The Chemical Context of Life

Atoms: The Smallest Units of Matter

All living organisms are composed of matter, which consists of elements organized into atoms. Understanding atomic structure is fundamental to biology, as it underpins the chemical properties of life.

• Atom: The smallest unit of an element, retaining its chemical properties.

• Element: A substance that cannot be broken down into simpler substances by chemical means.

• Compound: A substance formed when two or more elements are chemically bonded.

• Example: Diamond (pure carbon) and table salt (sodium chloride) are composed of atoms arranged in specific ways.

Atomic Structure

Atoms consist of subatomic particles: protons, neutrons, and electrons. Their arrangement determines the atom's identity and properties.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle found in the nucleus.

• Electron: Negatively charged particle orbiting the nucleus.

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

• Mass Number (A): Total number of protons and neutrons.

• Example: Carbon has 6 protons, 6 neutrons, and 6 electrons.

Elements Essential to Life

Life depends on a small subset of elements, each with unique chemical properties.

• Major Elements: Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N), Phosphorus (P), Sulfur (S).

• Trace Elements: Required in minute amounts (e.g., iron, iodine).

• Example: Water (H2O) is composed of hydrogen and oxygen.

The Periodic Table

The periodic table organizes elements by atomic number and chemical properties.

• Groups: Columns with elements sharing similar properties.

• Periods: Rows indicating increasing atomic number.

• Example: Group 1 elements (alkali metals) are highly reactive.

Electron Orbitals & Energy Shells

Electrons occupy energy levels (shells) around the nucleus. The arrangement of electrons determines chemical reactivity.

• First shell: Holds up to 2 electrons.

• Second shell: Holds up to 8 electrons.

• Valence electrons: Electrons in the outermost shell; involved in chemical bonding.

• Example: Carbon has 4 valence electrons.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

• Stable Isotopes: Do not decay over time.

• Radioactive Isotopes: Decay spontaneously, emitting particles and energy.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

• Application: Radioactive isotopes are used in radiometric dating and medical imaging.

Introduction to Chemical Bonding

Chemical bonds form when atoms share or transfer electrons, resulting in molecules and compounds essential for life.

• Molecule: Two or more atoms held together by covalent bonds.

• Compound: Substance formed from two or more different elements.

• Chemical Formula: Notation showing the elements and their ratios in a compound (e.g., H2O).

Types of Chemical Bonds

Atoms interact through various types of chemical bonds, each with distinct properties.

• Covalent Bonds: Atoms share pairs of electrons.

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

• Hydrogen Bonds: Weak attractions between a hydrogen atom and an electronegative atom (e.g., oxygen, nitrogen).

• Van der Waals Interactions: Weak, transient attractions between molecules due to temporary charge differences.

Covalent Bonds: Polar and Nonpolar

Covalent bonds can be classified based on the sharing of electrons and the electronegativity of the atoms involved.

• Nonpolar Covalent Bonds: Electrons are shared equally between atoms of similar electronegativity (e.g., H2).

• Polar Covalent Bonds: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

• Electronegativity: A measure of an atom's ability to attract electrons in a bond.

• Example: Water is a polar molecule due to unequal sharing of electrons between hydrogen and oxygen.

Ionic Bonds

Ionic bonds form when one atom donates an electron to another, creating oppositely charged ions that attract each other.

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

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

• Example: Sodium (Na) donates an electron to chlorine (Cl), forming Na+ and Cl- ions in NaCl.

Intermolecular Bonds

Intermolecular bonds are weak interactions between molecules, crucial for biological processes such as protein folding and DNA structure.

• Hydrogen Bonds: Important in water, DNA, and proteins.

• Van der Waals Interactions: Stabilize molecular structures.

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

Radioactive Decay and Biological Applications

Radioactive isotopes decay over time, releasing energy and particles. This property is used in biological research and medicine.

• Radiometric Dating: Determines the age of fossils and rocks using isotopes like Carbon-14.

• Medical Imaging: Radioisotopes are used to trace biological processes.

• Equation: , where is the decay constant.

Summary Table: Key Elements in Biology

Practice & Application

• Identify the number of protons, neutrons, and electrons in an atom given its atomic and mass numbers.

• Predict the type of bond formed between two atoms based on their electronegativity.

• Explain the biological significance of water's polarity and hydrogen bonding.

• Use isotopic data to estimate the age of biological samples.

Additional info: These notes expand upon the brief points and diagrams in the original study prep slides, providing definitions, examples, and academic context suitable for college-level General Biology students.