Chemistry Bonds and Nomenclature

BONDS AND NOMENCLATURE

PART 1

LINDA SUMMITT

TIERHEIM ASSOCIATES

PS1.13)

• Use the periodic table and electronegativity differences of elements

to predict the types of bonds that are formed between atoms
during chemical reactions and write the names of chemical
compounds, including polyatomic ions using the IUPAC criteria.

LEARNING TARGETS

1. Predict the type of
bond formed in a
compound based on
electronegativity
differences.

1

2. Write the names of chemical compounds, including
polyatomic ions using the IUPAC criteria.

• a. Predict the charge of an ion formed by main-group elements.
• b. Use prefixes to write chemical formulas for binary molecular compounds.
• c. Use Stock naming compounds containing metals with varying oxidation states.
• d. Utilize a polyatomic ion list to name and write the formulas for ionic compounds
• e. Identify the seven diatomic elements.

2

4 TYPES OF

CHEMICAL

BONDS

(FOCUS BIOLOGY)

Write 4 types
of bonds and a

brief

description for

each.

INTRODUCTION

• Some atoms, noble gases like helium, neon, and argon,

do not seem to bond with other atoms. They are known
as MONOATOMIC elements.

• Other atoms combine by shifting valence electrons so that

the atoms have complete outer energy levels and become
more stable. This may involve gaining, losing or sharing
electrons.

• Oppositely charged atoms are attracted to each other

or atoms that share electrons are held together by
the shared electrons by a chemical bond.

List 6 Monotomic elements!

YOU WILL NEED TO KNOW…

• Lewis Dot Structures

• Electronegativity

• IONs

• Oxidation numbers

• Valence Electrons

• Ionization Energy

• Electron Affinity

VALENCE ELECTRONS AND VALENCE SHELLS

The outer energy level is called the
VALENCE SHELL

Electrons in the outer energy
level are called VALENCE
ELECTRONS

VALENCE

ELECTRONS

The number of electrons

in the outer VALENCE
shell can be determined
by the location on the

periodic chart.

Group 1A: 1e- Group
2A: 2e- Group

3A: 3e- etc

Group 3B: 3e-

Group 4B: 4e-

What are Valence Electrons? Draw an arrow
and define!

ELECTRONEGATIVITY

•The most electronegative (most able to
attract electrons) are in the upper
right corner (F) and the least
electronegative are in the lower left
corner.

•The electronegativity is a scale of 0-4
with Fluorine being 4.0 the strongest
and Francium being 0.7 the weakest.

Fluorine

e- e
-

• Linus Pauling received the Nobel prize for developing the table that showed that
pattern for electronegativity increasing across the period and decreasing down the

group

low

high

What does electronegativity mean?

ELECTRONEGATIVITY

•Exceptions to the trend in electronegativity in the d
orbital block Group 8, 9 and 10 some are higher than
those after them. These elements are often
hybridized. (more later)

Which element is the most electronegative? The least
electronegative?

Identifying Bond Types

Understanding Oxidation numbers and electronegativity helps to
identify bond types. Watch Professor Dave’s Video and see if his
explanation helps you understand a little better.

•
If the difference in electronegativity is more than 1.7, the bond
will be ionic in nature.

•
If the difference in electronegativity is between 0.4 and 1.7, the
bond will have a polar covalent character.

•
Any difference below 0.4 is considered to be relatively equal and
therefore, covalent.

EXAMPLES OF

BOND Types

Bonding with Hydrogen
Electronegativity 2.1

Oxygen 3.5 3.5 – 2.1 = 1.4
polar covalent

Fluorine 4.0 4.0 – 2.1 = 1.9
ionic

Astatine 2.2 2.2 – 2.1 =
0.1 non-polar

2:31 / 3:32

The Chemical Bond: Covalent vs. Ionic and Polar vs. Nonpolar

https://www.youtube.com/watch?v=PoQjsnQmxok

Use your periodic table of electronegativity to identify these as polar, non=polar, of covalent. Drag them into the

correct column and fill in the electronegativity calculations.

Non-polar < 0.4

Polar 0.4-1.7(Up to 2.0)

Ionic >1.7

Cl -Br ∆En = 3.0 - 2.8 =0.2

H - F ∆En = 2.1 - 4.0 = 1.9

Li-F ∆En = 1.0 - 4.0 =3.0

H-H ∆En = 2.1 - 2.1 =0

C - O ∆En = 2.5 - 3.5 = 1.0

Na-Cl ∆En = 0.9 - 3.5 =2.6

∆En

∆En

∆En

∆En

∆En

∆En

∆En

∆En

∆En

∆En

∆En

∆En

Ca-O

C-Cl4

Mg-O

Na2-S

C-S2

O-O

C-H4

N-H3

Fe2- O3

H2-O

H-At

S- F6

OXIDATION NUMBERS

Atoms in a pure elemental state have an oxidation number of
zero. O2 P4 S8

Compounds have an oxidation number of zero. CO2 H2O

Atoms tend to gain or lose electrons to become stable or inert like the
noble gases that they are closest to. The noble gases are inert or
unreactive.

The oxidation number tells the number of electrons that
each atom must gain or lose to become stable.

OXIDATION

NUMBERS

Hydrogen usually has an oxidation number
of +1 because it loses one 1 electron.
Electrons have a negative charge so
Hydrogen will form a positive ion.

All Group 1 have oxidation number +1

Group 2 have oxidation number +2
and lose 2 electrons

Oxygen (and other group 6A
elements) gains 2 negative
electrons (becomes O-2) with and
oxidation number of -2.

POSITIVE OXIDATION NUMBERS

• +1: H, K, Na, Ag, Hg (Mercurous), Cu (Cuprous), Au, NH4

• +2: Ba, Ca, Co, Mg, Pb, Zn, Hg (Mercuric), Cu (Cupric), Fe

(Ferrous), Mn (Manganous), Sn (Stannous)

• +3: Al, Au (auric), As (Arsenous), Cr, Fe (Ferric), P

(Phosphorous), Sb (Antimonious), Bi (Bismuthus)

• +4: C, Si, Mn (Manganic), Sn (Stannic), Pt, S

• +5: As (Arsenic), P (Phosphoric), Sb (Antimonic), Bi

(Bismuthic)

“ous” endings are

the lower
oxidation

numbers; “ic”

endings are the
higher oxidation

numbers

NEGATIVE

OXIDATION

NUMBERS

• -1: F Fluoride, Cl Chloride, Br

Bromide, I Iodide

• -2: O Oxide, S Sulfide,

• -3: N Nitride, P Phosphide

• -4: C Carbide

Rules for assigning Oxidation Numbers

CATIONS AND

ANIONS

Positive ions are CATIONS

Negative ions are ANIONS

Cations are positive ions that move toward
the cathode (negative terminal) and accept
electrons.

Anions are negative ions that move toward
the anode (positive terminal) and give up
electrons.

CA+ION AND ANIONS( )

What is a Cation? What is an Anion?

Write the definition in your own words!

ELECTRON
AFFINITY

Electron affinity is the energy released when atoms
attract electrons.

Atoms generally release energy when they acquire an
electron. A + e- => A- + energy

Atoms that must be forced to accept an electron
release it immediately. A + e-
+ energy => A-

Electron affinity decreases down the group and
increases across the period generally.

Electron affinity, like other energy, is measured in
kilojoules

GRAPHING
ELECTRON
AFFINITY

IONIZATION

ENERGY

Ionization energy is the energy required to
remove an outer electron from a neutral
atom.

Ionization energy increases across a period.

Ionization energy decrease down a group.

The energy required to remove successive
electrons (2nd and 3rd ionization energy )
increases.

PERIODIC TRENDS

BOND TYPES

Intermolecular
• Ionic

• Covalent

• Polar covalent

• Coordinating Covalent

Intra-molecular
• Dipole-Dipole

• Van der Waals

• London Dispersion

BASICS OF CHEMICAL BONDING
https://www.youtube.com/watch?v=ttEBGT0CMsQ

https://youtu.be/WDc4436vlPs

Practice. Complete the practice
problems.

CHEMICAL BONDS

Chemical bonds occur when there is an
electrical attraction between the nuclei
of one atom and the valence electrons of
another atom.

The atoms arrange themselves to
lower the potential energy of the
particles and become more stable
like the noble gases.

CHEMICAL BONDS

8

Chemical bonds

are the

attractive

forces between

atoms

Atoms form

Chemical Bonds
according to the

OCTET RULE

THE OCTET

RULE

Atoms form chemical bonds with
other atoms in order to complete
the outer energy level (to become
like a noble gas)

Most atoms require an octet or
eight electrons in the outer level in
order to be stable

Hydrogen, & Helium need only two
electrons to complete their outer
energy level. (Li, Be, and B try to
become like He)

OCTET RULE

•O has 6 and needs 2 H has 1 and needs 1 so 2 H share 1
each with O giving O 8 and each H 2

THE OCTET RULE

To become like a noble gas &
require an octet or eight
electrons (or full valence shell)
atoms may share electrons like
oxygen and hydrogen.

Atoms like sodium and
chlorine may gain or lose
those electrons becoming
positively or negatively
charged IONS to become
stable .

Na has 1 valence electron.
It loses that electron. Na+
can bond with a negatively
charge atom

THE OCTET

RULE

• Chemical bonds tend to

form so that by gaining, or
losing, or sharing
electrons, each atom has
an octet (8) electrons in
the outer energy level.

• s2 p6

• Coordinate bonds are an

exception to the octet
rule.

• Expanded valence bonding

with electrons in the d
sublevel can occur with
elements bonding to F, O,
and Cl.

Don’t forget ones
like H, He, Li, Be,
and B only need 2
to have a full outer

shell

WHEN DOES

BONDING OCCUR?

When oppositely
charged ions can

position

themselves to
achieve a lower

energy state

When the intermolecular
proton-proton repulsion,

electron-electron

repulsion, and
electron-proton

attraction can arrange to
lower the total energy of
the system. (atoms align
to minimize the repulsive

forces and attractive

forces)

BOND
ENERGY &
BOND
LENGTH

BOND ENERGY

•

WHO BONDS? WHO
DOESN’T?

• Atoms shift valence electrons to

complete the outer energy level.

• Gaining, losing or sharing to reach

noble gas configuration

• Monatomic elements do not easily

form bonds

TYPES OF CHEMICAL BONDS

• Intramolecular or

Intermolecular bonds

•Bonds form in order to
lower the energy of being
separated atoms

COVALENT COMPOUNDS

Covalent
• Atoms having very similar or

identical electronegativity ( a
difference less than 0.4)
share electrons equally.

NO2 3.04 3.44

O2 3.44 3.44

Polar Covalent
• Atoms share electrons. The

electronegativity difference is
1.7 or less but greater than 0.4.
H2O 2.2 3.44

HCl 2.2 3.16

AlCl3 1.8 3.16

IONIC BONDS

Oppositely
charged ion

Metal

-non-metal

Strong

tight bond

IONIC BONDS
(NON-METALS & METALS)

• When an atom with a high electronegativity

is near an atom with a low electronegativity,
the atom with the high electronegativity will
“steal” electrons to complete its outer
energy level. (electronegativity must differ
by 1.7 or more)

• Electrostatic forces of opposite charges

hold ions together.

• Ex: Na (0.9) + Cl (3.0) => Na+ Cl-

IONIC SOLIDS

•Lattice

•High melting & Melting Boiling
point.

•Form ELECROLYTES (aqueous
solutions that carry electric
current) when dissolved

•Formula unit : simplest whole
number ratio of positive and
negative ions that will combine
so that the charges balance.

• Ex. 1 mole of ionic crystal

compound is formed from gaseous
ions calcium fluoride

Ca2+ + F1- + F1- CaF2

IONIC

COMPOUNDS

• Ionic compounds do

not form a single
molecule. They form
a crystal lattice of
charged particles.

• They tend to have

high melting points
and do not conduct
until they begin to
become molten.

The formula
unit tells the
ratio of the
elements.

USE THE INFORMATION BELOW TO DETERMINE THE TYPE OF BOND AND SHOW YOUR WORK BY GIVING

THE ELECTRONEGATIVITIES OF EACH ELEMENT IN THE BOND.

BINARY

COMPOUNDS

For binary compounds the cation
(positive metal ion) is written first.
They keep their name.

• Potassium, magnesium, sodium etc
• Some d block elements have 2 forms so retain

names like copper I and copper II, iron II and
iron III telling the charge.

The anion (negative ions) change their
ending to “-ide”

• Fluorine –fluoride oxygen – oxide
• Iodine –iodide sulfur -sulfide

BINARY COMPOUNDS

• The formula unit or molecular formula must always be neutral.

• The cation is written first and named first.

• Magnesium chloride 🡺 Mg2+ Cl- 🡺 MgCl2

• It would take 2 of the Cl- to balance the Mg2+

• Remember Al(OH)3

IONIC COMPOUNDS

• Remember oxidation numbers

• Look at electronegativity

• The formula unit must be

neutral

http://www.sliderbase.com/images/referats/119b/(8).PNG
Not for commercial use

NAMING SIMPLE BINARY
IONIC COMPOUNDS

• Potassium + chlorine => potassium chloride KCl

• Magnesium + oxygen => magnesium oxide MgO

• Aluminum + fluorine => aluminum fluoride AlF3

• Lithium + phosphorous => lithium phosphideLi3P

CRISS-CROSS
THE FORMULA

CRISS-CROSS

FORMULA
METHOD

YOU TRY IT.

STOCK

SYSTEM
NAMING

Using Roman numerals to identify the

charge of the ion is known as the Stock
system.

This is used for many of the d block

elements that have several possible
charges or oxidation numbers.

Mercury I & II, Copper I & II, Iron II &

III, Chromium II & III, Lead II & III &
IV, Tin II & IV, Vanadium II & IV

POSITIVE
OXIDATION
NUMBERS

• +1: H, K, Na, Ag, Hg (Mercurous), Cu (Cuprous), Au, NH4

• +2: Ba, Ca, Co, Mg, Pb, Zn, Hg (Mercuric), Cu (Cupric), Fe

(Ferrous), Mn (Manganous), Sn (Stannous)

• +3: Al, Au (auric), As (Arsenous), Cr, Fe (Ferric), P

(Phosphorous), Sb (Antimonious), Bi (Bismuthus)

• +4: C, Si, Mn (Manganic), Sn (Stannic), Pt, S

• +5: As (Arsenic), P (Phosphoric), Sb (Antimonic), Bi

(Bismuthic)

Aurora was the shining dawn
of the Romans and was placed
on the God coins

The moon was the silver
white companion Argentum
Ag

D BLOCK CONFIGURATION HELP

dblock2.gif (388×330) (4college.co.uk)

•https://chem.libretexts.org/@api/deki/files/42891/(Remake_1)_Oxidation_States_for_First_R
ow_Transition_Metals.jpg?revision=1

D ORBITALS TRANSITION METALS & OXIDATION
NUMBERS

• Most transition metals have multiple oxidation states

since transition metals have 5 d-orbitals.

• Keep in mind that negative electrons try to “spread out”.

• Looking at Iron, there are 4 unpaired electrons and 2 paired

electrons in the 3d orbital or 6 possible oxidation states
although we have learned that it is generally +2!

MORE ABOUT
TRANSITION

METALS

https://chem.libretexts.org
/ not for commercial use

TRADITIONAL

NAMES

• Traditional names use Latin suffixes for

endings on anions.

• “-ous” is used for the lower charge

• “ic” is used for the higher charge

• Cuprous – Copper I Stannous – Tin II

• Cupric – Copper II Stannic – Tin IV

• Ferrous – Iron II Plumbous – Lead

II

• Ferric – Iron III Plumbic - Lead IV

PROPERTIES OF

IONIC

COMPOUNDS

• Ionic compounds tend to have

high melting and boiling points
because of the crystal lattice
structures.

• They tend to form

ELECROLYTES (aqueous
solutions that carry electric
current) when dissolved.

POLYATOMIC IONS

• In a polyatomic ions, the atoms are tightly bonded into a single charged unit.

• Many polyatomic ions consist of the same elements “SO4

2-” and “SO3

2-”.

These are called OXOANIONS they differ I the number of oxygen atoms
they contain. The ionic charge stays the same.

• Oxoanions with are named by the number of oxygen atoms:

• Perchlorate

ClO4-1

• Chlorate ClO3-1

• Chlorite

ClO2-1

• Hypochlorite

ClO-1

FORMULAS WITH
POLYATOMIC
IONS

Name each component with the
positive component first.

Ammonium NH4

+ Nitrate NO3

-

Nitrite NO-

2 Sulfate SO4

2-

PhosphatePO4

3-

Ammonium chloride NH4Cl

POLYATOMIC

IONS

Polyatomic Ions are charged groups of atoms
that are covalently bonded together.

Polyatomic ions act as a single atom.

NH4

+ ammonium OH- hydroxide

NO3

- nitrate CO3

-2 carbonate

PO4

3- phosphate C2H3O2

- acetate

SO3

2- sulfate ClO3

- chlorate

THE 7 TYPES OF CHEMICAL BONDS

NAMING IONIC

AND MOLECULAR

COMPOUNDS