Building Blocks of Molecules

Introduction to Biochemistry

Biochemistry is the study of the chemistry of life, focusing on the molecules and reactions that occur within biological systems. Understanding these systems requires breaking them down into their component parts, studying these parts in isolation, and then reconstructing the system to see how the parts interact.

• Key biological molecules: Proteins, nucleic acids, carbohydrates, and lipids

• Energetics of life: Metabolism and free energy

Definitions and Fundamental Concepts

The study of chemistry begins with understanding the basic building blocks of matter.

• Matter: Physical material of the universe; anything that has mass and occupies space

• Mass: Amount of matter present

• Pure substance: Uniform chemical composition

• Mixture: Variable composition; can be physically separated

• Molecule: Smallest unit of a pure substance that retains its properties; made up of atoms

• Atom: Smallest particle of matter (element)

• Element: Homogeneous, pure substance, only one kind of atom

• Compound: Homogeneous, pure substance, multiple kinds of atoms

Examples of Elements and Compounds

• Water (H2O): Compound of 2 hydrogen atoms and 1 oxygen atom

• Table salt (NaCl): Compound of 1 sodium atom and 1 chlorine atom

• Glucose (C6H12O6): Compound of 6 carbon, 12 hydrogen, and 6 oxygen atoms

Elements in Biological Systems

The structure of biological molecules is determined by the elements that compose them.

• Most common elements in cells: Carbon (C), Nitrogen (N), Oxygen (O), Hydrogen (H)

• Structure determines function: The arrangement of atoms in molecules affects their properties and roles in biological systems

Abundance of Elements in the Human Body

The human body is composed primarily of a few key elements.

• Carbon (C): 61.7% dry weight

• Nitrogen (N): 11.0%

• Oxygen (O): 9.3%

• Hydrogen (H): 5.7%

• Calcium (Ca): 5.0%

• Phosphorus (P): 3.3%

• Potassium (K): 1.0%

• Sulfur (S): 1.0%

• Chlorine (Cl): 0.7%

• Sodium (Na): 0.7%

• Magnesium (Mg): 0.3%

Representations of Molecules

Chemists use several models to represent molecules, each providing different information about structure and properties.

• Molecular formula: Shows the types and numbers of atoms (e.g., H2O)

• Structural formula: Shows how atoms are connected

• Ball-and-stick model: Shows atoms as colored balls and bonds as sticks

• Space-filling model: Shows relative sizes of atoms in a molecule

Summary

• Elements are made of atoms that can combine to form molecules or compounds.

• The most prevalent elements in the human body are carbon (C), nitrogen (N), oxygen (O), and hydrogen (H).

Subatomic Particles and Isotopes

Atomic Structure

Atoms are composed of subatomic particles: protons, neutrons, and electrons.

• Proton: Positively charged, located in the nucleus

• Neutron: No charge, located in the nucleus

• Electron: Negatively charged, located in the electron cloud around the nucleus

Definitions

• Neutral atom: Equal numbers of protons and electrons

• Atomic number (Z): Number of protons in an element

• Isotopes: Atoms with the same number of protons but different numbers of neutrons

• Mass number (A): Number of protons + neutrons

• Atomic weight: Average mass of an element, based on isotope abundance

Hydrogen Isotopes

• Hydrogen-1 (protium): 1 proton, 0 neutrons

• Hydrogen-2 (deuterium): 1 proton, 1 neutron

• Hydrogen-3 (tritium): 1 proton, 2 neutrons

Calculating Atomic and Molecular Weights

• Atomic weight: Weighted average of isotopes

• Molecular weight: Sum of atomic weights for all atoms in a molecule

Example: Water (H2O): Glucose (C6H12O6):

Summary

• An atom is made of subatomic particles (protons, electrons, and neutrons), and its properties are defined primarily by its atomic number.

• A neutral atom has an equal number of protons and electrons.

• Isotopes are atoms with the same number of protons but different numbers of neutrons.

• The atomic weight of an element depends on the relative abundance of its isotopes.

Electronic Structure and the Periodic Table

Periodic Table Organization

The periodic table is organized by atomic number and electron configuration.

• Periods: Horizontal rows, correspond to principal energy shells (n)

• Groups/Families: Vertical columns, correspond to subshells and number of valence electrons

Electronic Structure

• Electrons are arranged in energy shells around the nucleus.

• Each shell can hold a specific number of electrons.

• Electrons in higher shells have higher energy and are further from the nucleus.

Valence Electrons

• Valence shell: Outermost electron shell

• Valence electron: Electron in the valence shell

• For main-group elements, the group number tells you the number of valence electrons (up to 8: s2 + p6)

Summary

• The periodic table is organized by the location of electrons in shells and subshells.

• The number of valence electrons determines the chemical properties and reactivity of elements.

Periodic Properties

Major Periodic Properties

• Metallic character: Ability to conduct heat/electricity, malleability, luster

• Atomic size: Increases down a group, decreases across a period

• Ionization energy: Energy required to remove an electron; increases across a period, decreases down a group

• Electronegativity: Tendency to attract electrons; increases across a period, decreases down a group

Summary

• Periodic trends include metallic character, atomic size, ionization energy, and electronegativity.

• Ionization energy and electronegativity define how hard it is for atoms to lose electrons.

The Octet Rule and Bonding

Noble Gases and the Octet Rule

• Noble gases: Group 18, have a full valence shell (8 electrons), are stable and unreactive

• Octet rule: Atoms are most stable when they have 8 valence electrons

• Atoms form bonds to achieve a full octet

Types of Chemical Bonds

• Ionic bonding: Atoms transfer electrons to form ions, which are held together by electrostatic attraction

• Covalent bonding: Atoms share electrons to achieve a full octet

Lewis Symbols

Lewis symbols use dots to represent valence electrons around an element's symbol.

Summary

• Noble gases have a full valence shell and are stable.

• The octet rule explains why atoms form bonds.

• Atoms can achieve stability by forming ionic or covalent bonds.

Ions and Ionic Bonding

Ion Formation

• Simple ion: Element with unequal numbers of protons and electrons

• Cation: Ion with more protons than electrons (positive charge)

• Anion: Ion with more electrons than protons (negative charge)

Polyatomic Ions

Polyatomic ions are covalently bonded groups of atoms with an overall charge.

• Examples: NH4+ (ammonium), H3O+ (hydronium), OH- (hydroxide), HCO3- (bicarbonate), CO32- (carbonate), SO42- (sulfate), PO43- (phosphate), CN- (cyanide)

Ionic Compounds

Ionic compounds are formed by the electrostatic attraction between cations and anions. The formula represents the lowest ratio of ions needed to form a neutral compound.

• Examples: NaCl, MgF2, K2O, Na2CO3, Mg(OH)2

Summary

• Ions form when atoms lose or gain electrons to reach an octet.

• Ionic bonds are formed between positive and negative ions.

• Common biological ions include ammonium, hydronium, hydroxide, bicarbonate, carbonate, sulfate, and phosphate.