Chemical Formulas and Bonding

Contributed by:

Jonathan James Fri, Jan 14, 2022 07:48 PM UTC

The highlights are:
1. Ionic bonding
2. The octet rule
3. Lewis dot structure
4. Types of ions
5. Properties of ionic compounds
6. Polarity
7. Special type of bonding
8. Exceptions to the octet rule
9. Naming compounds and acids

1. Chapter 7
“Chemical Formulas and
Bonding”
“How it all sticks together….”
T. Witherup 11/06

2. Some Questions to Consider….
 Why are so few elements (such as Au, Ag,
S, N, O) found in Nature in their free
atomic state?
 Why do atoms of different elements react
to form compounds?
 What is happening in this process?
 How can we explain the millions of
compounds that are known today?
 Answers to these questions will be found
in Chapter 7 (“Chemical Formulas and
Bonding”).

3. Chapter 7 Objectives
 Describe the characteristics of an ionic
bond.
 State and use the “Octet Rule.”
 Learn how to use “Lewis Dot” diagrams.
 Learn the types of ions.
 Describe the characteristics of a covalent
bond.
 Describe the difference between ‘polar’
and ‘non-polar’ covalent bonds.
 Write names for ionic compounds,
molecular compounds and acids.

4. 7-1 Ionic Bonding
 What’s an ‘ION’?
 An atom or group of atoms having a charge.
 Do you remember how ions form?
 Metals lose electrons to become
positive ions, called cations. (Which
electrons do they lose?)
M  M1+ + e1-
 Nonmetals gain electrons to become
negative ions, called anions. (Where
do the new electrons go?)
X + e1-  X1-

5. 7-1 Ionic Bonding (cont’d)
 Positively charged ions are attracted to
negatively charged ions.
 Why?
 Because ‘opposites attract.’
 Ionic Compound: A substance that is
composed entirely of ions.
 An ionic formula is the simplest whole-number
ratio of the ions, so the total charge balances
to zero.
 Total (+) charges & total (-) charges = Zero

6. Ionic Compound (General Example)
+ and - combine to form
an Ionic Compound
Cations Anions
These ions are held together in a solid by electrostatic attraction:
- + - + - + - +
+ - + - + - + -
- + - + - + - +

7. Ionic Bonding (Specific Example)
 Sodium (Na) is a poisonous, very reactive
metal.
 Chlorine (Cl2) is a poisonous, very reactive
nonmetal.
 They combine violently to form ordinary table
salt, NaCl, which is relatively harmless.
 NaCl is composed of Na1+ and Cl1- ions.
 Na  Na1+ + e1-
 Cl + e1-  Cl-
 Overall: Na + Cl  NaCl (Note the 1:1 ratio.)

8. The Octet Rule
 Atoms tend to gain, lose or share electrons in order to
acquire a full set (8) of valence electrons.
Na = [Ne]3s1
Loses a 3s1 electron 1+
to form Na1+
(Na1+)
([Ne] electron core).
e1-
e1-
Cl = [Ne]3s23p5 e1- Gains an electron 1-
e 1- e 1-
in 3p to form
e1- Cl1- (3s23p6) ([Ar]
(Cl1-)
electron core). e1- e1-
e 1-
e1-
e1- e1- e1- e1-
e1- e1-

9. The Role of Valence Electrons
 Note that only the valence electrons were
involved in this change, NOT the core
electrons.
 Why? (Which orbitals & electrons are
encountered first when two atoms interact?)
 Chemists focus on the valence electrons
(outer electrons) to understand the
chemistry of atoms.
 To aid us, we use shorthand diagrams,
called Lewis Dot Diagrams, where dots
represent the valence electrons around an
atom.
 Let’s do some examples.

10. Lewis Dot Diagram Method
 Write the element symbol.
 Use dots to show the valence electrons
(alone or in pairs) around the symbol.
 Sodium would be Na with one dot.
 Chlorine would be Cl with seven dots.
 Our previous reaction of sodium with
chlorine would be written as
¨ ¨ .¨
Na. + .Cl:  Na. .Cl:  Na1+ + .Cl:1-
¨ ¨ ¨

11. Lewis Dot Diagrams (Practice)
Element Electron Configuration Lewis Dot
Diagram
Li [He]2s1
Be [He]2s2
B [He]2s22p1
C [He]2s22p2
N [He]2s22p3
O [He]2s22p4
F [He]2s22p5
Ne [He]2s22p6
Al [Ne]3s23p1
P [Ne]3s23p3
Practice doing this! Remember, show only the valence electrons.

12. Types of Ions
 Monoatomic Cations
 Na1+, Mg2+, Al3+
 Fe2+ [Iron(II)], Fe3+ [Iron(III)]
 Monoatomic Anions
 F1-, Cl1-, Br1-
 Polyatomic Ions
 NH 1+, OH1-, NO31-, SO42-, CO32-, PO43-
4
 See list of ions you MUST learn!
 Pages 231 & 232
 http://
www.ausetute.com.au/wriiform.html

13. Facts About Ionic Compounds
 Binary Ionic Compound - contains ions
of only two elements. (e.g. NaCl, CaBr2)
 Empirical Formula – the formula of a
compound with the lowest whole-number
ratio of the elements.
 NaCl (NOT Na2Cl2 or Na3Cl3 or Na100Cl100)
 The “net charge” of a neutral compound
must equal zero, which tells us the ion
ratio. (Ca2+ & Cl1- needs CaCl2 as the
correct formula.)

14. Rules for Writing Ionic Formulas
 Use the simplest whole number ratio of Cation
and Anion.
 Since the net charge must be zero, balance the
number of cations and anions so the total
positive charge equals the total negative
charge.
 Use subscripts after each ion to indicate how
many are present. (Omit ‘1’ though.)
 Use parentheses around polyatomic ions and
indicate their number with a subscript outside
the parenthesis.
 Crisscross method helps write ionic formulas.
 See the next slides.

15. Crisscross Method for Writing Ionic
Compound Formulas
 Ionic compounds must have a net ionic charge of
zero (neutral).
 The total + and – charges must cancel.
 Always keep polyatomic ions intact!
 Use ‘crisscross’ method to write formulas.
 The charge superscript becomes the subscript of the
opposite ion, indicating the number of ions.
 Ba2+ & Br1- becomes BaBr2 [2+ with 2(1-)] = 0
 Al3+ & NO31- becomes Al(NO3)3 [3+ with 3(1-)] = 0
 NH41+ and SO42- becomes (NH4)2SO4 [2(1+) with 2-] = 0

16. Crisscross Method Examples
Barium bromide:
Ba2+ Br1- becomes BaBr2
Aluminum nitrate:
Al3+ NO31- becomes Al(NO3)3
Notice that ‘1’ is not written, that the nitrate
ion is kept intact, and that the net charge is zero.
(For example, barium bromide, 1(2+) + 2(1-) = 0)

17. Crisscross Method More Examples
Ammonium sulfate:
NH41+ SO42- becomes
(NH4)2SO4
Notice how the parentheses are used.
Aluminum oxide:
Al3+ O2- becomes Al2O3
Notice how the net charge is zero. [2(3+) + 3(2-) = 0
PRACTICE, PRACTICE, PRACTICE!

18. Naming Ionic Compounds
 Chemists name compounds on the basis of the atoms and
bonds present.
 Ionic compounds are named from their elements or
polyatomic ions.
 Cations (+) are named first (usually an element name).
 If it can have more than one charge, use Roman numerals to
indicate which ion is actually present.
 FeCl3 is iron(III) chloride & FeCl2 is iron(II) chloride.
 Change the ending of the anion to ‘ide’ (unless a
polyatomic ion is present).
 NaCl is sodium chloride.
 Al2O3 is aluminum oxide.
 Ba(NO3)2 is barium nitrate.
 K2SO4 is potassium sulfate.
 What is NiBr2? Sr3(PO4)2? FeI2?

19. Hydrates
 Hydrate – Ionic compound that absorbs water into their
crystals.
 Blue copper sulfate contains several water molecules in its
crystal. We will do a lab about this.
 Anhydrous – A water-free substance.
 These ionic compounds are named to reflect the ‘water of
hydration.’
 Name the compound in the normal way.
 Add the word ‘hydrate’ and a prefix term to show the number
of water molecules (degree of hydration).
 See Fig. 7-24 on page 246.
 Di-, tri- tetra-, penta- etc.
 MgSO4 *7 H2O is magnesium sulfate heptahydrate.
 What is the formula for copper(II) sulfate pentahydrate?

20. Properties of Ionic Compounds
 High melting points (usually).
 NaF (996 °C), NaCl (801 °C)
 This indicates very strong ionic bonding.
 Very brittle.
 Shatter, or cleave, in fixed paths rather than randomly.
 Example: Rock salt.
 Water soluble (usually).
 Water breaks the ionic bonds.
 Aqueous solutions conduct electricity because the ions
are free to move about in the water.
 Conduct electricity when molten (liquid).
 Ions are freed from the crystal structure (lattice).
 Do not conduct electricity when solid.
 Ions are held firmly in place, so they simply vibrate.

21. 7-2 Covalent Bonding
 A covalent bond is formed by a shared pair of
electrons between two atoms.
 Molecule – group of atoms united by a covalent
bond.
 Molecular Substance – a material made up of
molecules.
 Empirical Formula - the formula of a compound
with the lowest whole-number ratio of the elements.
 Molecular Formula – chemical description of a
molecular compound or molecule.
 Structural Formula – a formula that specifies which
atoms are bonded to each other in a molecule.
 Lewis Structures – molecular structure based on
Lewis Dot diagrams.

22. Covalent Bond Formation
Sharing of electrons, as in two chlorine atoms!
.. combines with
..
:Cl.
..
. ..
Cl:
to form a Cl2 molecule by sharing electrons.
.. ..
:Cl:Cl:
.. ..
This is a “diatomic” molecule, along with molecules of
fluorine, bromine, iodine, hydrogen, nitrogen, and oxygen.
“Professor BrINClHOF” will help you remember them!

23. Describing Covalent Bonds
 Draw Lewis dot diagrams, including unshared
pairs of electrons.
 Use a ‘dash’ for each pair of electrons in a
bond.
 Examples: Chlorine (Cl2) is written as Cl-Cl.
 Single covalent bonds
 C:C or simply C-C (Note the ‘dash.’)
 Double covalent bonds
 C::C or simply C=C (Note the ‘double dash.’)
 Triple covalent bonds
 C:::C or simply CΞC (Note the ‘triple dash.’)

24. Properties of Covalent Compounds
 Low melting points (usually).
 Methane, (CH4) is a gas at room temperature; oils are
liquids at room temperature; wax melts at ~100°C.
 This indicates very weak molecular association.
 Soft.
 Wax feels slippery and may be deformed even as a solid.
 Insoluble in water (usually).
 Water cannot break the covalent bonds.
 Aqueous solutions do not conduct electricity (no ions are
free to move about in the water).
 Do not conduct electricity when molten (liquid).
 Again, there are no ions to move about.
 Do not conduct electricity when solid.
 No ions!

25. Properties of Covalent Bonds (cont’d)
 Remember “electronegativity”? (What is it?)
 The ability of an atom to attract electrons in a
chemical bond.
 Fr has the lowest (0.7) and F has the highest (4.0)
on the Pauling scale.
 Electronegativity differences (“delta EN” or
∆EN) dictate which atom in a bond more
strongly attracts the electrons.
 See Fig 7-20, page 242, and the following slide.
 Chemists use lower case Greek letter delta
(δ) to mean a “partial” or “small difference.”

26. Polarity
 Refers to the unequal sharing of electrons in
covalent bonds of compounds.
 When both atoms in a bond are identical, they
form NONPOLAR bonds (e.g. Cl2 or F2)
because there is, equal sharing.
 When one atom has higher electronegativity
than the other, it forms a POLAR bond (e.g.
HCl), which means the electrons are not
shared equally.
 We use delta +/- (δ+ or δ-) or arrows (+)
to show polarity of a bond.
H-Cl
|

27. Bond Type by Electronegativity
(Use the electronegativity difference, ∆EN, to predict the bond type.)
∆EN Bond Type
≥ 2.0 Ionic
0.4 to 2.0 Polar Covalent
≤ 0.4 Pure Covalent
(Non-polar
Covalent)
Note that a large ∆EN means that it is an ionic bond.
Electrons have transferred from one atom to another.

28. A Special Type of Bonding
 Metallic Bonding – the force of attraction that
holds metals together.
 Positive metal ions are in a ‘sea of electrons’ (freely
floating valence electrons) that are shared.
 This accounts for metallic properties, such as
electrical conductivity, luster, ductility, malleability.
 Drifting electrons insulate the metal ions from one
another, so the ions can easily slide past each other
when stressed, unlike ionic solids, which shatter
when stressed.

29. Exceptions to the Octet Rule
 Atoms with less than an octet.
 Boron compounds.
 Atoms with more than an octet.
 Atoms with d-electrons, such as sulfur.
 Molecules with an odd number of
electrons.
 So called “Radicals” like nitroxyl, NO.

30. 7-3 Naming Chemical Compounds
 Ionic compounds are named from their
elements or polyatomic ions.
 Hydrates have water in their solid
structure, but anhydrous substances do
not.
 Molecular compounds are named using
prefixes to indicate the number of atom in
the formula.
 Acids have special names that must be
memorized (Fig 7-27, pg 249).
 PRACTICE, PRACTICE, PRACTICE!

31. Naming Molecular Compounds
 Use the element names and prefixes to indicate
the number of atoms in the formula.
 Di-, tri-, tetra-, etc.
 CO is carbon monoxide. (“Mono” is not used for the
first element generally.)
 CO2 is carbon dioxide.
 N2O is dinitrogen monoxide.
 N2O4 is dinitrogen tetroxide. (Not usually ‘tetraoxide’
because it is hard to say!)
 Name these: N2O5. SO3. BF3. PF5
 Many molecular compounds have ‘common
names.’
 Dihydrogen monoxide is ______?
 Trihydrogen mononitride is ‘ammonia.’

32. Naming Common Acids
 Acids are molecular substances that dissolve in
water to produce hydrogen ions (H+).
 Acids have special names that must be
memorized (Fig. 7-27, page 249), but focus on
these and their anions:
 Hydrofluoric, hydrochloric, hydrobromic, hydroiodic,
 Nitric
 Sulfuric
 Carbonic
 Phosphoric
 Acetic

33. Did we meet the Chapter 7 Objectives?
 Describe the characteristics of an ionic
bond.
 State and use the “Octet Rule.”
 Learn how to use “Lewis Dot” diagrams.
 Learn the types of ions.
 Describe the characteristics of a covalent
bond.
 Describe the difference between ‘polar’
and ‘non-polar’ covalent bonds.
 Write names for ionic compounds,
molecular compounds and acids.

Book a Trial Class

Download