The Periodic Table Basics &
Naming & Formulas of Compounds
Chemistry-CP
Periods 4 & 6
J.W. Dobereiner


In 1860, there were only 63 elements known
Classified some elements into triads--groups
of three


Triads had: Similar chemical properties & physical
properties that varied in an orderly way
Important because: He grouped elements with
similar properties revealing an orderly pattern in
the elements’ properties.
Examples of Triads



Halogen Triad: Chlorine, Bromine and
Iodine
Coinage Triad: Copper, Silver & Gold
Metal Triad: Calcium, Strontium &
Barium
J.A.R. Newlands (1865)


Realized that when the elements were
arranged by increasing atomic mass,
the properties of the 8th element were
similar to the 1st element.
Law of Octaves: The periodic pattern
repeats itself every 8 elements
Dmitri Mendeleev



Russian chemist who developed the first
periodic table
Listed the elements according to atomic
mass
Important because: He showed the
properties of the elements repeat in an
orderly way from row to row of the
table


Periodicity: the tendency to
recur at regular intervals
Things that are periodic:

Mendeleev’s periodic table was so successful because
it allowed him to predict the properties of still
unknown elements


Eka-Aluminum (Gallium)
Eka-Silicon (Germanium)
Lothar Meyer (1869)

Published almost the same element
classification scheme as Mendeleev but
did not receive credit because
Mendeleev revealed his first and
Mendeleev was more successful at
demonstrating its usefulness
Henry Moseley





Realized that the periodic table was not in the
perfect order
Arranged the modern periodic table.
Listed the elements according to atomic
number
Important because: once arranged by atomic
number all the elements were in order by
their chemical & physical properties
The modern periodic table is listed in order of
atomic number
Periodic Law

The physical and chemical properties of
the elements repeat in a regular pattern
when they are arranged in order of
increasing atomic number
Periodic Table

Arrangement of the
elements in order of
their atomic numbers
so that the elements
are periodic functions
of their atomic
numbers.
Element Key:
•Includes the element symbol, element name,
atomic mass and atomic number
•May include other information
Groups
(also called Families)


The vertical columns on the periodic table
There are 18 groups, labeled with the numbers 1-18.
1
18
2
13 14 15 16 17
3 4
5
6
7
8 9
10 11 12
Group Names
A
l
k
a
l
i
M
e
t
a
l
s
A
l
k
a
l
i
n
e
E
a
r
t
h
Transition
Metals
M
e
t.
Lanthanides
Actinides
B
o
r
o
n
C
a
r
b
o
n
G
r
o
u
p
G
r
o
u
p
N
i
t
r
o
g
e
n
O
x
y
g
e
n
N
o
H b
a
l
l
o e
g G
e a
G n
s
r
s
G o
r
u
o
u p
p
e
s
Periods


Horizontal Rows on the Periodic Table
There are 7 periods labeled with the numbers 1-7.
1
2
3
4
5
6
7
Examples
States of Matter (at Room Temp.)
Solids: Black lettering on
the wall periodic table
Liquids: Blue lettering on the
wall periodic table (Hg & Br)
Gases: Red lettering on the
wall periodic table (noble
gases, F, Cl, O, N, H)
METALS
Left of the zig-zag line
Exception: Hydrogen
PROPERTIES OF METALS






Typically solids at room temperature
Good conductors of heat & electricity
High melting points
Luster (shiny)
Malleable (can be hammered into sheets)
Ductile (can be pulled into wires)
Nonmetals-Located right of the zig-zag line
Exception: hydrogen
Nonmetals






Make up 99% of Earth’s atmosphere
(Oxygen & Nitrogen)
Do not conduct electricity and poor
conductors of heat
Brittle when solids
Many are gases at room temperature
Lack luster
Low melting points
Metalloids
Elements bordered by the zig-zag line
(exception: Al is a metal)
Metalloids

Properties of the Metalloids


Have some chemical and physical
properties of metals and other properties
of nonmetals
Some are semiconductors

Semiconductor: An element that does not
conduct electricity as well as a metal but does
conduct slightly better than a nonmetal

Computers, Handheld electronic devices, calculators
Radioactive Elements

Elements with atomic numbers higher
than 82


Radioactivity: Spontaneous emission of
radiation
Elements are radioactive because they
have too many or too few neutrons

The protons in the nucleus naturally repel
each other. The neutrons are the “glue”
that hold the nucleus together.
Synthetic Elements


The synthetic elements are the
elements with the outlined symbols on
the wall periodic table.
Synthetic elements are not found in
nature. They are man-made elements.
Atomic Model
Diagram
ENERGY
Electrons with the most
energy are located farthest
from the nucleus
Electrons with the
lowest energies are
located close to the
nucleus.
ENERGY lEVELS
The 4th energy level
contains a maximum of 32
electrons
Energy level 3 contains a
maximum of 18 electrons
Level 2 contains a
maximum of 8
electrons
Level 1 contains a
maximum of 2
electrons
EXAMPLES
Draw an atomic diagram of:
Hydrogen
Neon
Magnesium
Carbon
Lithium
Sodium
Atomic Model
Diagram
Valence Electrons

Group
1


The # of electrons in the highest (outermost) energy level
2
13
14
15
16
17
18
Transition Metals: The # of valence electrons for a
transition metal can vary due to the closeness of their s & d
sublevels
 Exceptions:
 Silver is always 1 valence electron
 Zinc is always 2 valence electrons
Inner Transition Metals: Typically have 3 valence electrons
Lewis Dot Diagrams



The element symbol, used to represent the element’s
inner level electrons, is surrounded by dots to
represent the element’s valence electrons
The # of dots must equal the # of valence electrons,
no more than 2 dots per side
Remember: The valence electrons can never be
greater than 8, therefore, there can never be more
than 8 dots.
Oxidation
Number

The charge an atom acquires when it
gains or loses electrons to become
stable


Ion: Atom that has a charge due to the
loss or gain of electrons
Octet Rule: Atoms tend to gain, lose or
share electrons so that each atom has a
full outermost energy level which is
typically 8 valence electrons (octet)
ION
• An atom becomes an ion when it
gains or loses electrons
–The protons in an atom never
change—an atom CANNOT gain or
lose protons
ION
An ion does not have equal numbers of
protons and electrons (the positive does
not = the negative)…therefore…
an ION is a CHARGED atom
ATOM
ION
Oxidation
Number

If an element loses electrons, its
oxidation # is a _______________
number because: there are more
positive protons than negative electrons
Ca+ion (a positively charged ion)

Elements with 1-3 valence electrons:


Lose electrons to become stable
Form ions with a positive charge
ATOM vs. CATION
Positively charged proton
Negatively charged electron
Oxidation Number
If an element gains electrons, its oxidation # is a
_______________ number because: there are more
negative electrons than positive protons.
A n ion

Elements with 5-7 valence electrons:



Gain electrons to become stable
Form ions with a negative charge
Elements with 4 valence electrons:


Metals will lose electrons, becoming positive ions
Nonmetals will gain electrons, becoming negative ions
ATOM vs. ANION
Positively charged proton
Negatively charged electron
Oxidation Number
Group
#
Li
Be
B
C
N
O
F
Ne
1
2
13
14
15
16
17
18
Oxidation Number

Transition Metals



Oxidation #s may vary
Except: Ag+1 & Zn+2
Inner Transition Metals:

Typically a +3 Oxidation Number
Examples
CA+IONS

To Name a Cation:

Name the Metal
Transition Metals: Name the Metal followed by
a Roman Numeral in Parentheses to Indicate the
Metal’s charge


Remember: The oxidation number can change for
transition metals, so it is important to indicate the
metal’s charge
Exceptions:

Zinc is always +2 and Silver is always +1 so they are
transition metals that do not require Roman Numerals
ROMAN NUMERALS
1
I
6
VI
2
II
7
VII
3
III
8
VIII
4
IV
9
IX
5
V
10
X
ANIONS

To Name: Change the ending of the
nonmetal’s name to –ide.
Examples: Sulfide, Iodide, Selenide
ION SYMBOL

ElementSymbolOxidation#
 The oxidation # is the charge the atom acquires
when it gains or loses electrons to become stable
(acquire 8 electrons)
OXIDATION NUMBERS ARE PERIODIC
Group
#
1
2
13
14
15
16
17
18
EXAMPLES

Write the ion symbol for the ions formed from
the following elements.
a) Lithium
b) Aluminum
c) Silver
d) Phosphorus
e) Selenium
f) Bromine
MORE EXAMPLES

Name the following ions
a) Fe2+
b) Clc) N3-
d) K+
e) Zn2+
f) P3-
POLYATOMIC IONS
Common Polyatomic Ions—This table is on your periodic tables
-1
-2
-3
Acetate, C2H3O2Bromate, BrO3Chlorate, ClO3Chlorite, ClO2Cyanide, CNBicarbonate, HCO3Hydroxide, OHHypochlorite, ClOIodate, IO3Nitrate, NO3Nitrite, NO2Permanganate, MnO4Perchlorate, ClO4Thiocyanate, SCN-
Carbonate, CO32Chromate, CrO42Dichromate, Cr2O72Oxalate, C2O42Peroxide, O22Sulfate, SO42Sulfite, SO32-
Phosphate, PO43Phosphite, PO33Arsenate, AsO43-
+1
Ammonium, NH4+
Polyatomic Ions

Names typically end in:
-ate or -ite

The only positively charged ion is:
Ammonium (NH4+)

Where should you look to find the polyatomic
ions?
Polyatomic Ion
Chart on your
Periodic Table
What do opposites do?
ATTRACT
A metal & a nonmetal in the same container both
become stable by gaining/losing electrons---what
charges will the ions in the container have?
Metal = +
Nonmetal = -
What will those charges want to do?
ATTRACT—Come together in an ionic bond
IONIC COMPOUNDS

An ionic bond is formed between an
_______________ and a ______________
because:


Opposite charges attract forming a bond
Therefore, ionic bonds form between
__________________ and ________________
because: ____________form ions with positive
charge and _________________ form ions with
negative charge
IONIC COMPOUNDS

Electrons are _______________________ in an
ionic bond because: one atom is trying to lose
electrons to become stable and the other atom is
trying to gain electrons to become stable
EMPIRICAL FORMULA


Chemical formula for an ionic compound
Lowest whole number ratio of ions in an ionic
compound
EMPIRICAL FORMULA
Al2O3
Subscript: # written to the lower right of a
chemical symbol that shows the number of
atoms of that element present in the
compound
3 WAYS TO DETERMINE AN EMPIRICAL FORMULA
1. Use Lewis Dot Diagrams
2. Use charges
3. Use the Crisscross Method
DETERMINING THE EMPIRICAL FORMULA BY USING
1


LEWIS DOT DIAGRAMS TO ILLUSTRATE THE IONIC BOND
Draw the Lewis Dot Diagram for each
element
Determine the number of each element
necessary to make each atom
DETERMINING THE EMPIRICAL FORMULA BY USING
2


THE OXIDATION NUMBERS (CHARGES) OF THE IONS
The overall charge of an ionic compound is 0
Determine the number of each ion necessary
so that the sum of the charges is 0.
DETERMINING THE EMPIRICAL FORMULA BY USING
3


THE CRISSCROSS METHOD
Write the symbol of charge of each ion in the
compound
Crisscross each ions numerical charge down to the
subscript of the other ion
Sr+2
N3-
Sr3N2
To Name an Ionic Compound
To Name an Ionic Compound
Name the metal
If the metal is a transition metal, add
a Roman numeral in parentheses to
indicate it’s charge
Name the nonmetal, changing
its suffix to –ide, or name the
polyatomic ion.
CaCl2
Fe2O3
NaOH
EXAMPLES
To Name an Ionic Compound
Name the metal (or the
polyatomic ion if it is
ammonium).
If the metal is a transition metal, add
a Roman numeral in parentheses to
indicate it’s charge
Name the nonmetal, changing
its suffix to –ide, or the
polyatomic ion (if it’s more
than 2 elements).
MgS
CaSO4
Cu(C2H3O2)2
Examples
Write the chemical formulas for:
Cobalt(II) chloride
Potassium bromide
Manganese(IV) oxide
IONIC VS. COVALENT BONDING
IONIC
www.blobs.org/science/article.php?article=17Remove frame
COVALENT
Covalent (Molecular) Compounds

Compounds formed from covalent bonds

Covalent Bonds: Bonds formed when atoms share
electrons

Covalent bonds are formed between 2 or more _________.

Ionic Compounds are formed between a _________ & _________.
Which type of bond would form between the following elements?
MgCl2
HI
AlN
CO2
F2
SnO2
Covalent Compound

Electrons are ______________ in covalent bonds.

Electrons are _________________ in ionic bonds.
Molecule

A group of atoms united by covalent bonds

Polyatomic Ions are molecules that have charge!
DIATOMIC MOLECULES:
Some elements only exist in nature as molecules consisting of 2
atoms of that element.
•Why? More stable as a pair!
•The 7 diatomic elements are: HONClBrIF
Molecular Substance
Substance made of molecules

Molecular Formula

Tells how many atoms of an element are in a single molecule of a
compound


The chemical formula of a covalent compound is not the lowest whole
number ratio
Different from an EMPIRICAL FORMULA which gives the lowest whole
number ratio of ions in an IONIC COMPOUND
To Name a Molecular Compound


Name the 1st element. If there is more than
one of that element, use the appropriate prefix
(mono- is not used on the first element)
Name the last element using the appropriate
prefix and changing its ending to –ide.

Since there are no metals in molecular
compounds, no Roman numerals are used.
NUMERICAL PREFIXES
1
2
3
4
5
6
7
8
9
10
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
Examples
N2O4
PCl5
Dinitrogen tetroxide
Phosphorous pentachloride
Nitrogen dioxide
NO2
Chemical Formulas for
Molecular Compounds

Use the prefixes to determine the
chemical formula
– Since molecular compounds do not
involve the transfer of electrons, there
are no ions—do not get charges and no
crisscrossing!
Examples
Nitrogen trifluoride
NF3
Diphosphorus pentoxide
P2O5
CCl4
Carbon tetrachloride
HYDRATES
•An ionic compound that contains water
within its crystal structure
•Hydrates contain water molecules that
are either bound to a metal center or
crystallized with the metal complex.
•The water may be evaporated off
Uses of Hydrates
•Hydrates replace the skin's moisture and repair tissue damaged by cold and
dryness.
•Desalination of Water
•Methane Hydrates
HYDRATES
To Identify a hydrate
BaCl2  3H2O
barium chloride trihydrate
To Name a hydrate
 Name the ionic compound followed by the word hydrate
with a prefix to indicate the # of water molecules attached.
– MgSO4  2H2O
To Write the Formula for a Hydrate
 Use the crisscross method to determine the formula of the
ionic compound, followed by a dot, followed by the # of
water molecules indicated by the prefix.
– Copper(II) chloride pentahydrate
Acids


A molecular substance that
dissolves in water to produce H+
ions
The chemical formula starts with H
(we will assume all are dissolved in
water)
To Name a Binary Acid

“Hydro”-root name of anion-“ic Acid”
– Ex: Hydrobromic Acid
To Name an Acid containing
a Polyatomic Ion

Root name of Polyatomic Ion- “ic” Acid
– Ex: Sulfuric Acid
To determine the formula of
an acid



Consider Hydrogen an Ion: H+
Determine the charge of the anion
Crisscross charges
– H+ S2- 
H2S
Name the
following:
MgI2
N2
CaCl2●2H2O
NO2
NaClO
H2 S
Write the
chemical
formula for:
hydroiodic acid
dinitrogen monoxide
magnesium carbonate
disulfur hexaoxide
manganese (IV) oxide
copper(II) chloride monohydrate
More Review: Once you have these correct & checked go to
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And take the on-line quiz. Password is ‘chemistry’.
Answers
Copper(II) iodide
Nitrogen
Calcium chloride dihydrate
Carbon tetrahydride
Sodium hypochlorite
Sulfuric acid
HI
N2O
MgCO3
S2O6
MnO2
CuCl2●H2O
SALT
An ionic compound typically formed from
reacting an acid with a base.
 The term halogen means “salt-former”