Chemistry and Structure of the Cell

Introduction

The chemistry underlying biological systems is essential for understanding cell structure and function. This section covers atomic structure, elements, isotopes, electron configuration, and chemical bonding, all foundational for cell biology.

Atoms and Subatomic Particles

Definition and Structure of Atoms

• Atom: The smallest unit of an element that retains its chemical properties.

• Subatomic particles:

  • Electron (e-): Negatively charged

  • Proton (p+): Positively charged

  • Neutron (n): No charge (neutral)

• Atoms consist of a nucleus (protons and neutrons) and an electron cloud (electrons).

Example: A hydrogen atom has one proton and one electron.

Elements and Atomic Number

Properties of Elements

• Element: A substance that cannot be broken down by ordinary chemical means.

• Elements are organized in the periodic table based on atomic number.

• Atomic Number: Number of protons in the nucleus.

• Atomic Mass: Sum of protons and neutrons in the nucleus.

Example: Carbon (C) has atomic number 6 and atomic mass approximately 12.

Isotopes

Definition and Examples

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Most elements exist as mixtures of isotopes.

Example: Carbon-14 is radioactive and used in radiometric dating.

Electrons and Chemical Properties

Electron Orbitals and Valence Electrons

• Electrons are found in regions called orbitals around the nucleus.

• Each orbital can contain up to two electrons.

• Valence electrons are those in the outermost shell and determine chemical reactivity.

Example: Oxygen has six valence electrons, making it highly reactive.

Electron Energy Levels

Energy and Electron Shells

• Energy: The capacity to do work or cause change.

• Potential energy: Energy due to position; electrons farther from the nucleus have higher potential energy.

• Electrons occupy energy levels or shells around the nucleus.

Example: Neon has two filled shells (10 electrons).

Additional info: The energy of an electron in a shell is quantized and can be calculated using Planck's equation above, where is energy, is Planck's constant, and is frequency.

Chemical Bonds

Types of Chemical Bonds

• Chemical bond: An attraction between atoms that holds them together in molecules or compounds.

• Ionic bonds: Formed when atoms transfer electrons, resulting in charged ions.

• Covalent bonds: Formed when atoms share electron pairs.

• Hydrogen bonds: Weak attractions between polar molecules, especially involving hydrogen.

Example: Water (H2O) molecules are held together by covalent bonds, and interact via hydrogen bonds.

Ionic Bonds

Formation and Properties

• Atoms lose or gain electrons to achieve a full outer shell (octet rule).

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Ionic bonds form between cations and anions due to electrostatic attraction.

Example: Sodium chloride (NaCl) forms when sodium donates an electron to chlorine.

Covalent Bonds

Single, Double, and Triple Bonds

• Single covalent bond: Two atoms share one pair of electrons.

• Double covalent bond: Two atoms share two pairs of electrons.

• Triple covalent bond: Two atoms share three pairs of electrons.

• Bond strength increases with the number of shared electron pairs.

Example: Oxygen gas (O2) has a double bond; nitrogen gas (N2) has a triple bond.

Polarity and Electronegativity

Polar and Nonpolar Covalent Bonds

• Electronegativity: An atom's tendency to attract electrons in a bond.

• Nonpolar covalent bond: Electrons are shared equally (similar electronegativity).

• Polar covalent bond: Electrons are shared unequally (different electronegativity).

Example: Water is a polar molecule due to the difference in electronegativity between oxygen and hydrogen.

Hydrogen Bonds

Importance in Biology

• Hydrogen bonds: Weak attractions between a hydrogen atom in one molecule and an electronegative atom in another.

• Hydrogen bonds are crucial for the structure of water, DNA, and proteins.

• They form between polar molecules and contribute to molecular stability.

Example: Hydrogen bonds hold the two strands of DNA together.