Chapter 3: Chemical Bonds

Answer:

Elements that lose electrons easily: Metals, especially alkali metals (Group 1: Li, Na, K, etc.) and alkaline earth metals (Group 2: Mg, Ca, etc.). These are electropositive elements with low ionization energies.

Elements that gain electrons easily: Non-metals, especially halogens (Group 17: F, Cl, Br, etc.) and oxygen group elements (Group 16: O, S, etc.). These are electronegative elements with high electron affinity.

Answer: Diamond (allotropic form of carbon).

Explanation: Diamond is a crystalline solid where each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement, forming a giant covalent network structure. Other examples include silicon, quartz (SiO₂), and graphite (different structure).

Answer: Yes, once formed, a coordinate covalent bond is as strong as a normal covalent bond.

Explanation: The difference between coordinate covalent bond and normal covalent bond is only in their formation (in coordinate bond, one atom donates both electrons). Once the bond is formed, it is identical in strength and properties to other covalent bonds in the molecule. For example, in ammonium ion (NH₄⁺), all four N-H bonds are equivalent despite one being originally a coordinate bond.

Answer: Nitric acid (HNO₃) has the following Lewis structure:

Detailed structure:

1. H bonded to O (single bond)
2. That O bonded to N (single bond)
3. N has double bond to one O
4. N has single bond to another O (with negative charge)
5. N has formal positive charge
6. Resonance exists between the two N-O bonds

Answer: HF is a liquid at room temperature while HCl is a gas due to differences in intermolecular forces.

1. HF has hydrogen bonding: Fluorine is highly electronegative (4.0) and hydrogen bonding occurs between H of one molecule and F of another (F-H···F). This strong intermolecular force increases boiling point to 19.5°C.
2. HCl has dipole-dipole forces: Chlorine is less electronegative (3.0) than fluorine, so hydrogen bonding is not significant in HCl. It has weaker dipole-dipole forces, resulting in lower boiling point (-85°C).
3. Molecular size: HF molecules are smaller but the strong hydrogen bonding overcomes this factor.
4. Comparison: Boiling points: HF (19.5°C) > HCl (-85°C) > HBr (-67°C) > HI (-35°C). The trend reverses after HF due to decreasing hydrogen bonding strength.

Answer: Covalent compounds are generally insoluble in water due to:

1. Polarity difference: Water is a highly polar solvent (due to O-H bonds and bent structure). Most covalent compounds are non-polar or have low polarity.
2. “Like dissolves like” principle: Polar solvents dissolve polar solutes; non-polar solvents dissolve non-polar solutes.
3. Intermolecular forces: Water molecules form strong hydrogen bonds with each other. To dissolve, solute must break some H-bonds and form new interactions. Non-polar covalent compounds cannot form strong interactions with water.
4. Energy consideration: The energy required to break water’s hydrogen bonds is not compensated by weak solute-water interactions.
5. Exceptions: Some covalent compounds with -OH, -COOH, -NH₂ groups can hydrogen bond with water and are soluble (e.g., ethanol, acetic acid).

Answer: Metals conduct heat through their delocalized electrons in the “electron sea”.

1. Metallic structure: Metals consist of positive ions in a sea of delocalized electrons.
2. Heat conduction mechanism: When heat is applied to one part of a metal:
   • Electrons near the heat source gain kinetic energy
   • These energetic electrons move rapidly throughout the metal
   • They collide with other electrons and positive ions, transferring kinetic energy
   • This rapid transfer of energy conducts heat efficiently
3. Comparison: In non-metals, heat is conducted by slower vibrational energy transfer through the lattice.
4. Relationship with electrical conductivity: The same mobile electrons that conduct heat also conduct electricity, explaining why good thermal conductors are usually good electrical conductors.
5. Examples: Silver and copper are excellent heat conductors used in cookware and heat sinks.

Answer: Nitrogen forms several oxides with different oxidation states:

1. Nitrous oxide: N₂O (oxidation state of N: +1)
2. Nitric oxide: NO (oxidation state of N: +2)
3. Dinitrogen trioxide: N₂O₃ (oxidation state of N: +3)
4. Nitrogen dioxide: NO₂ (oxidation state of N: +4)
5. Dinitrogen tetroxide: N₂O₄ (dimer of NO₂, oxidation state of N: +4)
6. Dinitrogen pentoxide: N₂O₅ (oxidation state of N: +5)

Common oxides: The most common nitrogen oxides are NO (nitric oxide), NO₂ (nitrogen dioxide), and N₂O (nitrous oxide).

Answer:

Ionic Bond Formation:

1. Definition: Ionic bond is formed by complete transfer of electrons from one atom to another.
2. Process: A metal atom loses electron(s) to form a positive ion (cation). A non-metal atom gains electron(s) to form a negative ion (anion).
3. Driving force: Atoms achieve stable noble gas electron configuration (octet/duplet).
4. Example (NaCl):
   • Na (2,8,1) loses 1 electron → Na⁺ (2,8) like Ne
   • Cl (2,8,7) gains 1 electron → Cl⁻ (2,8,8) like Ar
   • Na⁺ and Cl⁻ attract electrostatically → ionic bond
5. Properties: High melting/boiling points, soluble in water, conduct electricity when molten/dissolved.

Covalent Bond Formation:

1. Definition: Covalent bond is formed by mutual sharing of electron pairs between atoms.
2. Process: Atoms share electrons to achieve stable electron configuration.
3. Types:
   • Single bond: One shared pair (H-H)
   • Double bond: Two shared pairs (O=O)
   • Triple bond: Three shared pairs (N≡N)
4. Example (H₂O):
   • O (2,6) needs 2 electrons to complete octet
   • Each H (1) needs 1 electron to complete duplet
   • O shares 2 electrons (1 with each H) → 2 single covalent bonds
5. Properties: Low melting/boiling points, poor conductors, form discrete molecules.

Answer: Metallic properties explained by metallic bonding (“electron sea” model):

1. Metallic luster (shininess):
   • Delocalized electrons absorb light energy
   • They re-emit light at same frequency
   • Creates shiny appearance
2. High melting and boiling points:
   • Strong attraction between positive ions and electron sea
   • Requires high energy to overcome
   • Strength depends on: charge on ions and number of delocalized electrons
   • Example: Mg (2+ ions) has higher mp than Na (1+ ions)
3. Good electrical conductivity:
   • Delocalized electrons are free to move
   • When voltage applied, electrons drift toward positive terminal
   • Creates electric current
4. Good thermal conductivity:
   • Electrons transfer kinetic energy rapidly
   • Heat causes electrons to move faster
   • They collide with other electrons/ions, transferring energy
5. Malleability and ductility:
   • Layers of ions can slide over each other
   • Electron sea adjusts to new positions
   • Bonds not broken, just reformed
   • Allows shaping without fracturing
6. Hardness and density:
   • Close packing of atoms in regular arrangement
   • Strong metallic bonds
   • Results in generally hard, heavy materials

Answer: Metals are usually hard and heavy due to:

1. Close-packed structure:
   • Metal atoms arrange in regular, close-packed patterns (FCC, BCC, HCP)
   • Atoms are closely packed together
   • Minimizes empty space between atoms
2. Strong metallic bonding:
   • Attraction between positive ions and electron sea is strong
   • Requires force to separate atoms
   • Bond strength increases with more delocalized electrons
3. High density:
   • Close packing means more mass per unit volume
   • Metal atoms are relatively heavy (except Group 1 metals)
   • Examples: Osmium (22.59 g/cm³), platinum (21.45 g/cm³)
4. Exceptions:
   • Some metals are soft: Na, K (can be cut with knife)
   • Some are light: Li (0.534 g/cm³), Al (2.70 g/cm³)
   • Hardness varies with bond strength: W, Cr are very hard; Hg is liquid
5. Comparison with non-metals:
   • Non-metals often have open structures (graphite, diamond exception)
   • Weaker intermolecular forces
   • Generally lower densities