Chapter 2

Elements and the Periodic Table

Overview of Elements

The Periodic Table organizes all known chemical elements, which are fundamental to biological systems. Understanding the properties and organization of elements is essential for studying biological molecules and processes.

• 118 elements are currently recognized.

• 92 elements occur naturally on Earth; the rest are man-made.

• 25 elements are found in living organisms, with 4 elements (H, O, N, C) making up 96% of living matter.

• Other important elements in biology include Sulfur (S) and Phosphorus (P).

• Mnemonic: CHON (Carbon, Hydrogen, Oxygen, Nitrogen) and SPONCH (Sulfur, Phosphorus, Oxygen, Nitrogen, Carbon, Hydrogen).

Sub-atomic Particles

Structure of the Atom

Atoms are composed of three main sub-atomic particles: protons, neutrons, and electrons. The arrangement and number of these particles determine the properties of each element.

• Protons: Positively charged, located in the nucleus, mass ≈ 1 Dalton, defines the atomic number.

• Neutrons: Neutral charge, located in the nucleus, mass ≈ 1 Dalton (slightly heavier than protons), number can vary (isotopes).

• Electrons: Negatively charged, found outside the nucleus in orbitals, mass ≈ 1/2000 Dalton, number equals protons in a neutral atom.

Atomic Number: Number of protons in the nucleus (defines the element).

Atomic Mass: Sum of protons and neutrons.

• Example: Iron (Fe) has atomic number 26 (protons), atomic mass 56. Number of neutrons = 56 - 26 = 30.

Isotopes

Definition and Properties

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Some isotopes are stable, while others are radioactive.

• Radioactive isotopes can emit radiation and are used in biological research and medicine.

• Adding neutrons increases atomic mass but not the chemical properties.

Electron Orbitals and Energy Shells

Organization of Electrons

Electrons occupy specific regions called orbitals around the nucleus, organized into energy shells. The arrangement of electrons determines how atoms interact and bond.

• Orbitals: Each orbital holds up to 2 electrons.

• Energy shells:

  • 1st shell: 1 orbital = 2 electrons max

  • 2nd shell: 4 orbitals = 8 electrons max

  • 3rd shell: 4 orbitals = 8 electrons max (some elements can hold more in higher shells)

• Periodic Table: Rows = number of shells; Columns (groups) = number of electrons in the outer shell.

• Example: Calcium is in the 7th row, indicating 7 shells.

Valence Electrons

Importance in Chemical Bonding

Valence electrons are the electrons in the outermost shell of an atom. They determine the atom's chemical reactivity and bonding behavior.

• Valence shell: Outermost energy shell.

• Valency: Number of electrons needed to fill the outer shell.

• Example: Oxygen has a valency of 2 (needs 2 more electrons to fill its shell).

• Noble gases: Have full outer shells and are chemically inert.

• Octet Rule: Atoms tend to gain, lose, or share electrons to have 8 electrons in their outer shell.

Ions and Ionic Bonds

Formation and Properties

Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge. Ionic bonds form between oppositely charged ions.

• Cation: Positively charged ion (lost electrons).

• Anion: Negatively charged ion (gained electrons).

• Ionic bond: Attraction between cations and anions.

Electronegativity

Definition and Trends

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. It influences bond type and polarity.

• Oxygen is the most electronegative element in living systems.

• Fluorine is the most electronegative element overall.

• Electronegativity differences determine bond polarity.

Covalent Bonds

Types and Properties

Covalent bonds form when two atoms share one or more pairs of electrons. They are the strongest type of chemical bond in biological molecules.

• Single bond: 1 pair of shared electrons (e.g., H–H).

• Double bond: 2 pairs of shared electrons (e.g., O=O).

• Molecules with covalent bonds are often gases, liquids, or soft solids at room temperature.

Polarity

Polar vs. Non-polar Covalent Bonds

Polarity describes the distribution of electrical charge in a molecule.

• Non-polar covalent bond: Electrons are shared equally; no distinct poles (e.g., H2, O2).

• Polar covalent bond: Electrons are shared unequally, creating partial charges (e.g., H2O, H–Cl).

• Polar molecules can form hydrogen bonds and are hydrophilic.

• Non-polar molecules are hydrophobic.

Hydrogen Bonds

Definition and Biological Importance

Hydrogen bonds are weak attractions between a slightly positive hydrogen atom in one molecule and a slightly negative atom (often oxygen or nitrogen) in another molecule. They are crucial for the structure of water, DNA, and proteins.

Van der Waals Forces

Definition

Van der Waals forces are weak attractions between non-polar molecules due to temporary dipoles. They contribute to the stability of large biological molecules.

Bonding Patterns in Biological Molecules

Common Bonding Rules

Atoms form a predictable number of bonds to achieve a stable electron configuration (octet rule).

• Hydrogen (H): 1 bond

• Oxygen (O): 2 bonds

• Nitrogen (N): 3 bonds

• Carbon (C): 4 bonds

This pattern is often remembered as "HONC 1234": H = 1, O = 2, N = 3, C = 4.

Periodic Table and Electron Configuration

Understanding the Table

The periodic table arranges elements by increasing atomic number and groups them by similar chemical properties. Electron configuration diagrams show the arrangement of electrons in shells around the nucleus.

Element	Atomic Number	Electron Configuration (Shells)	Valence Electrons
Hydrogen	1	1	1
Oxygen	8	2, 6	6
Nitrogen	7	2, 5	5
Carbon	6	2, 4	4
Neon	10	2, 8	8

Example: Carbon (C) has 6 electrons: 2 in the first shell, 4 in the second shell, and thus 4 valence electrons.

Key Equations

• Atomic Mass Calculation:

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

• Number of Neutrons:

$\text{Number of Neutrons} = \text{Atomic Mass} - \text{Atomic Number}$

Summary Table: Types of Chemical Bonds

Bond Type	Definition	Relative Strength	Example
Covalent	Atoms share electron pairs	Strongest	H2O, O2
Ionic	Transfer of electrons between atoms	Strong (in dry conditions)	NaCl
Hydrogen	Attraction between polar molecules (H and O/N/F)	Weak	Between water molecules
Van der Waals	Temporary dipole attractions	Very weak	Between non-polar molecules