Valence Electrons and Chemical Reactions

Valence Electrons and Atomic Stability

Atoms contain electrons arranged in shells around the nucleus. The electrons in the outermost shell are called valence electrons, and they play a crucial role in chemical reactivity.

• Complete valence shells (as in noble gases) are energetically stable.

• Incomplete valence shells are energetically unstable, making atoms more likely to react to achieve stability.

• Atoms with unfilled valence shells can become stable by transferring (giving up or gaining) electrons or by sharing electrons with other atoms to complete their valence shell.

Example: Helium (He) has a full valence shell and is chemically inert, while sodium (Na) and chlorine (Cl) have incomplete shells and readily react to form compounds.

Electron Configuration and the Periodic Table

The arrangement of electrons in shells determines the chemical properties of elements. The periodic table groups elements with similar valence electron configurations together.

• First shell: up to 2 electrons

• Second shell: up to 8 electrons

• Third shell: up to 8 electrons (for main group elements)

Example: Oxygen (O) has 6 valence electrons and needs 2 more to complete its shell, making it highly reactive.

Chemical Bonds

Types of Chemical Bonds

Chemical bonds form when atoms transfer or share valence electrons to achieve stable electron configurations.

• Ionic bonds: Formed by transferring electrons from one atom to another, resulting in oppositely charged ions that attract each other.

• Covalent bonds: Formed by sharing pairs of electrons between atoms.

Example: Sodium (Na) transfers an electron to chlorine (Cl), forming Na+ and Cl- ions, which combine to make sodium chloride (NaCl).

Ionic Bonds

Ionic bonds occur when one atom donates an electron to another, creating ions with opposite charges.

• Cation: Positively charged ion (e.g., Na+).

• Anion: Negatively charged ion (e.g., Cl-).

• Ionic compounds often form crystalline solids and dissociate easily in water.

Example: Na (sodium) + Cl (chlorine) → NaCl (table salt).

Covalent Bonds

Covalent bonds involve the sharing of one or more pairs of electrons between atoms. Each shared pair constitutes a single covalent bond.

• Single bond: One pair of shared electrons (e.g., H–H in H2).

• Double bond: Two pairs of shared electrons (e.g., O=O in O2).

• Shared electrons count as part of each atom’s valence shell.

Example: Two hydrogen atoms share electrons with one oxygen atom to form water (H2O).

Bond Notation

Chemical bonds can be represented in several ways:

• Structural formula: Shows bonds as lines (e.g., H–H, O=O).

• Electron dot structure: Shows shared and unshared valence electrons (e.g., H:H).

• Molecular formula: Abbreviated notation showing the number and type of atoms (e.g., H2O, CH4).

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is an atom’s ability to attract shared electrons in a chemical bond. The greater the electronegativity difference between two atoms, the more polar the bond.

• Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Fluorine (F) is the most electronegative element.

Bond type by electronegativity difference:

• Difference > 2.0: Ionic bond

• Difference < 2.0: Covalent bond

Bond Polarity

• Nonpolar covalent bond: Electrons are shared equally; no partial charges (e.g., O2).

• Polar covalent bond: Electrons are shared unequally; partial charges develop (e.g., H2O).

• Ionic bond: Complete transfer of electrons; full charges on resulting ions (e.g., NaCl).

Hydrogen Bonds

Definition and Importance

Hydrogen bonds are weak attractions between a hydrogen atom covalently bonded to a highly electronegative atom (such as oxygen or nitrogen) and another electronegative atom. Though weaker than covalent or ionic bonds, hydrogen bonds are essential for the structure and function of biological molecules.

• Responsible for the unique properties of water.

• Stabilize the structures of DNA, RNA, and proteins.

Example: Hydrogen bonds between water molecules give water its high boiling point and surface tension.

Elements in the Human Body

Major and Trace Elements

The human body is composed of a limited number of elements, with a few making up the majority of body mass.

Additional info: Trace elements include boron, chromium, cobalt, copper, fluorine, iodine, iron, manganese, molybdenum, selenium, silicon, tin, and vanadium.

Properties of Water and Their Biological Importance

Emergent Properties of Water

Water exhibits several unique properties that make it essential for life on Earth. These properties arise from its molecular structure and hydrogen bonding.

• Ability to moderate temperature

• Cohesive behavior

• Expansion upon freezing

• Versatility as a solvent

High Specific Heat Capacity

Water has a high specific heat capacity, meaning it can absorb or release large amounts of heat with only a slight change in its own temperature.

• Specific heat of water:

• This property helps stabilize temperatures in organisms and environments.

Example: Oceans moderate coastal climates; human body temperature remains relatively constant.

Cohesion and Surface Tension

Cohesion is the attraction between water molecules due to hydrogen bonding, resulting in high surface tension.

• Allows water to form droplets and move through plant vessels (capillary action).

• Enables small insects to walk on water surfaces.

Expansion Upon Freezing

Unlike most substances, water expands as it freezes, making ice less dense than liquid water. This is due to the stable hydrogen bonds that form a crystalline lattice in ice.

• Ice floats on water, insulating aquatic life in winter.

• If ice sank, bodies of water would freeze solid from the bottom up, threatening aquatic ecosystems.

Water as a Solvent

Water is known as the "universal solvent" because it can dissolve a wide variety of substances, especially ionic and polar compounds.

• Solution: Homogeneous mixture of two or more substances.

• Solvent: The dissolving agent (water in aqueous solutions).

• Solute: The substance being dissolved.

• Water forms hydration shells around ions and polar molecules, facilitating their dissolution.

Example: Table salt (NaCl) dissolves in water to form Na+ and Cl- ions surrounded by water molecules.

Hydrophilic and Hydrophobic Substances

• Hydrophilic: Substances that are attracted to water and dissolve easily (e.g., salts, sugars, polar molecules).

• Hydrophobic: Substances that repel water and do not dissolve (e.g., oils, nonpolar molecules).

• Hydrophobic molecules aggregate in water (e.g., oil droplets in water).

Example: Cell membranes are composed of hydrophobic lipid bilayers that separate the cell from its environment.