Overview
This lecture covers the chemical context of life, focusing on the elements, atomic structure, types of chemical bonds, and chemical reactions that are foundational to biology. Understanding chemistry is essential for understanding biological processes, as all living things are composed of chemicals and their interactions.
Matter, Elements, and Compounds
- All living organisms are made of matter, which has mass and occupies space.
- Matter exists in different forms: solid, liquid, gas, and plasma.
- Elements are pure substances that cannot be broken down into other substances by chemical reactions.
- The six most common elements in living systems are carbon, hydrogen, nitrogen, oxygen, phosphorus, and sulfur (CHNOPS). These elements are essential for life and make up the majority of living matter.
- A mnemonic for remembering these elements is "CHNOPS."
- Compounds are substances composed of two or more elements in a fixed ratio (e.g., H₂O, NaCl, C₆H₁₂O₆).
- The properties of compounds are different from those of the individual elements that form them. For example, sodium is a reactive metal and chlorine is a poisonous gas, but together they form table salt (NaCl), which is safe to eat.
- Compounds are found in living systems in many forms, such as water, glucose, and carbon dioxide.
Essential and Trace Elements
- Of the 92 naturally occurring elements, only about 20–25 are essential for life.
- The CHNOPS elements account for about 96% of living matter.
- Most of the remaining 4% consists of elements like calcium and potassium, with phosphorus and sulfur being less abundant but still important.
- Trace elements are required in very small amounts but are still vital for life.
- Examples include magnesium, chlorine, boron, chromium, cobalt, copper, and iodine.
- Iodine is necessary for thyroid hormone production, and its deficiency can lead to health problems.
- Trace elements can be found in the human body and are important for various biological functions.
Atomic Structure
- Atoms are the smallest units of matter that retain the properties of an element.
- Each atom consists of three main subatomic particles:
- Protons (positive charge, p⁺)
- Neutrons (no charge, n⁰)
- Electrons (negative charge, e⁻)
- Protons and neutrons are located in the atomic nucleus at the center of the atom, while electrons form a cloud around the nucleus.
- The mass of protons and neutrons is measured in Daltons (Da); electrons have much less mass and are usually not counted in atomic mass.
- Atomic number = number of protons in the nucleus (defines the element).
- Changing the number of protons changes the element.
- Mass number = total number of protons and neutrons in the nucleus.
- The actual atomic mass is close to the mass number, as electrons contribute very little mass.
- Atoms of each element have a unique number of protons, but can have different numbers of neutrons or electrons.
Isotopes and Ions
- Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.
- All isotopes of an element have the same number of protons.
- Some isotopes are stable, while others are radioactive and decay spontaneously, emitting particles and energy.
- Radioactive isotopes can be used as tracers in medical imaging or for radiometric dating to determine the age of fossils.
- The rate of decay is measured by half-life, the time it takes for half of the isotope to decay.
- Ions are atoms or molecules with unequal numbers of protons and electrons, resulting in a net charge.
- Cations are positively charged ions (have lost electrons).
- Anions are negatively charged ions (have gained electrons).
- The transfer of electrons creates these charged species, which are important in ionic bonding.
- Ions are common in biological systems and play key roles in processes like nerve signaling and muscle contraction.
Energy and Electron Arrangement
- Energy is the capacity to cause change; potential energy is energy stored due to an object's position or structure.
- Electrons in atoms have different amounts of potential energy depending on their distance from the nucleus.
- Electrons occupy energy levels or shells; the first shell holds up to 2 electrons, the second up to 8, and so on.
- Electrons fill the lowest available energy levels first (closest to the nucleus), similar to filling seats closest to the stage at a concert.
- Valence electrons are those in the outermost shell and are most important for chemical bonding and reactivity.
- Atoms with full valence shells (such as noble gases) are chemically inert and do not readily react with other atoms.
- The number of valence electrons determines how an atom will interact with others.
- Electron shells are made up of orbitals, which are three-dimensional spaces where electrons are likely to be found.
- The first shell has one spherical (s) orbital; the second shell has one s orbital and three p orbitals (figure-eight shaped).
- Each orbital can hold up to two electrons.
- Orbitals can hybridize, affecting the shape and function of molecules.
- The arrangement of electrons in shells and orbitals influences the chemical behavior of atoms.
Chemical Bonds
- Chemical bonds form when atoms interact to achieve full valence shells, either by sharing or transferring electrons.
- Covalent bonds involve the sharing of electron pairs between atoms.
- Single, double, or triple covalent bonds can form depending on the number of shared electron pairs.
- The shared electrons count toward the valence shells of both atoms.
- Molecules are formed when two or more atoms are held together by covalent bonds.
- Electronegativity is an atom's tendency to attract shared electrons in a bond.
- Atoms with higher electronegativity (e.g., oxygen, nitrogen, fluorine) pull electrons more strongly.
- Polar covalent bonds occur when electrons are shared unequally, resulting in partial charges (e.g., water, H₂O).
- The more electronegative atom gains a partial negative charge, while the less electronegative atom becomes partially positive.
- Nonpolar covalent bonds occur when electrons are shared equally (e.g., O₂).
- Ionic bonds form when electrons are transferred from one atom to another, creating oppositely charged ions (cations and anions) that attract each other (e.g., NaCl).
- Cations are positive (lost electrons); anions are negative (gained electrons).
- Ionic compounds often form crystalline structures, such as table salt.
- Hydrogen bonds are weak attractions between a hydrogen atom covalently bonded to an electronegative atom (usually O or N) and another electronegative atom.
- These are important in biological molecules like water and proteins.
- Hydrogen bonds are shown as dotted lines and are easily broken, allowing for flexibility in biological structures.
- Van der Waals interactions are weak, temporary attractions between molecules due to transient charge differences.
- Individually weak, but collectively can be strong (e.g., gecko feet sticking to surfaces).
- These interactions occur when electrons are distributed unevenly in molecules that are close together.
Molecular Shape and Function
- The shape of a molecule is determined by the arrangement of its atoms and the repulsion between electron pairs in orbitals.
- Hybridization of orbitals can create specific molecular shapes (e.g., tetrahedral).
- The shape of a molecule affects its function in biological systems ("form defines function").
- The arrangement of electron pairs can cause molecules to bend or take on specific three-dimensional shapes.
- Molecules with similar shapes can mimic each other's effects in the body.
- For example, morphine mimics the shape of endorphins and can bind to the same receptors in the brain, producing similar effects.
- Other drugs work by mimicking or blocking natural molecules due to their shape.
- The ability of molecules to bind to specific receptors is determined by their shape and size.
Chemical Reactions
- Chemical reactions involve the making and breaking of chemical bonds, resulting in the transformation of reactants into products.
- Reactants are the starting substances; products are the substances formed.
- In chemical equations, reactants are written on the left, products on the right, with an arrow indicating the direction of the reaction ("yields").
- Some reactions are reversible, meaning they can proceed in both directions.
- Reversible reactions are indicated by a double arrow (⇌).
- Chemical equilibrium is reached when the forward and reverse reaction rates are equal, and the concentrations of reactants and products remain constant.
- Examples include photosynthesis (CO₂ + H₂O ⇌ C₆H₁₂O₆ + O₂) and the formation/breakdown of ammonia (N₂ + H₂ ⇌ NH₃).
- Chemical reactions are fundamental to all biological processes, including metabolism, energy transfer, and growth.
Key Terms & Definitions
- Element: A substance that cannot be broken down by chemical reactions.
- Compound: A substance composed of two or more elements in a fixed ratio.
- Atom: The smallest unit of matter retaining the properties of an element.
- Isotope: Atoms of the same element with different numbers of neutrons.
- Ion: An atom or molecule with a net electric charge due to unequal numbers of protons and electrons.
- Cation: A positively charged ion (lost electrons).
- Anion: A negatively charged ion (gained electrons).
- Covalent bond: A chemical bond formed by the sharing of electron pairs between atoms.
- Ionic bond: A chemical bond formed by the attraction between oppositely charged ions.
- Electronegativity: An atom's ability to attract shared electrons in a bond.
- Hydrogen bond: A weak attraction between a hydrogen atom and an electronegative atom (O or N).
- Van der Waals interaction: Weak, temporary attractions between molecules due to transient charge differences.
- Valence electron: An electron in the outermost shell, involved in chemical bonding.
- Chemical equilibrium: The state in which the forward and reverse reaction rates are equal, and concentrations of reactants and products remain constant.
- Potential energy: Energy stored due to position or structure.
- Radioactive isotope: An isotope that decays spontaneously, emitting particles and energy.
- Half-life: The time required for half of a radioactive isotope to decay.
Action Items / Next Steps
- Review atomic structure, subatomic particles, and electron configuration from previous chemistry courses.
- Study the different types of chemical bonds (covalent, ionic, hydrogen, van der Waals) and be able to identify examples of each.
- Practice determining the number of protons, neutrons, and electrons in isotopes and ions.
- Be prepared to discuss examples of chemical reactions, including reversible reactions and chemical equilibrium, in class.
- Understand how molecular shape influences biological function and be able to explain examples such as morphine and endorphins.
- Familiarize yourself with the key terms and definitions to build a strong foundation for upcoming chapters.