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Name: ____________________________________________ Period: __________

JHS KAP CHEMISTRY – IONIC BONDING NOTES

CHEMICAL BONDING OVERVIEW

The focus of this unit is on the IONIC BOND and naming IONIC compounds which are compounds involving atoms

bonded together by ionic bonds. We will briefly introduce metallic and covalent bonds as well.

Chemical Bond

A chemical bond is a force of attraction that holds atoms or ions together.

The Octet Rule

• The octet rule states that atoms form chemical bonds so that each atom has eight electrons (an octet) in its valence

shell (highest occupied energy level).

• Except for helium which

has a total of two electrons,

both in the valence shell,

the noble gases each have

eight electrons in their

valence shell. This is an

energetically stable valence

shell electron configuration.

• Atoms of other elements

achieve the stable electron

configuration of noble

gases by forming chemical

bonds so that each atom

ends up with an octet of

electrons in their valence

shell.

• A few atoms like hydrogen,

lithium, and beryllium will

become stable with two

electrons in their valence

shell like helium.

DEFINITIONS

If a substance is a pure solid, we use the symbol (s). If it is a pure liquid, we use the symbol (l). If it is a pure gas, we

use the symbol (g). When a substance dissolves in water, (is soluble in water), we use the symbol (aq) to show that it

has formed an aqueous solution.

Examples:

Solids: Na(s) NaCl(s) CaCO3(s)

Liquids: H2O(l) Br2(l) C2H5OH(l)

Gases: CO2(g) H2(g) NH3(g)

Aqueous solutions: LiF(aq) KNO3(aq) C6H12O6(aq)

See the flow chart on the next page for more definitions.

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The flowchart below defines the terms atoms, elements, molecules, compounds, mixtures.

Know and understand ALL the definitions.

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TYPES OF CHEMICAL BONDS

There are 3 types of chemical bonds: ionic, metallic and covalent bonds.

• Covalent bonds involve sharing of valence electrons between nonmetals and/or metalloids to achieve an octet.

Molecular compounds contain covalent bonds.

• Metallic bonds involve an electron sea model where the valence electrons move freely (are delocalized) throughout

the metal structure and are attracted to the positive nuclei of the metal atoms. Metals contain metallic bonds.

• Ionic bonds involve the transfer of valence electrons between a metal and nonmetal to achieve an octet. Metals

lose electrons and nonmetals gain electrons during the transfer. (Polyatomic ions may also be involved in ionic

bonding). Ionic compounds contain ionic bonds.

COVALENT BONDING (We will discuss covalent bonding in detail in the next unit. This is only a short overview.)

Sharing Valence Electrons

• Atoms of nonmetals

and/or metalloids share

electrons and form

covalent bonds to become

more stable by achieving

the valence electron

configuration of a noble

gas (an octet of electrons

in their highest occupied

energy level).

• The diagram shows how a

few molecules of

nonmetals share their

electrons to achieve an

octet. We will learn how

to draw those diagrams in

the next unit.

• Compounds consisting of covalently bonded atoms are called molecular compounds.

• The smallest particle of a molecular compound is a molecule.

• The chemical formula for a molecular compound is called the molecular formula.

• Examples of molecular compounds:

Properties of Molecular Compounds

• They can exist as solids, liquids or gases at room temperature.

• The molecules are held together by relatively weak intermolecular forces.

• When liquids, they tend to vaporize easily. This means they are volatile (change from a liquid to a gas at room

temperature).

• They have relatively low boiling and melting points due to weak intermolecular forces between molecules.

• They DO NOT conduct electricity either when pure or when dissolved in water (aqueous) (if they dissolve).

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***NOTE: For a substance to conduct electricity, it must contain mobile charged particles.

Electrons or protons do NOT have free movement in molecular compounds with covalent bonds.

METALLIC BONDING

Electrons Sea Model

• A metallic bond is the electrostatic attraction between the free moving (delocalized)

valence electrons, with a negative electrical charge, and the protons in the nuclei,

with a positive electrical charge.

• Metals have characteristically few valence electrons. They therefore often have

vacant (empty) p and d orbitals. This allows the valence electrons to move freely

throughout the metal and valence electrons are not held by one nucleus.

• Metallic bonding is modeled by a sea of delocalized valence electrons attracted to the

metal cations. This results in a lattice structure of metal atoms.

Properties of Metals

• Metals generally exist as solids at room temperature. Exception

is mercury which is a liquid at room temperature. (Hg(l)).

• Metals are lustrous (shiny in appearance)

• Metals have high melting and boiling points due to strong

attractions between the valence electrons and metal cations in

the lattice structures.

• Metals have high densities.

• The mobile sea of valence electrons accounts for many

properties of metals in their original solid state.

o High thermal (heat) conductivity.

o High electrical conductivity. Metals CAN conduct

electricity in the solid state because the valence

electrons are free to move in the electron sea model.

o Malleability. They can be hammered into different

shapes.

o Ductility. They can be pulled into wires.

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IONIC BONDING

Transfer of valence electrons

In the ionic bond, valence electrons of metal and nonmetal atoms are transferred (gained/lost) to achieve an octet.

The resulting ion is more stable than its parent atom because its electron configuration is more stable and matches that

of a Noble gas. The type of ion an element forms depends on how many valence electrons it has, which you can

determine by examining its position on the periodic table.

• Metals in groups 1,2 and 13 form positive ions, cations, with 1+, 2+ and 3+ charges respectively. These elements

are monovalent because they always form ions with only 1 possible charge.

Examples: Na+, Mg2+, Al3+

• Nonmetals in groups 15, 16 and 17 form negative ions, anions, with 3-, 2-, and 1- charges respectively. These

elements are monovalent because they always form ions with only 1 possible charge.

Examples: P3⁻, O2⁻, F3⁻

• Metals in the d block (groups 3-12) form ions with multiple possible charges. These elements are multivalent.

Roman numerals are used to indicate the charge in the written chemical name.

Examples: copper(I) is Cu+ and copper(II) is Cu2+

iron(II) is Fe2+ and iron(III) is Fe3+

Exceptions in the d block: zinc always forms Zn2+, cadmium always forms Cd2+, and

silver always forms Ag+. Zn, Cd, and Ag are therefore monovalent and DO NOT

have a roman numeral in the written in the name to indicate the charge.

• Tin(Sn) and lead(Pb) in group 14 and bismuth(Bi) in group 15 are also multivalent

metals.

Examples: tin(II) is Sn2+ and tin(IV) is Sn4+

lead(II) is Pb2+ and lead(IV) is Pb4+

We will not include bonding of elements in the f block in this course.

Metals always form positive ions (cations) because they lose valence electrons.

Nonmetals always form negative ions (anions) because they gain valence electrons.

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Examples of Metals and Nonmetals Forming Ions

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Polyatomic Ions

Ionic bonding can also include polyatomic ions.

A polyatomic ion is a group of bonded atoms having either lost or gained electron(s). The resulting ion has either a

positive charge (cation) or a negative charge (anion). The polyatomic ions’ names and formulas in the table below

should be memorized, including the correct charge. More polyatomic ions are included on your reference formula

chart.

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NOTE:

The “ite” ending

indicates one less O

than the “ate” ending

SULFITE SULFATE

NITRITE NITRATE

PHOSPHITE PHOSPHATE

Formation of Ionic Bonds and Ionic Compounds

Positive ions (cations), and negative ions (anions) are attracted to each other by an electrostatic force, also known as

a Coulombic attraction or Coulombic force. This electrostatic attraction is between cations and anions is called an

ionic bond which forms when metals and nonmetals react to produce an ionic compound.

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Examples and Practice:

Draw the Lewis dot structures (electron dot structures) for the formation of the compounds below.

a. Draw the Lewis dot structure of the atoms.

b. Write the electron configuration of each atom.

c. Show the transfer of electrons using arrow(s)

d. Draw the Lewis dot structures of the ions formed showing their ratio in the compound. Compounds are neutral;

therefore, their electrical charges should balance and cancel each other out.

e. Write the electron configuration of each ion.

f. Write the formula of the compound formed.

1. Lithium and bromine

2. Calcium and chlorine

Properties of Ionic Compounds

Ionic compounds are not composed of individual molecules. Instead, each ion is attracted to the oppositely charged

ions around it. A huge number of ions attract each other in a repeating orderly 3-dimensional pattern. This giant

extended structure of ions is called a crystal lattice. The electrostatic (Coulombic) attraction between the ions is

very strong. The properties of ionic compounds are a result of the strong electrostatic attractions between ions in ionic

bonding.

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1s2 2s1 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

1s2 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

Formula of compound: LiBr

1s2 2s2 2p6 3s2 3p6 4s2

1s2 2s2 2p6 3s2 3p5

1s2 2s2 2p6 3s2 3p6

1s2 2s2 2p6 3s2 3p6

Formula of compound: CaCl2

• Ionic compounds are all crystalline solids at room temperature.

• Ionic compounds are hard, brittle and rigid.

• The smallest particle of an ionic compound is a formula unit. Examples: NaCl, CaBr2.

• Ionic compounds have high melting points. Usually higher than 500 oC. This is because it takes a large amount or

energy to overcome the strong electrostatic attractions between ions in the crystal lattice.

• Ionic compounds are electrolytes. An electrolyte is a substance which conducts electricity when molten (melted –

a pure liquid) or as an aqueous solution (when dissolved in water). Remember that for a substance to conduct

electricity, it must contain mobile charged particles. The diagram below explains why ionic compounds can only

conduct electricity when aqueous or molten but not as solids, using sodium chloride (NaCl) as an example.

WRITING NAMES AND FORMULAS OF IONIC COMPOUNDS

The IUPAC system is the internationally accepted system for naming and writing formulas of compounds. It stands for

the International Union of Pure and Applied Chemistry.

All ionic compounds are neutral, the total positive and total negative charges must balance (they must be equal).

Ionic compounds are made of a metal ion (the positive ion/cation) and a nonmetal ion (the negative ion/anion).

Ionic compound names and formulas are written with the positive ion first followed by the negative ion.

The metal ion could be replaced by a positive polyatomic ion such as ammonium (NH4+).

The nonmetal ion could be replaced by a negative polyatomic ion such as nitrate (NO3⁻)

Naming Ions

• Monovalent metal ions have the same name as their parent atom.

Examples: Na = sodium atom and Na+ = sodium ion

Mg = magnesium atom and Mg2+ = magnesium ion

Zn = zinc atom and Zn2+ = zinc ion (d block monovalent exception)

• Multivalent metal ions have the same name as their parent atom including a roman numeral to indicate their

charge.

These include all metals in the d block and tin, lead and bismuth (exceptions in the d block are zinc, cadmium

and silver).

Examples: Cu = copper atom and Cu+ = copper(I) ion

Fe = iron atom and Fe2+ = iron(II) ion

• Nonmetal ions change the name ending of their parent atom to “ide”.

Examples: S = sulfur atom and S2⁻ = sulfide ion

N = nitrogen atom and N3⁻= nitride ion

• Polyatomic ions retain the exact name of the ion.

Examples: SO42⁻ = sulfate ion

NO2⁻ = nitrite ion

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Writing Formulas of Ionic Compounds

Binary Ionic Compounds (Composed of Only 2 Elements – A Metal and a Nonmetal)

To write the formula of a binary ionic compound we can use the “crisscross” (drop and swap) method.

1. Write the formula of the positive metal ion with its charge. If the metal is multivalent, the roman numeral

represents the charge of the ion. Write the formula of the negative nonmetal ion with its charge.

2. Balance the charges using the crisscross (drop and swap) method.

3. Reduce the subscripts to the smallest whole number ratio.

4. Confirm that the charges are balanced.

Examples

Practice

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Name of compound Positive ion w/ charge Negative ion w/ charge Formula of compound

Barium iodide Ba2+ I⁻ BaI2

Calcium oxide Ca2+ O2⁻ Ca2O2 → CaO

Copper (II) chloride Cu2+ Cl⁻ CuCl2

Iron (III) fluoride Fe3+ F⁻ FeF3

Lead (IV) sulfide Pb4+ S2⁻ Pb2S4 → PbS2

Ternary Ionic Compounds (Including a Polyatomic Ion)

To write the formula of a ternary ionic compound we can also use the “crisscross” (drop and swap) method.

1. Write the formula of the positive ion with its charge. This could be a positive polyatomic ion or a positive

metal ion. If the metal is multivalent, the roman numeral represents the charge of the ion. Write the formula of

the negative ion with its charge. This could be a nonmetal or a polyatomic ion.

2. Balance the charges using the crisscross (drop and swap) method.

3. Add parentheses around the polyatomic ion to show that the subscript applies to the entire ion.

4. Reduce the subscripts to the smallest whole number ratio if necessary.

5. Confirm that the charges are balanced.

Example:

Practice

Name of compound Positive ion w/ charge Negative ion w/ charge Formula of compound

Lithium cyanide Li+ CN⁻ LiCN

Magnesium sulfate Mg2+ SO42⁻ Mg2(SO4)2 → MgSO4

Aluminum carbonate Al3+ CO32⁻ Al2(CO3)3

Sodium nitrate Na+ NO3⁻ NaNO3

Barium hydroxide Ba2+ OH⁻ Ba(OH)2

Ammonium phosphate NH4 + PO43⁻ (NH4)3PO4

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Naming Ionic Compounds

To name ionic compounds:

1. Combine the names of the cation and the anion in that order.

2. Make sure to change the ending of nonmetal ions to “ide”.

3. If the formula includes a multivalent metal cation, use the charge of the anion (you will always know the

charge of the anion) to find the charge of the cation. Add the charge as a roman numeral after the name of the

metal ion.

4. If the positive or negative ion is a polyatomic ion, write the name of the polyatomic ion.

Practice: Name the following ionic compounds

1. LiCl lithium chloride

2. SnF2 tin(II) fluoride

3. PbS2 lead(IV) sulfide

4. NaNO3 sodium nitrate

5. ZnCO3 zinc carbonate

6. NH4Br ammonium bromide

7. Ag2SO3 silver sulfite

8. Fe3(PO4)2 iron(II) phosphate

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Example with multivalent metal cation:

Name MnO2

Manganese(IV) oxide

Example:

Name CaCl2

Calcium chloride

The oxide ion is the anion. Charge of O (oxide) = 2-

There are 2 oxide ions in MnO2

Total negative charge = 2(2-) = 4-

To balance this and make the compound neutral:

Total positive charge must equal 4+

Charge of compound must = 0 (4-) + (4+) = 0

There is 1 manganese ion in the formula MnO2

Mn charge = 4+/1 = 4+ Manganese(IV) oxide

NAMING AND WRITING FORMULAS OF ACIDS

Acids can be thought of as a hydrogen ion (H+) followed by an anion (negative ion). The anion can be a nonmetal or

a polyatomic anion. Examples of acids are HCl, HNO3, H2SO4

This is not strictly correct, as acids are NOT ionic compounds, they are molecular compounds which when dissolved

in water (in the aqueous state) can form ions.

Example:

HCl is a molecular compound since it is a made of 2 nonmetals. It is a gas. HCl(g)

When HCl is dissolved in water, H+ and Cl⁻ ions are formed.

HCl(g) → HCl(aq) → H+(aq) + Cl⁻(aq)

We will learn more about properties of acids. For now, we are only learning how to name them and write their formulas.

If the anion is a nonmetal ion such as in HCl, the acid is a binary acid.

If the anion is a polyatomic ion containing oxygen, such as in H2SO4, the acid is a ternary acid (an oxyacid).

Naming Acids

There are three rules for naming acids, based on the anion name.

The cation is always H+

1) Binary Acids

H+ and a nonmetal anion ending in –ide: Acid name is “hydro_____ic acid”

Take the root from the anion name and fill in the blank.

Helpful acronym: "My rIDE has HYDRaulICs"

Examples: HCl Cl⁻ is the anion, its name is chloride

Name of acid is: hydrochloric acid

H2S S2⁻ is the anion, its name is sulfide

Name of acid is: hydrosulfuric acid

Note: when the anion contains sulfur, the root is sulfur, not sulf. H2S is hydrosulfuric acid, NOT hydrosulfic acid.

2) Ternary Acids

H+ and a polyatomic anion ending in –ate: Acid name is “______ic acid.

Take the root from the anion name and fill in the blank.

Helpful acronyms: “I ATE organIC” or “What I ATE was ICky”

Examples: HNO3 NO3⁻ is the anion, its name is nitrate

Name of acid is: nitric acid

H2CO3 CO32⁻ is the anion, its name is carbonate

Name of acid is: carbonic acid

Note: when the anion contains sulfur or phosphorus, the roots are sulfur- and phosphor-, respectively, NOT

sulf- and phosph-. H2SO4 is sulfuric acid, NOT sulfic acid and H3PO4 is phosphoric acid, NOT phosphic acid.

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water

3) Ternary Acids

H+ and a polyatomic anion ending in –ite: Acid name is “________ous acid.”

Take the root from the anion name and fill in the blank.

"Every bITE was deliciOUS" or “The snake bITE was poisonous”

Examples: HNO2 NO2⁻ is the anion, its name is nitrite

Name of acid is: nitrous acid

H2SO3 SO32 ⁻ is the anion, its name is sulfite

Name of acid is: sulfurous acid

Note: when the anion contains sulfur or phosphorus, the roots are sulfur- and phosphor-, respectively, NOT sulf-

and phosph-. H2SO3 is sulfurous acid, NOT sulfous acid and H3PO3 is phosphorous acid, NOT phosphous acid.

4) Where does HCN fit?

CN⁻ is a polyatomic ion. Its name is cyanIDE

Cyanide does not end it ATE or ITE!

We use the rule for binary acids because cyanide ends in IDE.

H+ and a nonmetal anion ending in –ide: Acid name is “hydro_____ic acid”

HCN is hydrocyanic acid!

Writing Formulas of Acids

To write the formula for an acid, given the name, you would just work backwards, using the 3 rules for naming

acids. Acid formulas must be crisscrossed as with ionic compounds, so that the charges balance and equal out

to zero.

The cation is always H+

Examples:

Hydrobromic acid anion is bromide H+ and Br⁻ formula: HBr

Acetic acid anion is acetate H+ and C2H3O2⁻ formula: HC2H3O2

Phosphorous acid anion is phosphite H+ and PO33⁻ formula: H3PO3

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