Atomic Structure, Elements, and Chemical Bonds: Foundations of General Biology

Atomic Structure and Matter

Introduction to Matter and Atoms

All living and non-living things are composed of matter, which is governed by the laws of physics and chemistry. Understanding the structure of matter is fundamental to biology, as it underpins the composition and behavior of all biological molecules.

• Matter: Anything that takes up space and has mass.

• Element: A pure substance consisting of only one kind of atom (e.g., oxygen, hydrogen).

• Atom: The smallest unit of an element that retains the properties of that element (e.g., O, H, C).

• Molecule: Two or more atoms joined together (e.g., O2, H2, H2O).

• Compound: A molecule that contains more than one element (e.g., H2O, C6H12O6).

Elements Essential for Life

About 25 elements are essential for life, with a few making up the majority of living matter.

• Major elements: Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N) constitute about 96% of the human body mass (excluding water).

• Other essential elements: Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), Chlorine (Cl), Magnesium (Mg), and trace elements (e.g., Iron (Fe), Iodine (I)).

Atomic Structure

Subatomic Particles

Atoms are composed of three main subatomic particles:

• Protons: Positively charged particles found in the nucleus; determine the atomic number and identity of the element.

• Neutrons: Neutral particles found in the nucleus; contribute to atomic mass and isotope formation.

• Electrons: Negatively charged particles that orbit the nucleus in electron shells; determine chemical behavior and bonding.

Atomic Number, Mass, and Isotopes

• Atomic Number (Z): Number of protons in the nucleus; unique to each element.

• Atomic Mass: Total number of protons and neutrons in an atom; measured in Daltons (Da).

• Isotopes: Atoms of the same element with different numbers of neutrons. Some isotopes are stable, while others are radioactive and decay over time, emitting particles and energy.

Example: Carbon-12 (12C) and Carbon-14 (14C) are isotopes of carbon; 14C is radioactive and used in radiometric dating.

Electron Shells and Valence

Electrons are arranged in shells around the nucleus, each with a specific energy level. The arrangement of electrons determines an atom's reactivity.

• The first shell holds up to 2 electrons; the second and third shells hold up to 8 electrons each.

• Valence electrons: Electrons in the outermost shell; determine the chemical properties and reactivity of the atom.

• Atoms with full valence shells are chemically inert (noble gases), while those with incomplete shells are reactive.

Chemical Bonds and Electronegativity

Types of Chemical Bonds

Atoms form chemical bonds to achieve stable electron configurations. The main types of bonds are:

• Covalent Bonds: Atoms share pairs of valence electrons. Can be single, double, or triple bonds depending on the number of shared electron pairs.

• Ionic Bonds: Atoms transfer electrons, resulting in oppositely charged ions that attract each other.

• Hydrogen Bonds: Weak attractions between a hydrogen atom covalently bonded to an electronegative atom (like O or N) and another electronegative atom.

Electronegativity and Bond Polarity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. It depends on atomic number and the distance of valence electrons from the nucleus.

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

• When two atoms with different electronegativities form a bond, the electrons may be shared unequally, resulting in a polar covalent bond.

• If the electronegativity difference is large, an ionic bond forms; if small or zero, a nonpolar covalent bond forms.

Example: In water (H2O), oxygen is more electronegative than hydrogen, so the shared electrons are pulled closer to oxygen, making the molecule polar.

Bond Classification by Electronegativity Difference

• Nonpolar Covalent: Electronegativity difference < 0.4 (e.g., H2, O2).

• Polar Covalent: Electronegativity difference between 0.4 and 2.0 (e.g., H2O).

• Ionic: Electronegativity difference > 2.0 (e.g., NaCl).

Common Ions in Biology

Atoms that gain or lose electrons become ions. Cations are positively charged (lost electrons), and anions are negatively charged (gained electrons).

Ion Name	Symbol	Charge	Electron Change
Sodium ion	Na+	+1	Lost 1 electron
Chloride ion	Cl-	-1	Gained 1 electron
Magnesium ion	Mg2+	+2	Lost 2 electrons
Potassium ion	K+	+1	Lost 1 electron
Hydrogen ion	H+	+1	Lost 1 electron

Summary Table: Key Terms and Definitions

Term	Definition	Example
Matter	Anything that takes up space and has mass	Water, air, rocks
Element	Pure substance of one kind of atom	Oxygen (O), Hydrogen (H)
Atom	Smallest unit of an element	O, H, C
Molecule	Two or more atoms joined together	O2, H2O
Compound	Molecule with more than one element	H2O, C6H12O6
Isotope	Atoms of the same element with different numbers of neutrons	12C, 14C

Key Equations

• Atomic Mass: $\text{Atomic Mass} = \text{Number of Protons} + \text{Number of Neutrons}$

• Electronegativity Difference (ΔEN): $\Delta EN = |EN_1 - EN_2|$

Additional info:

• Some context and examples were inferred to provide a complete, self-contained study guide.

• Tables were reconstructed and expanded for clarity and completeness.