The Periodic Table
Periodic Table of Elements
• There are 117 elements (January, 2007)
– Your table contains 113
• 94 of the elements are naturally occurring,
the rest are man-made
• Most of the elements were discovered
between 1735-1843
History/Development
• Development of the table has occurred
over 300 years and continues today
• Dmitri Mendeleev is commonly credited
with creating the periodic table in 1869
Original Table (Mendeleev)
• Classified elements in horizontal rows
based on their atomic mass
• When there was a repeat of properties
the elements were placed in the next
row
– Concluded that similar properties appear at
regular intervals when the elements are
listed in order of increasing atomic mass
Original Table (Mendeleev)
• Elements with similar properties were
located in the same vertical columns, if no
known element had the expected
properties to fit the particular space he left
that space empty
– Assumed that elements not yet discovered
would fit into the empty spaces
– Predicted some properties of these unknown
elements (gallium, scandium, and germanium)
Modern Periodic Law - 1914
• Henry Moselely bombarded
elements with high speed electrons
and they emitted X-rays with a
certain wavelength
• He found that each element differed
by one proton
• Concluded that the regularity, or
periodicity, of the properties is a
function of the atomic number
(Modern Periodic Law)
Questions
1. The modern periodic table is arranged by
increasing _______________________.
2. The original periodic table was arranged
by increasing _____________________.
Organization of the Modern Periodic Table
• The periodic table is arranged in order of
increasing ATOMIC NUMBER
• Horizontal rows are called PERIODS
– There are 7 periods
• Vertical columns are called GROUPS
– There are 18 groups
Periods
• Horizontal Rows
• The number of each period indicates the
principle energy level in which the valence
electrons are located
• The number of valence electrons
increases as you go from left to right
• The properties of the elements change
systematically through a period
Groups (Families)
• Vertical Columns
• The outermost shell of an atom contains
the same number and arrangement of
valence electrons
• Elements that have similar chemical
properties are located in the same
group
Group/Period Examples
1. Are elements in the same period or group
more similar? Explain why.
2. Which elements have the most similar
chemical properties?
a.
b.
c.
d.
K and Na
K and Cl
K and Ca
K and S
3. Why do elements of a given group on the
periodic table show similar chemical
properties?
States of Matter
• Solid
– The majority of the elements are solids
• Liquid
– The only liquids are Hg and Br, which are
found on the right
• Gas
– H, O, N, F, Cl and the Noble Gases (Group
18), located on the right
Nonmetals
Metals
Metals/Nonmetals/Metalloids
• Where are the metals located?
• Where are the nonmetals located?
• Where are the metalloids located?
– List all of the Metalloids
Properties of
Metals and Nonmetals
• Lose electrons to form
positive ions
• Solids at STP
– Except Hg
• High melting and boiling
points
• Good thermal (heat) and
electrical conductors
• Luster (shine)
• Malleable (bendable)
• Ductile – can be made
into wires
• Gain electrons to form
negative ions
• Solids, 1 liquid (Br),
gases at STP
• Low melting and boiling
points
• Poor conductors of heat
and electricity
• Dull
• Brittle
• Not ductile
Metalloids
• Properties of both metals and nonmetals
• B, Si, As, Te, Ge, and Sb (front + middle
back of staircase)
Metal/Nonmetal Questions
1. Atoms of metals tend to
a.
b.
c.
d.
lose electrons to form positive ions
lose electrons to form negative ions
gain electrons to form positive ions
gain electrons to form negative ions
2. The majority of elements on the table are classified as
a. metals
b. nonmetals
c. metalloids
3. Which property is generally characteristic of metallic elements?
a.
b.
c.
d.
low electrical conductivity
high heat conductivity
existence as brittle solids
low melting points
Metal/Nonmetal Examples
4. At room temperature, which substance is the best conductor of electricity?
a.
b.
c.
d.
nitrogen
neon
sulfur
silver
5. Which element is brittle in the solid phase and a poor conductor of electricity?
a.
b.
c.
d.
calcium
strontium
sulfur
copper
6. The majority of elements on the table are in what physical state at STP?
a. solid
b. liquid
c. gas
Atomic Radius
• ½ the distance between any two nuclei
• Given on Reference Table S in picometers
– 1pm = 1x10-12m
Trend within a Period (Left to Right)
• Atomic Radius decreases
• As you move across a period the number of
protons increases, resulting in a stronger
nuclear charge therefore, electrons are pulled
closer to the nucleus
Trend within a Group (Top to Bottom)
• Atomic Radius increases
– Remember period number = PEL
• As you move down a group there are
additional rings, therefore the valence
electrons are further away from the
nucleus, resulting in a larger radius
Atomic Radius Examples
1. Which sequence of elements is arranged in
order of decreasing atomic radii?
a.
b.
c.
d.
Al, Si, P
Li, Na, K
Cl, Br, I
N, C, B
2. What is the radius of Ca?
3. What is the radius of Sr?
4. Explain why Sr has a larger atomic radius than
Ca.
Ionic Radius (IR)
• Radius that results from the loss or gain of
electrons
Metals (Left Side)
• Tend to lose 1 or more electrons when
forming positive ions
• Radius will decrease
• Ionic Radius < Atomic Radius
Ex: Na+ is smaller than Na
Non-metals (Right Side)
• Tend to gain 1 or more electrons when
forming negative ions
• Radius will increase
• Ionic Radius > Atomic Radius
Ex: Cl- is larger than Cl
Ionic Radius Examples
1. Which ion has the largest radius?
a. Na+ b. Mg2+
c. K+
d. Ca2+
2. Which of the following elements has an
ionic radius smaller than its atomic
radius?
a.
Neon b. Nitrogen
c. Sodium
d. Sulfur
3. The Na+ ion has a smaller radius than
the Ne atom, even though they both
contain 10 electrons. Explain why this is
so.
Ionization Energy
• The energy required to remove the most
loosely bound electron from an atom
• Low IE = greater tendency to lose
electrons and form positive ions
• High IE = greater tendency to gain
electrons and form negative ions
• Given on Reference Table S in kilojoules
per mole (kJ/mol)
IE within a Period (Left to Right)
• IE increases
• Number of protons increases, resulting in
a stronger nuclear charge
• The nucleus has a better hold on the
electrons, therefore more energy is
required to remove an electron
IE within a Group (Top to Bottom)
• IE decreases
• The principle energy levels increase, so
the valence electrons are further away
• Protons cannot hold onto the valence
electrons as well, therefore less energy is
required to remove an electron
Electronegativity (e-neg)
• Scale that measures the ability of an atom
to attract electrons from another atom
• Reference Table S
• Scale ranges from 0.0-4.0
• Fluorine is the highest = 4.0
• Difference between electronegativity
between two atoms can be used to
determine the type of bond
Trend within a Period (Left to Right)
• Electronegativity Increases
• More protons, resulting in a stronger
nucleus, therefore the nucleus is better
able to attract electrons
Trend within a Group (Top to
Bottom)
• Electronegativity decreases
• The atom is larger, so the nucleus is
further away from the valence shell,
therefore the nucleus is less able to attract
electrons
IE/e-neg Examples
1. Which element will lose electrons the
easiest?
a. Na
b. Cl
c. K
d. F
2. Which element would be most likely to
gain electrons?
a. Na
b.Cl
c. K
d. F
Reactivity
Metals
• Most reactive = loses electrons the easiest
(low ionization energy)
– Lower left corner of the table (Fr)
Nonmetals
• Most reactive = gains electrons the easiest
(high electronegativity)
– Upper right corner of the table (F, Cl, O) – not
group 18
Periodic Trends Questions
1. What is the ionization energy of K?
2. What is the ionization energy of Ca?
3. Explain why K has a lower ionization
energy than Ca.
4. According to the reference table, which
of the following elements has the
smallest radius?
a. Ni
b. Co
c. Ca
d. K
Periodic Trends Questions
5. An element with high ionization energy would most likely
be?
a. A nonmetal with low electronegativity
b. A nonmetal with high electronegativity
c. A metal with low electronegativity
d. A metal with high electronegativity
6. What happens to S when it becomes S2-?
a. It loses two electrons and the radius increases
b. It loses two electrons and the radius decreases
c. It gains two electrons and the radius increases
d. It gains two electrons and the radius decreases
Periodic Trends Questions
7. Which of the following would have the
largest radius?
a. Na b. Na+1 c. Cl Cl-1
8. Which of the following has the greater
ionization energy, Na or Na+? Explain
your answer.
Groups 1 and 2
• Properties: typical metallic characteristics
• High reactivity (valence electrons are easily lost)
– Only occur in nature as compounds
• Reactivity increases as you move down the
group
• Group 1 elements are more reactive than group
2 elements
• For metals – Low IE = high reactivity (electrons
are easily lost) – Fr is the most reactive
• Na and Water
• K and Water
Examples:
1. Which atom is the most reactive?
a. Na
b. Mg
c. K
d. Ca
2. Which group 15 element has the least
metallic character?
a. N
b. P
c. Asd. Sb
3. Explain why reactivity increases as you
move down group 1.
Group 17 - Halogens
•
•
•
•
Typical Nonmetals
High electronegativity – F is the highest
High ionization energy
Are so reactive that they cannot exist as in the
monoatomic form
• Exist in nature as diatomics (HOFBrINCl)
– F2 and Cl2 are gases
– Br2 is a liquid
– I2 and At2 are solids
• F is the most reactive nonmetal
• For nonmetals – high electronegativity = high
reactivity
Group 18 – Noble Gases
• Exist as gases at STP
• Exist as monatomic molecules (not
combined with anything
Example: He, Ne, Ar
• The outermost ring is complete, therefore
they are VERY STABLE (unreactive)
Groups 3-12
(Transition Elements)
• Hard solids
• High melting points (except mercury)
• Multiple Positive Oxidation States
– They can react with electrons from both s and
d sublevels
– Different numbers of electrons can be lost
• Colored Ions
– Easily excited (since d and s sublevels are
close)
Group Examples
1.
The presence of which ion usually produces a colored
solution?
a. K+
2.
b. Ca(NO3)2
c. Cu(NO3)2
d. Al(NO3)3
b. N
c. Cl
d. O
Which is a solid at STP?
a. F2
5.
d. S2-
Which element at STP exists as monatomic
molecules?
a. Ne
4.
c. Fe2+
Which solution would be colored?
a. KNO3
3.
b. F-
b. Cl2
c. Br2
d. I2
Which element in Period 4 is classified as an active
nonmetal?
a. Ga
b. Ge
c. Br
Kr
Group Examples
6. Which noble gas would most likely form a
compound with fluorine?
a. He
b. Ne
c. Ar
d. Kr
7. Which element in Period 3 is the most reactive
metal?
a. Na
b. Mg
c. N
d. Cl
8. Which element in Group 15 has the most
metallic character?
a. Bi
b. As
c. P
d. N
Group Examples
9. Why is hydrogen not considered to be a
member of Group 1?
10. Why is hydrogen considered to be a
member of Group 1?
11. Why is it unlikely for sodium to form the
Na2+ ion?
Allotropes
• Some nonmetals can exist in 2 or more forms
in the same phase
• Allotropes have different physical and chemical
properties because their atoms are arranged
differently
Examples:
– oxygen and ozone (O2 and O3)
– Graphite and Diamond (carbon)
Diamond
Graphite