Chapter 3: Ionic Compounds – Structured Study Notes for GOB Chemistry

Chapter 3: Ionic Compounds

3.1 Ions

Ions are fundamental to many biological and chemical processes, including nerve signal transmission. An ion is formed when a neutral atom gains or loses electrons, resulting in a charged particle.

• Cation: A positively charged ion formed by the loss of electrons.

• Anion: A negatively charged ion formed by the gain of electrons.

• Example: Sodium (Na) loses one electron to become Na+; chlorine (Cl) gains one electron to become Cl−.

Electricity flows through solutions containing free-moving ions, as demonstrated by sodium chloride solution conducting electricity.

The charge of an ion is indicated as a superscript (e.g., Na+, Cl−, Ca2+, N3−).

3.2 Ions and the Octet Rule

The octet rule states that main group elements tend to react to achieve eight valence electrons, resembling noble gas configurations.

• Metals lose electrons to form cations with noble gas configurations before them.

• Nonmetals gain electrons to form anions with noble gas configurations after them.

Example: Magnesium (Mg, Z=12) loses two electrons to form Mg2+ with a neon configuration. Nitrogen (N, Z=7) gains three electrons to form N3− with a neon configuration.

• Electron configuration for Mg: $1s^2\ 2s^2\ 2p^6\ 3s^2$

• Electron configuration for Mg2+: $1s^2\ 2s^2\ 2p^6$

• Electron configuration for N: $1s^2\ 2s^2\ 2p^3$

• Electron configuration for N3−: $1s^2\ 2s^2\ 2p^6$

3.3 Ions of Some Common Elements

The periodic table can be used to predict the charges of ions formed by main group elements.

• Group 1A: +1 charge

• Group 2A: +2 charge

• Group 3A: +3 charge (Al only)

• Group 6A: −2 charge (O, S)

• Group 7A: −1 charge (halogens)

• Transition metals: Variable charges, do not always follow the octet rule.

Charges for main group elements can be predicted:

• Cation charge = group number (for 1A, 2A metals)

• Anion charge = 8 − group number (for 5A, 6A, 7A nonmetals)

3.4 Periodic Properties and Ion Formation

Periodic trends influence ion formation. Metals tend to lose electrons (form cations), while nonmetals tend to gain electrons (form anions).

• Ionization energy: Energy required to remove an electron; lower for metals.

• Electron affinity: Energy released when an electron is added; higher for halogens.

Halogens gain electrons most easily, alkali metals lose electrons most easily, and noble gases are generally unreactive.

3.5 Naming Monoatomic Ions

Monoatomic ions are named based on their element and charge.

• Main group cations: Element name + "ion" (e.g., potassium ion, magnesium ion).

• Transition metal cations: Old system (-ous, -ic) or new system (Roman numerals).

• Anions: Element name ending replaced with "-ide" (e.g., chloride ion, oxide ion).

Element	Symbol	Old Name	New Name
Chromium	Cr2+	Chromous	Chromium(II) ion
Chromium	Cr3+	Chromic	Chromium(III) ion
Copper	Cu+	Cuprous	Copper(I) ion
Copper	Cu2+	Cupric	Copper(II) ion
Iron	Fe2+	Ferrous	Iron(II) ion
Iron	Fe3+	Ferric	Iron(III) ion
Mercury	Hg22+	Mercurous	Mercury(I) ion
Mercury	Hg2+	Mercuric	Mercury(II) ion
Tin	Sn2+	Stannous	Tin(II) ion
Tin	Sn4+	Stannic	Tin(IV) ion

Element	Symbol	Name
Bromine	Br−	Bromide ion
Chlorine	Cl−	Chloride ion
Fluorine	F−	Fluoride ion
Iodine	I−	Iodide ion
Oxygen	O2−	Oxide ion
Sulfur	S2−	Sulfide ion

3.6 Polyatomic Ions

Polyatomic ions are ions composed of more than one atom, held together by covalent bonds. Their names and formulas must be memorized.

Name	Formula
Hydronium ion	H3O+
Ammonium ion	NH4+
Acetate ion	CH3COO−
Carbonate ion	CO32−
Nitrate ion	NO3−
Sulfate ion	SO42−
Phosphate ion	PO43−
Hydroxide ion	OH−

Polyatomic ions are important in biological and chemical systems, such as bicarbonate (HCO3−) for acid–base balance.

3.7 Ionic Bonds

Ionic bonds are formed by the electrostatic attraction between oppositely charged ions. In sodium chloride, Na+ and Cl− ions arrange in a crystal lattice.

Ionic compounds are held together by these bonds, forming ionic solids.

3.8 Formulas of Ionic Compounds

Ionic compounds are electrically neutral. The formula shows the ratio of cations to anions needed to balance charges.

• If ions have the same charge, one of each is needed (e.g., KF).

• If ions have different charges, unequal numbers are needed (e.g., CaCl2).

The formula unit is the smallest neutral unit (e.g., NaCl, CaF2).

• Cation is listed first, anion second.

• Charges are not shown in the formula.

• Parentheses are used for polyatomic ions with subscripts (e.g., Ca(NO3)2).

3.9 Naming Ionic Compounds

Ionic compounds are named by citing the cation first, then the anion.

• Type I: Main group cations; charge does not vary (e.g., sodium chloride).

• Type II: Transition metals; charge specified with Roman numerals (e.g., iron(III) chloride).

Chemical Name	Formula	Applications
Ammonium carbonate	(NH4)2CO3	Smelling salts
Calcium hydroxide	Ca(OH)2	Mortar, plaster, whitewash
Magnesium hydroxide	Mg(OH)2	Antacid
Potassium permanganate	KMnO4	Antiseptic, disinfectant
Sodium bicarbonate	NaHCO3	Baking powder, antacid

Example: FeCl3 is iron(III) chloride; MgCl2 is magnesium chloride.

3.10 Properties of Ionic Compounds

Ionic compounds exhibit distinct physical properties due to their ionic bonds.

• High melting and boiling points (e.g., NaCl melts at 801°C).

• Rigid crystal structure; solids shatter if struck.

• Conduct electricity when dissolved in water.

• Solubility depends on the attraction between water and ions.

3.11 H+ and OH− Ions: Introduction to Acids and Bases

Hydrogen cation (H+) and hydroxide anion (OH−) are central to acid–base chemistry.

• Acid: Provides H+ ions in water; H+ attaches to water to form hydronium (H3O+).

• Base: Provides OH− ions in water.

• Different acids and bases provide varying numbers of H+ or OH− ions per molecule.

Acid	Anion Derived
Acetic acid (CH3COOH)	Acetate (CH3COO−)
Carbonic acid (H2CO3)	Bicarbonate (HCO3−), Carbonate (CO32−)
Hydrochloric acid (HCl)	Chloride (Cl−)
Nitric acid (HNO3)	Nitrate (NO3−)
Phosphoric acid (H3PO4)	Dihydrogen phosphate (H2PO4−), Hydrogen phosphate (HPO42−), Phosphate (PO43−)
Sulfuric acid (H2SO4)	Hydrogen sulfate (HSO4−), Sulfate (SO42−)

Concept Map: Ionic vs. Covalent Bonds

This concept map summarizes the differences between ionic and covalent bonds, and the factors influencing ion formation and properties of ionic compounds.

Additional info: These notes cover all major learning objectives for Chapter 3: Ionic Compounds, including ion formation, periodic trends, naming conventions, polyatomic ions, ionic bonding, formulas, properties, and an introduction to acids and bases. All images included are directly relevant to the explanations provided.