Inorganic

INORGANIC CHEMISTRY REVIEWER

Hydrogen Bonding • Hydrogen-Bond Theories • Metallic Bond

1. Hydrogen Bonding

Easy meaning: A hydrogen bond is a weak attraction between a slightly positive hydrogen atom and a nearby electronegative atom, usually oxygen (O), nitrogen (N), or fluorine (F). It is often shown with a dotted line (···).

Think of it as a small “pull” between molecules. It is weaker than a normal covalent bond, but it can strongly affect a substance’s properties.

How does a hydrogen bond form?

• Hydrogen is covalently bonded to a very electronegative atom such as O, N, or F.

• The bond makes hydrogen slightly positive (δ+), while the electronegative atom is slightly negative (δ−).

• The δ+ hydrogen is attracted to a lone pair or negative region on a nearby electronegative atom.

General pattern: X–H···Y, where X and Y are usually O, N, or F. The dotted line represents the hydrogen bond.

Examples

• Water (H₂O): hydrogen bonds between neighboring water molecules.

• Hydrogen fluoride (HF): hydrogen bonding can occur between HF molecules.

• Ammonia (NH₃): hydrogen bonding can occur between ammonia molecules, though it is generally weaker than in water or HF.

Why does hydrogen bonding happen?

• Hydrogen is very small, so it can get close to another atom.

• When hydrogen is attached to O, N, or F, it becomes strongly partially positive.

• The nearby electronegative atom has a partially negative region and attracts the hydrogen.

2. Types of Hydrogen Bonds

A. Intermolecular hydrogen bond

“Inter” means between. This bond forms between two separate molecules.

• Example: one water molecule attracts another water molecule.

• It can make molecules stick together in groups or chains.

• It often raises boiling point and melting point because more energy is needed to separate the molecules.

B. Intramolecular hydrogen bond

“Intra” means within. This bond forms inside the same molecule, between parts of that one molecule.

• It often happens when the molecule has suitable groups positioned close to each other.

• Examples shown in the reference: o-nitrophenol and o-hydroxybenzaldehyde.

• It can affect the molecule’s shape, acidity, and physical properties.

Memory tip: INTER = between molecules; INTRA = inside one molecule.

3. Properties of Hydrogen Bonds

• Weak compared with covalent bonds: the reference gives about 3–10 kcal/mol for hydrogen bonds versus about 50–100 kcal/mol for ordinary covalent bonds. Exact values vary by substance and conditions.

• Hydrogen bonding is a type of dipole–dipole attraction involving hydrogen attached to a strongly electronegative atom.

• The hydrogen is usually closer to one of the two electronegative atoms involved.

• Small, highly electronegative atoms (especially N, O, and F) are most effective at forming strong hydrogen bonds.

• Hydrogen bonds have direction: the atoms need to be arranged in a suitable way for the attraction to be strong.

Important comparison: H₂O is liquid at room temperature, while H₂S is a gas. Hydrogen bonding between water molecules helps explain water’s unusually high boiling point.

4. Effects and Importance of Hydrogen Bonding

1. Water and ice: Hydrogen bonds create an open structure in ice, so ice occupies more volume than the same mass of liquid water. This makes ice less dense, allowing it to float.

2. Water density: Liquid water reaches its maximum density at about 4°C. This is why ice can float on liquid water.

3. Mixing of liquids: Hydrogen bonding helps water mix with alcohols such as ethanol. Water and alcohol can form hydrogen bonds with each other.

4. DNA and RNA: Hydrogen bonds help hold complementary parts of DNA together and contribute to the structures of biological molecules.

5. Proteins: Hydrogen bonds help stabilize the shapes of proteins.

6. Paints, dyes, and fibres: Hydrogen bonding can contribute to adhesion and the strength or structure of some materials.

7. Hydrated crystals: Water can be held in some crystals, such as CuSO₄·5H₂O, as part of their crystal structure. Do not assume every hydrate is explained only by hydrogen bonding.

How can hydrogen bonding be detected?

• Infrared (IR) spectroscopy: hydrogen bonding can change or broaden the O–H or N–H stretching signal.

• Proton magnetic resonance / NMR: hydrogen bonding can change the signal position of a hydrogen nucleus.

• X-ray and electronic absorption methods can also provide evidence, depending on the substance.

5. Theories of Hydrogen Bonding

A. Pimentel theory

Easy meaning: This explanation uses Molecular Orbital Theory. Atomic orbitals combine to form molecular orbitals, including bonding, non-bonding, and antibonding orbitals.

• The reference discusses the interaction of hydrogen’s 1s orbital with p orbitals on atoms A and B.

• The theory explains hydrogen bonding through orbital interactions and the arrangement of electrons.

B. Coulson theory

Easy meaning: This explanation focuses on electrostatic attraction.

• When hydrogen is bonded to a strongly electronegative atom (such as O, N, or F), hydrogen becomes partially positive (δ+).

• A nearby electronegative atom has a partially negative region (δ−).

• The attraction between these regions forms the weak hydrogen bond, sometimes called an H-bridge.

• The reference notes that this theory explains intermolecular hydrogen bonding better than intramolecular hydrogen bonding.

C. McClellan theory

Easy meaning: This theory explains hydrogen bonding using resonance—different electron arrangements contribute to the overall structure.

• The molecule is described as having several contributing structures, rather than just one fixed arrangement.

• The reference mentions electrostatic interactions, attraction, electron–electron repulsion, and weak van der Waals forces.

• The combined effects help explain the character of the hydrogen bond.

D. Valence Bond Theory (VBT)

Easy meaning: VBT explains hydrogen bonding using electrostatic attraction and resonance structures.

• Different possible bonding arrangements contribute to the overall bond.

• The reference gives an equation for estimating hydrogen-bond energy, but for a basic review, focus on the idea rather than memorizing the formula unless your teacher requires it.

6. Metallic Bond

Easy meaning: A metallic bond is the attraction between positive metal ions and a “sea” of mobile, delocalized valence electrons shared throughout the metal.

Imagine positive metal ions arranged together while electrons can move around them. The attraction holds the metal together.

How metallic bonding works

• Metal atoms can release or share their outer (valence) electrons relatively easily.

• The positive metal ions stay arranged in a structure.

• The mobile electrons move through the structure and attract the positive ions, helping hold the metal together.

Properties of metals explained by metallic bonding

• Electrical conductivity: mobile electrons move when an electric field is applied, carrying charge.

• Thermal conductivity: mobile electrons transfer energy through the metal.

• Metallic lustre: metals reflect light and appear shiny.

• Malleability: metals can be hammered into sheets because layers of ions can shift while the electron sea continues to hold the structure together.

• Ductility: metals can be drawn into wires for a similar reason.

• Strength and often high density: many metals have closely packed structures and strong attractions, although values vary by metal.

• Alloy formation: metals can mix with other elements to form alloys.

• Opacity: most bulk metals do not allow visible light to pass through them.

Common metallic crystal structures

• FCC – Face-Centred Cubic

• BCC – Body-Centred Cubic

• HCP – Hexagonal Close-Packed

Remember: not every metal has all these structures. Different metals may adopt different crystal arrangements.

7. Quick Comparison

8. Practice Questions

8. What is a hydrogen bond?

9. Which three atoms most commonly participate in strong hydrogen bonding?

10. What does the dotted line (···) represent?

11. What is the difference between intermolecular and intramolecular hydrogen bonding?

12. Why does ice float on liquid water?

13. How does hydrogen bonding help explain water’s high boiling point?

14. Name two methods that can help detect hydrogen bonding.

15. What is the main idea of Coulson theory?

16. What does Pimentel theory use to explain hydrogen bonding?

17. What is a metallic bond?

18. Why do metals conduct electricity?

19. Why can metals be hammered into sheets or drawn into wires?

20. What do FCC, BCC, and HCP stand for?

21. Give two biological roles of hydrogen bonds.

22. What is the difference between a covalent bond and a hydrogen bond?

Answer Key

1. A weak attraction between a partially positive hydrogen and a nearby electronegative atom/region.

2. Nitrogen (N), oxygen (O), and fluorine (F).

3. A hydrogen bond.

4. Intermolecular is between molecules; intramolecular is within one molecule.

5. Hydrogen bonds create an open structure in ice, making it less dense than liquid water.

6. Extra energy is needed to separate water molecules connected by hydrogen bonds.

7. Infrared spectroscopy and proton NMR; X-ray methods can also help.

8. It explains hydrogen bonding mainly through electrostatic attraction between partial charges.

9. Molecular Orbital Theory and interactions among orbitals.

10. Attraction between positive metal ions and delocalized mobile electrons.

11. Mobile electrons can move and carry electric charge.

12. Layers can shift while the mobile electrons continue to hold the metal together.

13. Face-Centred Cubic, Body-Centred Cubic, and Hexagonal Close-Packed.

14. They help stabilize DNA/RNA structures and protein shapes.

15. Covalent bonds involve sharing electron pairs within a molecule and are generally stronger; hydrogen bonds are weaker attractions between suitable polar regions.

Inorganic Chemistry Reviewer | Simplified study notes