Chapter 3: Ionic Compounds

Concepts to Review

The Periodic Table: Referenced Sections $2.4$ and $2.5$ .
Electron Configuration: Referenced Sections $2.7$ and $2.8$ .

3.1 Ions

Role of Ions: Ions are crucial in many cellular processes, such as signal transmission between nerve cells.
Ion Formation: An ion is formed when a neutral atom either gains or loses electrons.
- Cation: A positively charged ion formed when a neutral atom loses one or more electrons.
- Anion: A negatively charged ion formed when a neutral atom gains one or more electrons.
Properties of Ionic Compounds (e.g., Alkali Metals with Halogens):
- High melting points.
- Stable, white, crystalline solids.
- Soluble in water.
- Conduct electricity when dissolved in water.
Sodium Chloride Formation: Table salt (sodium chloride) is formed through the reaction of sodium (Na) with chlorine (Cl).
- Aqueous solutions of sodium chloride conduct electricity, demonstrating the presence of mobile charged particles (ions).
Alkali Metal Ion Formation (Group 1A):
- Alkali metals have a single valence electron ( $ns^1$ electron configuration).
- They lose this single electron to form a positively charged cation, achieving a stable noble gas electron configuration.
- Example: Sodium (Na) loses one electron to form $Na^+$ ( $Na ightarrow Na^+ + e^-$ ).
Halogen Ion Formation (Group 7A):
- Halogen atoms have an $ns^2np^5$ electron configuration (seven valence electrons).
- They gain one electron to form a negatively charged anion, achieving a stable noble gas electron configuration.
- Example: Chlorine (Cl) gains one electron to form $Cl^-$ ( $Cl + e^- ightarrow Cl^-$ ).
Ion Symbolism:
- Cation: Written by adding the positive charge as a superscript to the element symbol (e.g., $Na^+$ ).
- Anion: Written by adding the negative charge as a superscript (e.g., $Cl^-$ ).
- If the charge is greater than $1$ , the number is included (e.g., $Ca^{2+}$ , $N^{3-}$ ).

3.2 Ions and the Octet Rule

Octet Rule: Main-group elements tend to react in ways that leave them with eight valence electrons.
Electron Transfer in Ionic Reactions: When an alkali metal (e.g., sodium) reacts with a halogen (e.g., chlorine), the metal atom transfers an electron from its valence shell to the valence shell of the halogen.
Cation Formation (Main-Group Metals):
- Main-group metals lose electrons to form cations.
- They attain an electron configuration identical to the noble gas located before them in the periodic table.
Anion Formation (Main-Group Nonmetals):
- Main-group nonmetals gain electrons to form anions.
- They attain an electron configuration identical to the noble gas located after them in the periodic table.
Worked Example 3.1: Magnesium Ion Formation (Z=12):
- Electron Configuration of Neutral Mg: $1s^22s^22p^63s^2$ .
- Valence Electrons: Two electrons in the $3s^2$ subshell.
- Octet Achievement: Magnesium can achieve a valence-shell octet by losing the two electrons in the $3s^2$ subshell. The second shell ( $2s^22p^6$ ) already contains an octet.
- Ion Configuration: $Mg^{2+}$ : $1s^22s^22p^6$ (neon configuration or $[Ne]$ ).
- Reason for Charge: A neutral magnesium atom has $12$ protons and $12$ electrons. Losing $2$ electrons results in $12$ protons and $10$ electrons, leading to an excess of $2$ protons and thus a $+2$ charge.
- Ion Symbol: $Mg^{2+}$ .
Worked Example 3.2: Nitrogen Ion Formation (Z=7):
- Electron Configuration of Neutral N: $1s^22s^22p^3$ .
- Valence Electrons: Five electrons in the second shell ( $2s^22p^3$ ).
- Octet Achievement: Nitrogen needs to gain $3$ more electrons to reach an octet (matching neon's configuration).
- Ion Configuration: $N^{3-}$ : $1s^22s^22p^6$ (neon configuration).
- Ion Symbol: $N^{3-}$ .

3.3 Ions of Some Common Elements

Predicting Charges (Summary based on Octet Rule):
- Group 1A (Alkali Metals): Lose $1e^-$ to form $M^+$ ions ( $M ightarrow M^+ + e^-$ ).
- Group 2A (Alkaline Earth Metals): Lose $2e^-$ to form $M^{2+}$ ions ( $M ightarrow M^{2+} + 2e^-$ ).
- Group 3A: Aluminum (Al) typically forms $Al^{3+}$ by losing $3e^-$ ; other common ions are rare for this group.
- Group 4A, 5A (Nonmetals): Generally do not form common ions by gaining/losing electrons, as it requires gaining/losing too many electrons.
- Group 6A (Chalcogens): Gain $2e^-$ to form $X^{2-}$ ions (e.g., $O^{2-}$ , $S^{2-}$ ) ( $X + 2e^- ightarrow X^{2-}$ ).
- Group 7A (Halogens): Gain $1e^-$ to form $X^-$ ions ( $X + e^- ightarrow X^-$ ).
- Group 8A (Noble Gases): Unreactive; neither lose nor gain electrons.
Transition Metals: Form cations but do not strictly follow the octet rule. They can lose one or more $d$ electrons in addition to valence $s$ electrons, leading to multiple possible charges (e.g., $Fe^{2+}$ and $Fe^{3+}$ ).
Ionic Charges Prediction for Main-Group Elements:
- Cation charge for Group 1A and 2A metals: Equal to the group number ( $+1$ for Group 1A, $+2$ for Group 2A).
- Anion charge for nonmetals in Groups 5A, 6A, and 7A: Calculated as $8 - ( ext{group number})$ (e.g., Group 7A: $8-7 = -1$ ; Group 6A: $8-6 = -2$ ; Group 5A: $8-5 = -3$ ).
Worked Example 3.3: Likely Ion Formation:
- (a) $S^{3-}$ : Sulfur (S) is in Group 6A, has $6$ valence electrons, and needs to gain $2$ to reach an octet ( $S^{2-}$ ). Gaining $3$ electrons does not result in a noble gas configuration, so $S^{3-}$ is unlikely to form.
- (b) $Si^{2+}$ : Silicon (Si) is a nonmetal in Group 4A. It needs to gain or lose $4$ electrons to achieve a noble gas configuration, which is energetically unfavorable for ion formation. Thus, $Si^{2+}$ is unlikely to form.
- (c) $Sr^{2+}$ : Strontium (Sr) is a metal in Group 2A. It has $2$ outer-shell electrons and can easily lose both to achieve a noble gas configuration ( $Sr^{2+}$ ). This ion forms readily.

3.4 Periodic Properties and Ion Formation

General Trend: Metals on the left side of the periodic table tend to lose electrons, while nonmetals on the right side tend to gain electrons.
Ionization Energy: The energy required to remove one electron from a single atom in the gaseous state.
- Small values indicate that elements easily lose electrons to form cations.
- Decreases down a group and increases across a period.
Electron Affinity: The energy released when an electron is added to a single atom in the gaseous state.
- Large (more negative) values indicate that elements gain electrons most easily.
- Halogens have the largest electron affinities among main group elements.
Predicting Ion Formation Based on Properties:
- Alkali Metals: Have very low ionization energies, making it easy to lose electrons.
- Halogens: Have very high electron affinities, making it easy to gain electrons.
- Noble Gases: Have very high ionization energies and very low (or positive) electron affinities, so they rarely lose or gain electrons.
- Elements in the middle of the periodic table: Do not easily form ions because they require significant energy input to either gain or lose multiple electrons.
Worked Example 3.4: Rubidium's Ionization Energy:
- Rubidium (Rb) is an alkali metal (Group 1A), located below potassium (K) in the periodic table.
- Alkali metals (Li, Na, K) generally have low ionization energies.
- Therefore, rubidium's ionization energy is predicted to be low, similar to other alkali metals.
Worked Example 3.5: Ease of Electron Loss (Mg vs. S):
- Magnesium (Mg): A Group 2A element on the left side of the periodic table. It has a relatively low ionization energy and readily loses electrons.
- Sulfur (S): A Group 6A element on the right side of the periodic table. It has a higher ionization energy than magnesium and tends to gain electrons rather than lose them. Therefore, Mg is more likely to lose an electron than S.

3.5 Naming Monoatomic Ions

Main-Group Metal Cations (Type I):
- Named by simply identifying the metal, followed by the word "ion."
- Examples:
  - $K^+$ : Potassium ion
  - $Mg^{2+}$ : Magnesium ion
  - $Al^{3+}$ : Aluminum ion
Transition Metal Cations (Type II):
- Can form more than one type of cation (i.e., multiple charges).
- Two Naming Systems:
  - Old System: Uses suffixes to indicate charge.
    - -ous: For the ion with the smaller charge.
    - -ic: For the ion with the larger charge.
    - Examples:
      - $Cr^{2+}$ : Chromous ion
      - $Cr^{3+}$ : Chromic ion
  - New (Stock) System: Uses Roman numerals in parentheses immediately after the metal name.
    - Examples:
      - $Cr^{2+}$ : Chromium(II) ion
      - $Cr^{3+}$ : Chromium(III) ion
- Table 3.1: Names of Some Transition Metal Cations (summarized):
  - Chromium:
    - $Cr^{2+}$ : Chromous / Chromium(II)
    - $Cr^{3+}$ : Chromic / Chromium(III)
  - Copper:
    - $Cu^+$ : Cuprous / Copper(I)
    - $Cu^{2+}$ : Cupric / Copper(II)
  - Iron:
    - $Fe^{2+}$ : Ferrous / Iron(II)
    - $Fe^{3+}$ : Ferric / Iron(III)
  - Mercury:
    - $[Hg_2]^{2+}$ : Mercurous / Mercury(I) (Note: Dimer, average charge $+1$ per Hg)
    - $Hg^{2+}$ : Mercuric / Mercury(II)
  - Tin:
    - $Sn^{2+}$ : Stannous / Tin(II)
    - $Sn^{4+}$ : Stannic / Tin(IV)
Anions:
- Named by replacing the ending of the element name with "-ide," followed by the word "ion."
- Examples:
  - $Cl^-$ : Chloride ion (from chlorine)
  - $O^{2-}$ : Oxide ion (from oxygen)
  - $N^{3-}$ : Nitride ion (from nitrogen)

3.6 Polyatomic Ions

Definition: An ion composed of more than one atom.
Bonding: The atoms within a polyatomic ion are held together by covalent bonds.
Charge: A polyatomic ion carries a net charge because the total number of electrons is different from the total number of protons in its combined atoms.
Importance: These ions are frequently encountered and their names and formulas must be memorized.
Table 3.3: Some Common Polyatomic Ions (summarized):
- $H_3O^+$ : Hydronium ion
- $NH_4^+$ : Ammonium ion
- $CH3CO2^-$ : Acetate ion
- $CO_3^{2-}$ : Carbonate ion
- $HCO_3^-$ : Hydrogen carbonate ion (bicarbonate ion)
- $CrO_4^{2-}$ : Chromate ion
- $Cr2O7^{2-}$ : Dichromate ion
- $CN^-$ : Cyanide ion
- $OH^-$ : Hydroxide ion
- $ClO^-$ : Hypochlorite ion
- $NO_3^-$ : Nitrate ion
- $NO_2^-$ : Nitrite ion
- $C2O4^{2-}$ : Oxalate ion
- $MnO_4^-$ : Permanganate ion
- $PO_4^{3-}$ : Phosphate ion
- $HPO_4^{2-}$ : Hydrogen phosphate ion (biphosphate ion)
- $H2PO4^-$ : Dihydrogen phosphate ion
- $SO_4^{2-}$ : Sulfate ion
- $HSO_4^-$ : Hydrogen sulfate ion (bisulfate ion)
- $SO_3^{2-}$ : Sulfite ion
Table 3.2: Some Biologically Important Ions (summarized):
- $Ca^{2+}$ : Outside cell (bones/teeth); bone/tooth structure, blood clotting, muscle contraction, nerve impulses. Dietary Source: Milk, whole grains, leafy vegetables.
- $Fe^{2+}$ : Blood hemoglobin; transports oxygen. Dietary Source: Liver, red meat, leafy green vegetables.
- $K^+$ : Inside cells; maintains ion concentrations, regulates insulin release, heartbeat. Dietary Source: Milk, oranges, bananas, meat.
- $Na^+$ : Outside cells; protects against fluid loss, muscle contraction, nerve impulses. Dietary Source: Table salt, seafood.
- $Mg^{2+}$ : Inside cells (bone); in enzymes, needed for energy/muscle contraction. Dietary Source: Leafy green plants, seafood, nuts.
- $Cl^-$ : Outside cells (gastric juice); maintains fluid balance, $CO_2$ transfer. Dietary Source: Table salt, seafood.
- $HCO_3^-$ : Outside cells; controls acid-base balance in blood. Dietary Source: By-product of food metabolism.
- $HPO_4^{2-}$ : Inside cells (bones/teeth); controls acid-base balance in cells. Dietary Source: Fish, poultry, milk.

3.7 Ionic Bonds

Nature of the Bond: Ionic bonds form due to the strong electrostatic attraction between oppositely charged ions.
- This attraction holds positive ions (cations) and negative ions (anions) together.
Ionic Solids (Crystals):
- In an ionic solid, many ions are attracted by ionic bonds to their nearest neighbors, forming a crystal lattice structure.
- Example: The arrangement of $Na^+$ and $Cl^-$ ions in a sodium chloride crystal.
Ionic Compounds: Compounds formed by ionic bonds are referred to as ionic compounds.
Formation: Ion-transfer reactions between metals and nonmetals result in products with properties distinct from the elements themselves.

3.8 Formulas of Ionic Compounds

Neutrality Principle: All chemical compounds are electrically neutral. The total positive charge from cations must balance the total negative charge from anions.
Determining Formulas: Once the ions and their charges are identified, determine the ratio of each ion needed to achieve a total charge of zero.
Chemical Formula: Represents the lowest whole-number ratio of anions and cations in the compound.
Combinations:
- Ions with Same Magnitude of Charge: One of each ion is needed.
 - Example: $K^+ + F^- ightarrow KF$ ( $(+1) + (-1) = 0$ ).
- Ions with Different Magnitudes of Charge: Unequal numbers of anions and cations combine.
 - The number of one ion is typically equal to the charge of the other ion (cross-multiplication of charges, then simplification).
 - Example: $2 K^+ + O^{2-} ightarrow K_2O$ ( $2(+1) + (-2) = 0$ ).
 - Example: $Ca^{2+} + 2 Cl^- ightarrow CaCl_2$ ( $(+2) + 2(-1) = 0$ ).
Simplest Formula: Ionic compound formulas represent the lowest possible ratio of atoms.
Formula Unit: The smallest neutral unit of an ionic compound.
- For NaCl, the formula unit is one $Na^+$ ion and one $Cl^-$ ion.
- For $CaF_2$ , the formula unit is one $Ca^{2+}$ ion and two $F^-$ ions.
Rules for Writing Formulas:
- List the cation first and the anion second.
- Do not write the charges of the ions in the final formula.
- Use parentheses around a polyatomic ion formula if it has a subscript greater than $1$ .
Worked Example 3.6: Formula for Calcium Nitrate:
- Ions: $Ca^{2+}$ and $NO_3^-$ .
- Balancing Charges: One $Ca^{2+}$ ( $+2$ charge) requires two $NO_3^-$ ions ( $2 imes (-1) = -2$ charge) to achieve neutrality.
- Formula: $Ca(NO3)2$ .

3.9 Naming Ionic Compounds

General Rule: Name the cation first, then the anion, with a space between the words.
Types of Ionic Compounds:
- Type I Ionic Compounds: Contain cations of main-group elements whose charges do not vary.
  - Do not specify the charge on the cation.
  - Examples:
    - NaCl: Sodium chloride
    - $MgCO_3$ : Magnesium carbonate
- Type II Ionic Compounds: Contain metals that can exhibit more than one charge (typically transition metals).
  - Must specify the charge on the cation.
  - Naming Systems:
    - Old System: Uses suffixes "-ous" and "-ic" (e.g., ferrous, ferric).
    - New System: Uses Roman numerals in parentheses after the metal name.
    - Examples:
      - $FeCl_2$ : Iron(II) chloride or ferrous chloride ( $Fe^{2+}$ )
      - $FeCl_3$ : Iron(III) chloride or ferric chloride ( $Fe^{3+}$ )
  - Important Note: Do not use prefixes like "di-" or "tri-" to indicate the number of anions (e.g., not "iron dichloride"). The charge on the metal implicitly dictates the number of anions needed for a neutral compound.
Worked Example 3.8: Naming Ionic Compounds:
- (a) KF: Potassium fluoride (K is Group 1A, fixed charge).
- (b) $MgCl_2$ : Magnesium chloride (Mg is Group 2A, fixed charge).
- (c) $AuCl_3$ : Gold(III) chloride (Each Cl has charge -1. Three Cl means total -3. Gold must be $Au^{3+}$ ).
- (d) $Fe2O3$ : Iron(III) oxide (Each O has charge -2. Three O means total -6. Two Fe must have total +6, so each Fe is $Fe^{3+}$ ).
Table 3.4: Some Common Ionic Compounds and Their Applications (summarized):
- $(NH4)2CO_3$ : Ammonium carbonate (Smelling salts)
- $Ca(OH)_2$ : Calcium hydroxide (Mortar, plaster, whitewash)
- CaO: Calcium oxide (Lawn treatment, industrial chemical)
- $Li2CO3$ : Lithium carbonate (Treatment of bipolar disorder)
- $Mg(OH)_2$ : Magnesium hydroxide (Antacid)
- $MgSO_4$ : Magnesium sulfate (Laxative, anticonvulsant)
- $KMnO_4$ : Potassium permanganate (Antiseptic, disinfectant)
- $KNO_3$ : Potassium nitrate (Fireworks, matches, desensitizer for teeth)
- $AgNO_3$ : Silver nitrate (Antiseptic, germicide)
- $NaHCO_3$ : Sodium bicarbonate (Baking powder, antacid, mouthwash, deodorizer)
- NaClO: Sodium hypochlorite (Disinfectant, active ingredient in household bleach)
- ZnO: Zinc oxide (Skin protection, in calamine lotion)
Worked Example 3.7: Writing Formulas (Sodium & Calcium Compounds):
- Cations: Sodium ( $Na^+$ , Group 1A) and Calcium ( $Ca^{2+}$ , Group 2A).
- (a) Bromide ( $Br^-$ ):
 - Sodium bromide: NaBr ( $Na^+$ and $Br^-$ )
 - Calcium bromide: $CaBr_2$ ( $Ca^{2+}$ and two $Br^-$ )
- (b) Sulfide ( $S^{2-}$ ):
 - Sodium sulfide: $Na_2S$ (two $Na^+$ and $S^{2-}$ ; from Group 6A)
 - Calcium sulfide: CaS ( $Ca^{2+}$ and $S^{2-}$ ; from Group 6A)
- (c) Sulfate ( $SO_4^{2-}$ ):
 - Sodium sulfate: $Na2SO4$ (two $Na^+$ and $SO_4^{2-}$ )
 - Calcium sulfate: $CaSO4$ ( $Ca^{2+}$ and $SO4^{2-}$ )
- (d) Phosphate ( $PO_4^{3-}$ ):
 - Sodium phosphate: $Na3PO4$ (three $Na^+$ and $PO_4^{3-}$ )
 - Calcium phosphate: $Ca3(PO4)2$ (three $Ca^{2+}$ and two $PO4^{3-}$ ; $3(+2) + 2(-3) = 0$ )

3.10 Some Properties of Ionic Compounds

Crystal Structure: Ions in an ionic compound arrange themselves into a regular, repeating pattern that efficiently fills space and maximizes ionic bonding, forming a crystal lattice.
Rigidity: Ions in an ionic solid are held rigidly in place by strong attractions to their neighbors, giving them structural integrity.
Electrical Conductivity:
- Ionic solids are generally poor conductors of electricity because their ions are fixed in the lattice.
- When dissolved in water (or molten), the ions become free to move, allowing the solution (or melt) to conduct electricity.
Melting and Boiling Points: Ionic compounds typically have very high melting and boiling points due to the strong electrostatic forces that need to be overcome.
- Example: Sodium chloride (NaCl) melts at $801^ ext{o}C$ and boils at $1413^ ext{o}C$ .
Brittleness: Ionic solids are typically brittle and shatter if struck sharply. This is because a sharp blow can cause layers of ions to shift, bringing like-charged ions into repulsion, which causes the crystal to cleave.
Solubility in Water: Ionic compounds dissolve in water if the attractive forces between water molecules and the ions are strong enough to overcome the attractive forces between the ions themselves in the crystal lattice. Not all ionic compounds are water-soluble.

3.11 H $^+$ and OH $^-$ Ions: An Introduction to Acids and Bases

Significance: Hydrogen cation ( $H^+$ ) and hydroxide anion ( $OH^-$ ) are two of the most important ions, fundamental to the concepts of acids and bases.
Acids: A substance that provides $H^+$ ions in water.
- Hydrogen Cation ( $H^+$ ): Simply a proton (a hydrogen atom that lost its single electron).
- Hydronium Ion ( $H3O^+$ ): When an acid dissolves in water, the proton ( $H^+$ ) attaches to a water molecule ( $H2O$ ) to form a hydronium ion ( $H^+ + H2O ightarrow H3O^+$ ).
- Chemists often use $H^+$ and $H_3O^+$ interchangeably to represent the acidic species in aqueous solution.
- Varying H $^+$ Production: Different acids can provide different numbers of $H^+$ ions per acid molecule:
 - Monoprotic: Hydrochloric acid (HCl) provides one $H^+$ ion.
 - Diprotic: Sulfuric acid ( $H2SO4$ ) can provide two $H^+$ ions.
 - Triprotic: Phosphoric acid ( $H3PO4$ ) can provide three $H^+$ ions.
Bases: A substance that provides $OH^-$ ions in water.
- Hydroxide Anion ( $OH^-$ ): A polyatomic ion where an oxygen atom is covalently bonded to a hydrogen atom, carrying a negative charge.
- Examples: Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are common bases.
- When bases like NaOH or KOH dissolve, $OH^-$ anions enter the solution along with the metal cation.
- Varying OH $^-$ Production: Different bases can provide different numbers of $OH^-$ ions per formula unit:
  - Monobasic: Sodium hydroxide (NaOH) provides one $OH^-$ ion.
  - Dibasic: Barium hydroxide ( $Ba(OH)_2$ ) can provide two $OH^-$ ions.
Table 3.5: Some Common Acids and the Anions Derived from Them (summarized):
- Acetic acid ( $CH3COOH$ ): Acetate ion ( $CH3COO^-$ )
- Carbonic acid ( $H2CO3$ ): Hydrogen carbonate ion ( $HCO3^-$ ) / Carbonate ion ( $CO3^{2-}$ )
- Hydrochloric acid (HCl): Chloride ion ( $Cl^-$ )
- Nitric acid ( $HNO3$ ): Nitrate ion ( $NO3^-$ )
- Nitrous acid ( $HNO2$ ): Nitrite ion ( $NO2^-$ )
- Phosphoric acid ( $H3PO4$ ): Dihydrogen phosphate ion ( $H2PO4^-$ ) / Hydrogen phosphate ion ( $HPO4^{2-}$ ) / Phosphate ion ( $PO4^{3-}$ )
- Sulfuric acid ( $H2SO4$ ): Hydrogen sulfate ion ( $HSO4^-$ ) / Sulfate ion ( $SO4^{2-}$ )

Concept Map: Intramolecular Forces

Ionic Bonds (Chapter 3):
- Formation involves the transfer of electrons.
- Formation of ions depends on:
  - Ionization energy: Energy required to remove an electron. Small values facilitate cation formation.
  - Electron affinity: Energy released when an electron is added. Large values facilitate anion formation.
  - Electron configurations: Atoms achieve stable noble gas configurations.
  - Octet rule: Tendency of main-group elements to achieve eight valence electrons.
- Ionic Compounds:
  - Formed by the balancing of charges between cations and anions.
  - Properties:
    - Solubility in polar solvents (e.g., water).
    - High melting points.
    - Solution conductivity (when dissolved or molten).
Covalent Bonds (Chapter 4 - Future Topic):
- Formation involves the sharing of electrons.

Chapter 3: Ionic Compounds

Chapter 3: Ionic Compounds

Concepts to Review

3.1 Ions

3.2 Ions and the Octet Rule

3.3 Ions of Some Common Elements

3.4 Periodic Properties and Ion Formation

3.5 Naming Monoatomic Ions

3.6 Polyatomic Ions

3.7 Ionic Bonds

3.8 Formulas of Ionic Compounds

3.9 Naming Ionic Compounds

3.10 Some Properties of Ionic Compounds

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

Concept Map: Intramolecular Forces

3.11 H $^+$ and OH $^-$ Ions: An Introduction to Acids and Bases