Notes on Chemical Bonding and Properties of Compounds (Pages 10–18)

Page 10 — Lesson 1: Chemical Bonding — Properties of Compounds

Context: Building on the Quantum Mechanical Model of the atom to describe electron behavior; focus on recognizing ionic vs covalent compounds through their properties.
Key questions to ponder after the module:
- 1) How to identify and describe ionic and covalent compounds from their chemical formula and chemical name?
- 2) What are the different physical properties of ionic and covalent compounds?
- 3) How to distinguish ionic from covalent compounds based on physical properties?
- 4) What natural phenomena rely on the distinct properties of ionic vs covalent compounds?
The modern periodic table and the concept of groups:
- Elements in the same vertical column (group) share common properties.
- Activity: Classify the elements below into their representative-group categories using chemical symbols. Elements provided: Pb (Lead), He (Helium), Ba (Barium), B (Boron), O (Oxygen), Al (Aluminum), Ar (Argon), Cl (Chlorine), K (Potassium), Rn (Radon), N (Nitrogen), Mg (Magnesium), Na (Sodium), H (Hydrogen), P (Phosphorus).
- Note: Answer is to be completed on separate paper as part of the exercise; this section reinforces group patterns among representative elements.
Figure reference: Page mentions "Figure 1: Periodic Table of the Representative Elements" (contextual visualization).
Summary takeaway: Recognize how groups relate to element properties and how to map elements to the correct representative groups for quick property inference.

Page 11 — What’s New: Getting to Know Chemical Compounds

Focus: Identify the name of elements in given compounds and determine whether the elements are metals or nonmetals.
The table (example item given):
- 1. Water (H₂O): Elements involved — Hydrogen (H) and Oxygen (O); Type: Nonmetal and Nonmetal.
Items to classify (activities to be written on separate paper):
- 2. Sugar (C₁₂H₂₂O₁₁): Elements involved — C, H, O; All nonmetals; Type: Covalent compound.
- 3. Potassium chloride (KCl): Elements involved — K (potassium, metal) and Cl (chlorine, nonmetal); Type: Ionic compound.
- 4. Carbon Dioxide (CO₂): Elements involved — C and O; Both nonmetals; Type: Covalent compound.
- 5. Sodium Chloride (NaCl): Elements involved — Na (sodium, metal) and Cl (chlorine, nonmetal); Type: Ionic compound.
Companion questions:
- 1) How do compounds form?
- 2) What does a compound contain?
Figure reference:
- "What's New" section alludes to a compound table used for practice with element-type identification.
Takeaway: Practice recognizing ionic vs covalent character by inspecting element types within common compounds and by noting whether the elements involved are metals or nonmetals.

Page 12 — What is It? Let’s Bond with Ionic and Covalent Compounds

Core definitions:
- An ionic compound forms when a metal (cation) transfers valence electrons to a nonmetal (anion).
- A covalent compound forms when nonmetals share valence electrons with another nonmetal.
Visual aid: A diagram illustrating electron transfer or sharing to achieve stability (as described in the module).
Examples to classify by formula/name:
- Carbon Dioxide, CO₂, is covalent (nonmetals: C and O).
- Sodium Chloride, NaCl, is ionic (Na is a metal; Cl is a nonmetal).
Practical instruction: Use chemical formula and the types of elements involved to determine the bond type; consult the representative compounds table in the module for practice.
Takeaway: The distinction hinges on whether electrons are transferred (ionic) or shared (covalent) between atoms.

Page 13 — Understanding the Different Properties of Compounds

Major concept: Ionic and covalent compounds exhibit distinct physical properties that help distinguish them.
1) At room temperature and pressure:
- Covalent compounds can exist as solids, liquids, or gases.
- Ionic compounds exist as crystalline solids.
- Reason: Covalent bonds involve sharing electrons; weaker inter-molecular attractions lead to less rigid overall attraction than in ionic crystals.
- Implication example: Alcohol (a covalent compound) can be a liquid at room temperature and widely used as disinfectant, fuel, and component of beverages.
- Ionic solids form a rigid crystal lattice (e.g., NaCl) and are typically solids at room temperature.
2) Melting and boiling points:
- Ionic compounds generally have higher melting and boiling points due to the energy required to overcome strong ionic lattice forces.
- Covalent compounds generally have lower melting and boiling points due to weaker intermolecular forces between molecules.
Mara’s scenario (salt vs sugar):
- Two identical white powders at room temperature: one is NaCl (ionic) and the other C₁₂H₂₂O₁₁ (sugar, covalent).
- Without tasting, use properties such as solubility, conductivity, conductivity in solution, melting behavior, and crystallinity to distinguish.
Takeaway: Physical-property differences are practical diagnostic tools for identifying ionic vs covalent substances in everyday contexts.

Page 14 — Ionic vs Covalent: Hardness, Polarity, and Bonding Characteristics

3) Hardness and brittleness:
- Ionic compounds: hard and brittle due to rigid crystal lattices.
- Covalent compounds: comparatively soft and flexible due to weaker inter-molecular forces and molecular discreteness.
- Crystal lattices influence hardness and brittleness; ions arranged in a lattice lead to rigidity but brittleness under distortion.
4) Polarity and electronegativity:
- Polarity arises from differences in electronegativity between bonded atoms.
- Ionic character criterion: when electronegativity difference (ΔEN) between the metallic and nonmetallic elements is greater than 1.9, the compound is ionic.
- Covalent bonds may be polar or nonpolar:
- Polar covalent: ΔEN between 0.5 and 1.9.
- Nonpolar covalent: ΔEN less than 0.5.
- Diatomic molecules of identical nonmetals (e.g., N₂, O₂, H₂) are nonpolar covalent.
5) Solubility in water:
- Ionic compounds are usually highly soluble in water because water's polarity stabilizes and separates ions.
- Covalent compounds are often less soluble in water; solubility correlates with polarity and “like dissolves like.”
- Note: When a solvent contains a common ion, solubility can decrease (common-ion effect).
4(b)/(5) Practical example: salt vs sugar solubility and polarity considerations align with these principles.
Figure reference: Conceptual illustration of crystal lattices for NaCl is noted in the module.

Page 15 — Practical Example: Electronegativity, Polarity, and Bonding

Example table (selected items):
- 1) Sodium Chloride (NaCl): EN(Na) = 0.9; EN(Cl) = 3.0; ΔEN = 3.0 − 0.9 = 2.1 → Ionic.
- 2) Water (H₂O): EN(H) = 2.1; EN(O) = 3.5; ΔEN = 3.5 − 2.1 = 1.4 → Polar covalent.
- 3) Hydrogen gas (H₂): EN(H) = 2.1; EN(H) = 2.1; ΔEN = 0 → Nonpolar covalent.
- 4) (Not explicitly shown in the excerpt) [Potential item]: commonly used covalent compounds with varying ΔEN.
5) Ionic compounds are usually soluble in water; covalent compounds tend to be less soluble; “Like dissolves like” as a guiding principle when considering polarity alignment.
Takeaway: Quantitative electronegativity differences provide a framework to classify bonds as ionic, polar covalent, or nonpolar covalent.

Page 16 — Fluency in Flammability and Electrical/Thermal Conductivity

6) Flammability:
- Ionic compounds tend to be less flammable than covalent compounds.
- Combustion overview: hydrocarbons (common in covalent compounds) burn in the presence of oxygen to form CO₂ and H₂O.
- Notable exception: water (H₂O), a covalent compound with strong intramolecular bonds and polar characteristics, is not easily ignited.
- Example: Liquefied Petroleum Gas (LPG) is a flammable covalent mixture used for cooking and heating; safety regulations (e.g., DOE circular) underscore handling requirements.
7) Conductivity of heat and electricity:
- Ionic compounds: good conductors of electricity when dissolved in water or molten, due to mobile ions; good conductors of heat because ions in lattice can transfer energy efficiently.
- Covalent compounds: generally good insulators for electricity and heat; electrons are shared, and molecules are less tightly bound, reducing energy transfer efficiency.
Important caveat: These are general properties with exceptions; context (state, solvent, temperature) matters.

Page 16–17 — Natural Phenomena That Use Different Properties

A. Frozen Fractals of Snowflakes

Philippines climate context:
- Tropical country with high humidity (PAGASA data): 71% in March to 85% in September.
- Rare hailstone events noted in 2017 (not typical snowfall).
Snow vs hail:
- Hailstone: ice pellets formed during thunderstorms via updrafts suspending ice particles; distinct from snow.
- Snow: formed when temperature is low and moisture in the atmosphere forms tiny ice crystals in clouds that stick together to become snowflakes; typically hexagonal symmetry.
Snowflake chemistry and bonding (as described):
- Snowflakes are composed of water; the intermolecular interaction described is a covalent bond in the slide's framing, though standard chemistry classically emphasizes hydrogen bonding (an intermolecular force) in ice crystals.
- Structure: the hexagonal six-sided pattern arises from the bent shape of the H–O–H arrangement and hydrogen bonding leading to a dipole character (oxygen side is more negative; hydrogen side is more positive).
Takeaway: Snowflakes illustrate covalent-influenced molecular geometry (via water's V-shaped molecule) and hydrogen-bonding networks underpin crystal structure. Note: the slide attributes a covalent bond to snowflakes; scientifically, hydrogen bonding is the key intermolecular interaction, while the covalent O–H bonds hold the H₂O units together internally.

B. Chemical Reactions as a Source of Power

Electrochemistry: study of processes that move electrons to generate electricity; a means to address environmental concerns from burning fossil fuels.
Galvanic (Voltaic) cells:
- Definition: Electrochemical cells that convert chemical energy into electrical energy via oxidation-reduction (redox) reactions.
- Everyday example: Batteries powering devices like cell phones.
Basic setup described:
- Anode (oxidation) and cathode (reduction) electrodes immersed in an electrolyte solution.
- Zinc and copper electrodes are used in two separate solutions: ZnSO₄ and CuSO₄, respectively.
- A connecting wire allows electrons to flow from the zinc electrode (anode) to the copper electrode (cathode).
Redox processes described (as per the module):
- Oxidation at the zinc electrode:
  ext{Zn}
  ightarrow ext{Zn}^{2+} + 2e^-
- Reduction at the copper electrode:
  ext{Cu}^{2+} + 2e^-
  ightarrow ext{Cu}
Overall cell reaction (net reaction):
- ext{Zn} + ext{Cu}^{2+}
  ightarrow ext{Zn}^{2+} + ext{Cu}
Takeaway: Galvanic cells demonstrate practical energy generation through controlled electron transfer, linking chemical bonding concepts to real-world electricity production.

Page 18 — Galvanic Cell Details (continued) and Connections

Recap of the galvanic cell setup and operation:
- Two different metal electrodes in aqueous metal salt solutions, connected by a conductor (wire) to enable electron flow.
- The electrode with lower tendency to gain electrons (zinc) is oxidized (loses electrons); the electrode with higher tendency to gain electrons (copper) is reduced (gains electrons).
- Mobile metal ions in solution balance charge changes as the reaction proceeds.
Significance for exam preparation:
- Recognize the key components and the flow of electrons in a galvanic cell.
- Be able to write oxidation and reduction half-reactions and the overall cell reaction.
Formula recap (for quick reference):
- Oxidation: ext{Zn}
  ightarrow ext{Zn}^{2+} + 2e^-
- Reduction: ext{Cu}^{2+} + 2e^-
  ightarrow ext{Cu}
- Overall: ext{Zn} + ext{Cu}^{2+}
  ightarrow ext{Zn}^{2+} + ext{Cu}
Overall study-oriented synthesis:
- Ionic vs covalent bonding is reflected across multiple properties: melting/boiling points, hardness, solubility, polarity, conductivity, and flammability.
- These properties connect to real-world phenomena (snow, hail, electrochemistry) and everyday materials (water, salt, sugar, LPG).
- The module emphasizes using chemical formulas and the identities of elements (metals vs nonmetals) to classify compounds and predict behavior.
Quick reference formulas and constants:
- Covalent compound example: CO2, H2O, C{12}H{22}O_{11}
- Ionic compound example: NaCl
- Polarity thresholds:
- ext{If } riangle EN > 1.9 ext{ then Ionic}
- 0.5
  le riangle EN
  < 1.9 ext{ then Polar covalent}
- riangle EN < 0.5 ext{ then Nonpolar covalent}
- Hydrogen bonds (ice/snow context) involve strong intermolecular interactions despite individual O–H covalent bonds within H₂O molecules.
Key takeaways for exam readiness:
- Be able to classify compounds from formulas and names by identifying metal vs nonmetal participation.
- Distinguish ionic and covalent compounds by examining melting points, state at room temperature, hardness, solubility, polarity, and conductivity.
- Apply electronegativity differences to categorize bonds and predict solubility and conductivity tendencies.
- Understand real-world phenomena (snow vs hail, LPG safety) as practical illustrations of compound properties in action.