Chapter 15 Notes: Ionic Bonding and Ionic Compounds
Reading Assignment C15!
1. Read C15 pp. 412-435 and while reading, continue building your science vocabulary
table that includes all terms in bold face type and all terms you are unfamiliar with or
unsure of. Your vocabulary table needs to include three things:
1.
The term
2.
The definition
3.
A picture of what the term means to you.
2. On page 431, organize the Concept Map into a form that correctly connects the major
ideas of the chapter.
3. Read C15 pp 413-425. Answer section review questions 1-14 and write a list of 5
precious gems (not diamonds) and/or minerals including their chemical formulas.
Example. Sapphires belong to the mineral class corundum. The formula is Be3Al2SiO6
I.
Atoms and Ions
a. Atoms are electrically neutral- equal number of electrons and protons.
b. Ion- atoms which have either gained or lost electrons- contain a net
charge due to imbalance between protons and electrons.
i. Cation- electrically positive ion- due to a loss of electrons.
ii. Anion- electrically negative ion- due to a gain of electrons.
c. Valence shell - Highest energy level occupied by electrons.
i. Valence electrons - electrons in the valence shell.
1. Elements in the same groups usually have the same number
of valence electrons.
2. The group number for representative elements is the same
as the number valence electrons.
d. Electron dot structures or Lewis-Dot structures (G.N. Lewis)
i. Lewis, G.N., Published The Atom and the Molecule in the Journal
of the American Chemical Society, 1916.
1. Describes the valence electrons in an atom utilizing dots
and symbols (x can be used to represent valence electrons,
especially when representing those involved in chemical
bonds)
2. Drawing Lewis Structures C16.1 p.441.
e. Octet Rule- Atoms have a tendency to gain/lose electrons and achieve
electron configurations of the noble gases.
i. First proposed by Irving Langmuir in 1919. Published in The
structure of atoms and the octet theory of valence. PNAS, 1919.
f. Stable ion configurations - addition or subtraction of electrons from
orbitals that yield electron configurations of the noble gases.- pseudonoble gas configurations.
i. Isoelectric species- atoms/ions that have similar electron
configurations.
g.
Oxidation Number- net charge of the stable ion.
i. Identifying oxidation states:
Periodic Variations of Oxidation Numbers
1. The maximum positive oxidation number for any representative
element is equal to the group number, from +1 (Alkali metals) to
+7 (Halogens). Noble Gases have an oxidation number of 0
2. Metallic elements usually exhibit only positive oxidation
numbers due to low Zeff values
3. The most negative oxidation number for any representative
element is equal to the group number minus eight.
4. Negative oxidation numbers are commonly limited to nonmetals.
Semi-metals are negative only when these elements are
combined with less electronegative elements.
5. Elements commonly exhibit positive oxidation numbers only
when combined with more electronegative elements.
6. With the exception of Carbon, Nitrogen, Oxygen, and Mercury,
each representative element that exhibits multiple oxidation
numbers in its compounds commonly have oxidation numbers
that are either all even or all odd.
II.
Compounds
a. Compound- a pure substance that is composed of two or more elements
that are chemically bonded.
b. Law of Definite Proportions- All samples of the same compound contain
the same elements in the same proportion by mass.
i. Coined by Joseph Proust.
ii. John Dalton used this as a foundation to his Atomic Theory
Calculating Percent Mass
Percent Mass =Mass of the element x
100%
Mass of the sample
c. Law of Multiple Proportions. When two elements reacts to form more
than one compound, the ratios of the element's masses. Can be reduced to
a small whole number.
i. Molecules- a group of atoms that are chemically bonded which act
as an individual unit.
1. Typically formed from covalent bonds.
d. Ionic Bond - chemical bond that occurs by a transfer of electrons and the
resulting electrostatic attraction between the charged particles (ions).
i. Due to larger differences in electronegativities (metal/nonmetal).
1. Electronegativity is the pull on electrons involved in a
covalent bond. Derived by Linus Pauling.
2. Classifying Bond types based on electronegativity
differences.
a. Nonpolar covalent = 0.0 to 0.4
b. Moderately polar covalent = 0.4 to 1.0
c. Very polar covalent = 1.0 to 2.0
d. Ionic = greater than 2.0
ii. Ionic Compounds- Compounds formed from ionic bonds.
1. Characteristics:
a. Brittle Crystalline Solids.
b. High melting points- smaller ions produce higher
melting points.
c. Conduct electricity in the molten state (barring no
decomposition).
d. Dissolve in high-polarity solvents (H2O) which
results in conduction of electricity- electrolyte.
e. Lattice Energy- the energy associated with bonded
ions in a crystal.
i. Lattice Energy is the energy needed to
vaporize the ions in a compound. This is
basically related to the amount of energy.
Released when the ions come together
forming the compound.
Calculating Lattice Energy
Lattice Energy =
k
(qc qa)
d
q is the charge of the cation or
anion
d is the S of the ion radii
Print this off a table of Solubilities of Ionic Compounds
e. Chemical formula- shows that kinds and numbers of atoms in a
compound.
i. Empirical formula- shows the simplest whole-number ration of
atoms in a molecule.
ii. Molecular formula- shows the actual ratio of elements in a
molecule.
iii. Formula Unit- lowest whole-number ratio of ions in an ionic
compound.
f. Percent Ionic Character (Pauling)- Based on D-electronegativities between
cations and anions.
i. Allen Factor (Leland C. Allen - Princeton)- used to predict Percent
Ionic character.
1. High Allen Factors predict ionic compounds ( A.F. > 0.4).
2. Low Allen Factors predict covalent molecules (A.F. < 0.4)
3. Modified Allen Factor equation - Using Allen Factor to
calculate percent ionic character:
a. Percent Ion character = Allen Factor x 1.3 x
100%
Calculating Allen Factor
A.F. =
 e.n. /  e.n.
g. Writing & Naming Ionic Compounds
i. Naming Cations
1. Cations formed from metals have the same name as the
metal
a. Single atom ions with or without multiple oxidative
states are called monatomic ions
2. Cations that express multiple oxidation states (typically the
transition metals) are named using:
a. -ic/-ous- archaic scheme used to differentiate
between an elements ions with only 2 oxidation
states.
b. Roman Numeral Nomenclature- the Roman
numeral is used that corresponds to the oxidation
state.
c. Cations formed from nonmetal atoms have names
that end in –ium.
3. Ions that are created from groups of covalently bonded
atoms are called polyatomic ions
a. Print off a list of Common Polyatomic Ions.
ii. Naming Anions.
1. Monatomic anion name endings are changed to –ide.
a. NOTE: There are some polyatomic anions that also
end in -ide, so don't confuse them.
2. Polyatomic anions containing oxygen atoms are called
oxyanions and have name endings of -ite or –ate.
a. NOTE: -ate anions have 1 more oxygen atom than ite anions.
b. Some groups of anions need prefixes to differentiate
between anions. (Ex. chlorates & chlorites).
III.
3. Polyatomic anions that are made from adding an H+ to an
oxyanion have the prefix hydrogen- or bi-.
iii. Rules for Writing & Naming Ionic Compounds.
1. Cations are always written first.
2. The sum of the oxidation numbers must equal zeroElectrically neutral species.
a. Use subscripts in order to get the correct proportion
of atoms.
b. The criss-cross method (oxidation #s to subscripts)
works if you remember to reduce when necessary.
3. Name the compound using correct naming scheme for the
cation and anion.
a. Binary compounds- Composed of two different
types of elements (metal/nonmetal).
b. Ternary compounds- Composed of more than two
types of elements (compounds containing
polyatomic ions).
c. NOTE: A polyatomic anion name ending does not
change.
Naming Acids.
a. Acid- A compound which donates H+ ions in an aqueous solution. The
hydrogen functions as the cation in the formula.
i. Binary acid- hydrogen is bonded to a single non-metal and
dissolved in water.
1. To name Binary Acids:
a. use the prefix hydrob. followed by the root of the non-metal
c. end with the suffix -ic acid.
d. Examples:
i. HCl(aq) -- hydrochloric acid
ii. HBr(aq) -- hydrobromic acid
iii. HF(aq) -- hydrofluoric acid
iv. HI(aq) -- hydroiodic acid
v. H2Se(aq) -- hydroselenic acid
ii. Ternary acids- hydrogen is bonded to a polyatomic acid and
dissolved in water.
1. To name Ternary Acids:
a. DO NOT use the hydro- prefix
b. use the root of the polyatomic ion
c. change -ate endings to -ic acid or -ite ending to ous acid
d. Examples:
i.
ii.
iii.
iv.
v.
vi.
HNO3(aq) -- nitric acid
HNO2(aq) -- nitrous acid
HClO4(aq) -- perchloric acid
HClO3(aq) -- chloric acid
HClO2(aq) -- chlorous acid
HClO(aq) -- hypochlorous acid
Want to see another approach in the form of flowcharts?
Download and print flowcharts for:
Naming anions
Naming cations
Naming compounds with Acids
Naming compounds