This document provides information on ionic compounds and metals. It discusses how ions are formed through the gain or loss of valence electrons to achieve stable octet configurations. Ionic compounds contain oppositely charged ions that are attracted to each other. Their crystal lattices give them high melting and boiling points. Metals form lattices with cations surrounded by a sea of delocalized electrons, giving them malleability, ductility, and high conductivity.

1. Ionic Compounds and
Metals
Chemistry Unit 6

2. Main Ideas
Ions are formed when atoms gain or lose valence
electrons to achieve a stable octet electron
configuration.
Oppositely charged ion attract each other, forming
electrically neutral ionic compounds.
In written names and formulas for ionic
compounds, the cation appears first, followed by
the anion.
Metals form crystal lattices and can be modeled as
cations surrounded by a “sea’” of freely moving
valence electrons.

3. Ion Formation
Ions are formed when atoms gain or lose
valence electrons to achieve a stable octet
electron configuration.
Goals and Objectives:
Define a chemical bond.
Describe the formation of positive and
negative ions.
Relate ion formation to electron
configuration.

4. Valence Electrons and
Chemical Bonds
Chemical Bond – is a force that holds two
atoms together.
They can form between the positive nucleus
of one atom and the negative valence
electrons of another atom or between two
oppositely charged ions.

5. Valence Electrons and
Chemical Bonds
Atom’s try to form the octet – the stable
arrangement of eight valence electrons in the
outer energy level – by gaining or losing
valence electrons.
The transfer of valence electrons between
two atoms is based on the ionization energy
and electron affinity of the two atoms.
Noble gases- high ionization energy + low
electron affinity = little chemical reactivity.

6. Positive Ion Formation
Cation – positively charged ion
Example:
Sodium atom:
1s2 2s2p6 3s1
Sodium ion:
1s2 2s2 p6 = neon

7. Positive Ion Formation
Metal atoms are reactive because they lose
valence electrons easily.
Group 1: commonly form +1 ions
Group 2: commonly form +2 ions
Group 13: sometimes +3 ions

8. Positive Ion Formation
Transition metal ions have an outer shell of s2
They will lose their s electrons and
occasionally a d electron.
Typically form +2 or +3 ions but can form
greater than +3 ions

9. Positive Ion Formation
Other relatively stable electron arrangements
are referred to as pseudo-noble gas
configurations.
Groups 11-14 will lose electrons to form full
outer shells: s, p, and d.

10. Negative Ion Formation
An anion is a negatively charged ion. (Add “-ide”
to the end of the root atom name.)
Nonmetals easily gain electrons.
Example:
Chlorine atom:
1s2 2s2p6 3s2p5
Chlorine ion:
1s2 2s2p6 3s2p6 = Argon

11. Negative Ion Formation
Nonmetal ions gain the number of electrons
required to fill an octet.
Some nonmetals can gain or lose electrons to
complete an octet.
Phosphorus can gain 3 or lose 5
Group 15 usually gains 3 electrons
Group 16 usually gains 2 electrons
Group 17 usually gains 1 electron

12. Practice Problems
CALM 6:1

13. Ionic Bonds and Ionic
Compounds

14. Ionic Bonds and Ionic
Compounds
Oppositely charged ions attract each other, forming
electrically neutral ionic compounds.
Goals and Objectives:
Describe the formation of ionic bonds and the structure
of ionic compounds.
Generalize about the strength of ionic bonds based on
the physical properties of ionic compounds.
Categorize ionic bond formation as exothermic or
endothermic.

15. Ionic Bond
Ionic Bond is the electrostatic force that
holds oppositely charged particles together.
Ionic Compound is a compound that contains
an ionic bond.
Ionic bonds between metals and oxygen
are called oxides.
Most other ionic compounds are
considered salts.

16. Binary Ionic Compound
Binary Ionic compound is an ionic compound
that contains two different elements.
One metallic cation and a nonmetallic anion.
Examples: NaCl, MgO, KBr, LiF

17. Ionic Bond Formation
Electrons gained and lost in each element
must be equal. (conservation of electrons)
Calcium and Fluorine
Aluminum and Oxygen
Sodium and Chlorine

18. Properties of Ionic
Compounds
Compounds are organized such that a pattern
repeats to balance attraction and repulsion
Total charge of a compound is neutral
Often highly organized
Example: NaCl crystal

19. Crystal Lattice
A crystal lattice is a three dimensional
geometric arrangement of particles.
Each negative ion is surrounded by a
positive ion which results in strong
attractions between ions.
Size and shape are dependent on relative
numbers of ions.

20. Crystal Lattice
Physical Properties:
Characteristics of bond strength – ionic
bonds are relatively strong and take a large
amount of energy to break.
Melting point – high
Boiling point – high
Hardness of material is high: rigid and
brittle solids.

21. Properties of Ionic
Compounds

22. Crystal Lattice
Characteristics of the compound:
Conducts electricity, conditionally
Ions in solid state ionic compounds are locked
in place and they do not have free electrons
in order to conduct electricity.
Ionic compounds that are melted or dissolved
in aqueous solutions have ions that are free
to move and therefore do conduct electricity
Electrolyte – an ionic compound that
conducts electricity in an aqueous solution.

23. Energy and the Ionic
Bond
Formation of ionic compounds forms a
more stable system and therefore reduces
the energy required to sustain it.
Since the creating of bonds lowers energy,
energy is released in the process. The
creation of bonds is said to be Exothermic.
Exothermic – energy is released during a
chemical reaction.

24. Energy and the Ionic
Bond
Breaking of ionic compounds reduces the
stability of a system and therefore increases
the energy required to sustain it .
Since the creating of bonds lowers energy, the
breaking of bonds increases energy and
therefore it is required for the process.
Endothermic – energy is absorbed during a
chemical reaction.

25. Lattice Energy
Lattice energy is the energy required to separate 1
mole of the ions of an ionic compound.
higher the lattice energy the stronger the bond
strength.
Directly related to the size of the ions bonded.
smaller ions form compounds more closely
because attraction increases with decreased
distance.
Also affected by charge of ions
Higher ion charge typically has higher lattice
energy.

26. Lattice Energy
Lattice µ
Q1Q2
r

27. Practice Problems
CALM 6:2

28. Names and Formulas
for Ionic Compounds

29. Names and Formulas
for Ionic Compounds
In written names and formulas for ionic compounds, the
cation appears first, followed by the anion.
Goals and Objectives:
Relate a formula unit of an ionic compound to its
composition.
Write formulas for ionic compounds and oxyanions.
Apply naming conventions to ionic compounds and
oxyanions.

30. Formulas for Ionic
Compounds
A standardized system for naming compounds
was developed for much the same reason as the
SI unit system. This serves as a universal naming
system for communication among the science
community.

31. Formula Unit
A formula unit is the chemical formula for an ionic
compound and represents the simplest ratio of
ions.
MgCl2 not Mg4Cl8
A monoatomic ion is a one atom ion.

32. Oxidation number
Oxidation number (oxidation state) the
charge of a monatomic ion.

33. Formulas for Ionic
Compounds
The symbol for the cation is written first with
the anion second.
Subscripts represent the number of atoms of
each element in a compound.
The total charge must equal zero in an ionic
compound.

34. Polyatomic Ions
Polyatomic ions are made up of more than one
atom.
Formulas for polyatomic ionic compounds
Charge applies to the entire group of atoms.
Parentheses are used if more than one
polyatomic ion is needed to balance a
compound.
Do not change subscripts within the ion group
Example (NH4)O

35. Polyatomic Ions

36. Oxyanion
An oxyanion is a polyatomic negative ion
composed of an element, usually a nonmetal,
bonded to one or more oxygen.

37. Oxyanion Naming Rules

38. Names for Ions and
Ionic Compounds
1. Name the cation followed by the anion.
2. For monatomic cations, use the element
name.
3. For monatomic anions, use the root of the
element with the suffix –ide.

39. Names for Ions and
Ionic Compounds
4. Multiple oxidation states are represented by a
Roman numeral in paranthesis after the cation.
a) This applies to transition metals with more than
one oxidation state and not the Group 1 and 2
cations.
b) Example: FeO is Iron (II) oxide; Fe2O3 Iron (III)
oxide.
5. With a polyatomic ion, name the cation followed by
the name of the polyatomic ion.
a) Example: NaOH is sodium hydroxide.

40. Problem Solving

41. Practice Problems
CALM 6:3

42. Metallic Bonds and the
Properties of Metals

43. Metallic Bonds and the
Properties of Metals
Metals form crystal lattices and can be modeled as
cations surrounded by a “sea” of freely moving valence
electrons.
Goals and Objectives:
Describe a metallic bond.
Relate the electron sea model to the physical
properties of metals.
Define alloys, and categorize them into two basic
types.

44. Metals
Metals are not ionic but share several properties
with ionic compounds.
Metals also form lattices in the solid state,
where 8 to 12 other atoms closely surround each
metal atom.
Within the crowded lattice, the outer energy
levels of metal atoms overlap.

45. Electron Sea Model
The electron sea model proposes that all metal
atoms in a metallic solid contribute their
valence electrons to form a "sea" of electron.
The electrons are free to move around and
are referred to as delocalized electrons,
forming a metallic cation.

46. Metallic Bonds
A metallic bond is the attraction of an metallic
cation for delocalized electrons

47. Properties of Metals
Boiling points are much more extreme than
melting points because of the energy
required to separate atoms from the groups
of cations and electrons.

48. Properties of Metals
Metals are malleable because they can be
hammered into sheets.
Metals are ductile because they can be drawn
into wires.

49. Properties of Metals
Mobile electrons around cations make metals
good conductors of electricity and heat.
As the number of delocalized electrons
increases, so does hardness and strength.

50. Metal Alloys
An alloy is a mixture of elements that has
metallic properties.
The properties of alloys differ from the
elements they contain.

51. Metal Alloys

52. Metal Alloys
Substitutional alloys are formed when some
atoms in the original metallic solid are
replaced by other metals of similar atomic
structure.
Interstitial alloys are formed when small
holes in a metallic crystal are filled with
smaller atoms.

53. Practice Problems
CALM 6:4

54. Accumulating Content
How does the electron configuration of a neutral
element compare to that of its ion configuration?

55. Accumulating Content
How does energy and stability relate to ion
formation and bond formation?

56. Accumulating Content
What are some physical and chemical properties of
metals that are caused by the way they ionize or
bond?

57. Study Guide
Key Concepts

58. Key Concepts
A chemical bond is the force that holds two atoms
together.
Some atoms form ions to gain stability. This stable
configuration involves a complete outer energy
level, usually consisting of eight valence
electrons.
Ions are formed by the loss or gain of valence
electrons.

59. Key Concepts
The number of protons remains unchanged during
ion formation.
Ionic compounds contain ionic bonds formed by
the attraction of oppositely charged ions.
Ions in an ionic compound are arranged in a
repeating pattern known as a crystal lattice

60. Key Concepts
Ionic compound properties are related to
ionic bond strength.
Ionic compounds are electrolytes; they
conduct an electric current in the liquid phase
and in aqueous solution.
Lattice energy is the energy needed to
remove 1 mol of ions from its crystal lattice.

61. Key Concepts
A formula unit gives the ratio of cations to
anions in the ionic compound.
A monatomic ion is formed from one atom.
The charge of a monatomic ion is its
oxidation number.
Roman numerals indicate the oxidation
number of cations having multiple possible
oxidation states.

62. Key Concepts
Polyatomic ions consist of more than one
atom and act as a single unit.
To indicate more than one polyatomic ion in a
chemical formula, place parentheses around
the polyatomic ion and use a subscript.
A metallic bond forms when metal cations
attract freely moving, delocalized valence
electrons.

63. Key Concepts
In the electron sea model, electrons move
through the metallic crystal and are not held
by any particular atom.
The electron sea model explains the physical
properties of metallic solids.
Metal alloys are formed when a metal is
mixed with one or more other elements.

64. Questions
Cations form when atoms _______ electrons.
A. gain
B. lose
C. charge
D. delocalize

65. Questions
What is the repeating pattern of atoms in an ionic solid
called?
A. crystal lattice
B. ionic lattice
C. energy lattice
D. ionic bonding

66. Questions
Give the name of the following: NaClO4
A. sodium hypochlorite
B. sodium chlorite
C. sodium chlorate
D. sodium perchlorate

67. Questions
As the distance between ions in an ionic bond is shortened,
A. the energy to break the bond decreases.
B. the electrostatic attraction decreases.
C. the electrostatic attraction increases.
D. the ionic bond changes to a metallic bond.

68. Questions
An alloy is what type of substance?
A. heterogeneous mixture
B. compound
C. mixture of elements
D. element

69. Questions
Which is NOT true about metallic solids?
A. Metals are shiny.
B. Metals are good conductors of heat and electricity.
C. Metals are ductile.
D. Metals have relatively low boiling points.

70. Questions
Electrons in an atom’s outer most energy level are
referred to as what?
A. ions
B. cations
C. valence electrons
D. noble-gas electrons

71. Questions
What is the oxidation state of copper in Cu(II)Cl2?
A. 1+
B. 2+
C. 2–
D. unable to determine

72. Questions
Which elements naturally occur with a full octet of
valence electrons?
A. alkali metals
B. alkali earth metals
C. halogens
D. noble gases

73. Questions
How many electrons are in a full octet?
A. 10
B. 8
C. 6
D. 4

74. The End