles together in an
ionic compound. Formed by losing or
gaining electrons.
Polar Covalent – Bond formed when
electrons are not shared equally.
Non-polar Covalent – bond formed from
the equal sharing of valence electrons.
Like Dissolves Like
Water is a polar covalent solvent (meaning
one end of the molecule has a slight
charge).
– Only Ionic compounds and Polar compounds
will dissolve in water because they also have a
charge.
– Non-polar covalent compounds will not
dissolve in water because it does not have a
charge.
E.N. differences
Subtract the lowest E.N. element in a
compound from the highest E.N. element
(the subscripts do not matter)
Difference greater than 1.70 = IONIC
Difference 0.5 to 1.70 = Polar
Difference less than .5 =Non-polar
Use your book pg. 403 to find the E.N. values of
each element. What kind of bond will the
compounds likely have?
 SO2
 .94 = polar covalent
 BaCl2
 2.27 = ionic
 H2S
 0.3 = non-polar
 Ca2N3
 2.04 = ionic
 As2O3
 1.26 = polar covalent