New textbook
Pages 145-165
 By
the 1860’s, scientists knew some of the
physical and chemical properties of more
than 60 elements. But, there was no general
system of organizing the elements.
 So, Dmitri Mendeleev, a Russian chemist, was
one of the first scientists to design a way of
organizing the elements.
 In his periodic table, Mendeleev arranged
elements in rows by increasing atomic mass.
 He started a new row each time the chemical
properties of the elements repeated.
 Mendeleev
was able to predict the properties
of undiscovered elements because he noticed
that the properties of elements seemed to
repeat. He left gaps in his list for these
elements.
 There were a few exceptions of elements
that did not quite fit the pattern. Their
masses did not follow in order. Mendeleev
thought that the masses were not accurate
but as it turns out, they were.
 Due to his work, the element, Mendelevium,
was named in his honor.
 Today,
the modern periodic table is
organized by increasing atomic number.
When arranged in this way, the elements that
have similar properties appear at regular
intervals.
 This principle is known as “periodic law”.
 Each row of the periodic table is called a
“period”. There are 7 of them.
 Each column of the of the periodic table is
called a “group”. There are 18 groups.
 As
you move from left to right across a
period, properties such as reactivity and
conductivity change. The elements become
less metallic.
 For each group, all of the elements in that
group have similar chemical properties.
 See color-coded periodic table on pages
148-149 or the back cover of textbook.
 Why
are some elements reactive and others
are not?
 The answer lies in the number of valence
electrons of each element!!
 The periodic trends in the periodic table are
the result of electron arrangement.
 Elements in Group 1 have 1 valence electron.
 Example, lithium and sodium have similar
chemical properties because they both have
1 valence electron.
You can find out how the electrons are arranged
in the atom of an element if you know where the
element is located in the periodic table.
 Color code your periodic table according to the
s, p, d and f blocks.
 Group 1 and 2 have 1 or 2 valence electrons
respectively.
 Group 13 has 3 valence electrons, Group 14 has
4, Group 15 has 5, Group 16 has 6, Group 17 has
7, Group 18 has 8 except for helium which has 2.
 Transition metals in the middle of the table
(groups 3-12) are different. They do not follow
this pattern.

 Atoms
whose s and p orbitals are not filled
may undergo a process called “ionization”.
 Unfilled orbitals cause atoms to be unstable.
 To become more stable, atoms will gain or
lose electrons and become ions.
 Atoms which lose electrons become
positively charged “cations”.
 Atoms which gain electrons become
negatively charged “anions”.
 To write an ion, you write the symbol and
the charge/amount in the upper right corner.
 Group
1 elements tend to lose 1 electron and
Group 2 elements tend to lose 2 electrons.
 Example: Li 1+
Ca 2+
 Losing each electron require a small amount
of energy. Electrons which are closer to the
nucleus are “harder” to lose because they
are attracted to the positive nucleus.
 When elements lose these outermost
electrons, the energy level is “dropped” and
the next energy level down is filled. This
makes the element more stable even though
it is charged.
 Elements
in group 16 have 6 valence
electrons and tend to gain 2 electrons. This
gives them a filled outermost energy level of
8 exactly.
 Elements in group 17 have 7 valence
electrons and tend to gain 1 electron. Again,
this gives them a total of 8 valence
electrons.
 Example: O2- Cl 1 Energy is required to attract electrons. This
is easier for atoms who energy level being
filled is closer to the positive nucleus.
 Besides
being in groups and periods,
elements are further classified by whether
they are a metal, non-metal or
semiconductor.

Color code your periodic table using your text pg 148-149 (lab).
 Metals
are shiny solids (1 liquid – mercury),
that can be stretched (ductile) and shaped
(malleable). They are good conductors of
heat and electricity.
 Metals are located on the left and center of
the periodic table.
 Ex gold, platinum and copper.
 Non-metals
are on the right side of the table.
(except hydrogen).
 They may be solids, liquids or gases at room
temp.
 Solid nonmetals are often dull and brittle.
 They are poor conductors of heat and
electricity. (insulators).
 Ex. Carbon, oxygen, and helium.
 Semiconductors
are located near the
“staircase” line dividing metals from nonmetals.
 There are 6 semiconductors (aka metalloids):
Boron, Silicon, Germanium, Arsenic,
Antimony, and Tellurium.
 These elements have some properties of
metals and non-metals. Some can conduct
electricity under certain conditions.
 Sometimes,
one or more groups in the
periodic table are categorized as being
members of a unit called a “family”.
 Elements in a family have the same number
of valence electrons. This causes them to
have similar chemical properties.
 Group 1 – alkali metals
 Group 2 – alkaline earth metals
 Groups 3-12 – transition metals
 Group 17 – halogens
 Group 18 – noble gases
 Group
1, except for hydrogen, are called
“alkali metals”.
 They are highly reactive and need to be
stored in oil to prevent them from reacting
with moisture in the air.
 They are soft and shiny
 Lose 1 valence electron
 They are not found in nature as uncombined
elements
 Examples: sodium, potassium
 Group
2 is called “alkaline earth metals”.
 In general, they are harder, denser, stronger
and have higher melting points than alkali
metals.
 They have 2 valence electrons and are less
reactive than alkali metals.
 Examples: Beryllium, Magnesium
Groups 3-12 are called “transition metals”.
 These metals have valence electrons that they
lose to become positive metal ions (cations).
 Some elements lose different numbers of
valence electrons causing them to have several
versions of ions.
 Transition metals are much less reactive than
sodium or calcium.
 Except mercury, transition metals tend to be
harder, more dense and have higher melting
points than groups 1 or 2 elements.
 Examples: Copper, iron, gold, uranium,
europium

 Some
elements are made in a laboratory and
are called synthetic.
 All elements that have atomic numbers
greater than 92 are synthetic.
 Some elements are also radioactive, meaning
the nuclei of their atoms are continually
decaying (falling apart) to produce different
elements. Ex. Tc (#43) and Pm (#61)
 Some elements have stable isotopes, like Tc99, which are used to create images during
brain scans.
 Non-metals
are located on the right hand
side of the periodic table, except hydrogen.
 Group 18 elements are called “noble gases”.
 The noble gases are different from most
elements because they exist as single
uncombined atoms instead as molecules or
compounds, aka “inert”. (hence the name
noble)
 Noble gases are stable because they have
filled energy levels, and 8 valence electrons
(except helium that is filled with 2).
 Group
17 elements are called “halogens”.
 These elements are the most reactive nonmetals.
 They have 7 valence electrons and tend to
react with Group 1 elements. When this
happens the compound is called a “salt”.
 Examples: Fluorine, Chlorine, Bromine.
 In
addition to the noble gases and halogens,
there are 6 more non-metals on the right
side of the periodic table. (don’t forget
about hydrogen, which is on the left)
 One very important non-metal is carbon.
Carbon is what makes up graphite (in your
pencil) and diamonds. It can also form
millions of carbon-containing compounds.
 The
six semiconductors are also known as
metalloids and are located along the
“staircase line” dividing the metals and nonmetals.
 As their name suggests, semiconductors can
conduct heat and electricity under certain
conditions.
 The six semiconductors are boron, silicon,
germanium, arsenic, antimony, tellurium.
 Hydrogen,
which has one proton and one
electron, does not behave like any other
elements.
 As a result, hydrogen is in a class by itself.
 Hydrogen is the most abundant element in
the universe.
 Hydrogen can react with many other
elements to form compounds.