Ch 2 Life's Chemistry and the Importance of Water
Life's Chemistry and the Importance of Water
Fundamental Composition of Life: Living organisms are composed of the same chemical elements as the nonliving universe.
Abundant Elements: There are 92 naturally occurring chemical elements, but living organisms are largely composed of remarkably few of them.
Hydrogen (H), Oxygen (O), Carbon (C), and Nitrogen (N) account for more than 95\% of the weight of all living organisms.
The properties of these four elements are critical for life, as discussed in detail throughout the text.
KEY CONCEPT 2.1 An Element's Atomic Structure Determines Its Properties
Matter and Atoms: All matter is composed of atoms.
An element is a fundamental substance consisting of only one kind of atom. There are over 100 different elements, with 92 occurring naturally.
Elements interact differently, determined by the number and charge of particles in their atoms.
Atomic Structure: Atoms consist of three fundamental particles:
Protons: Positively charged, located in the nucleus. The number of protons is the atomic number, which determines the element's identity.
Neutrons: Uncharged (neutral), located in the nucleus. (Except for hydrogen, which typically has no neutrons).
Electrons: Negatively charged, rapidly moving outside the nucleus in electron shells.
Atoms are mostly empty space due to the vast distance between the nucleus and electron shells.
Mass: Protons and neutrons are \approx 2,000 times larger than electrons and contribute most of an atom's mass.
The unit for atomic mass is the dalton (Da).
Proton mass \approx 1\text{ Da}; Neutron mass \approx 1\text{ Da}; Electron mass \approx 0.0005\text{ Da} (negligible for atomic mass).
Periodic Table and Element Identity:
Elements are presented in order of atomic number and organized into vertical columns (groups).
Elements in the same column have the same number of electrons in their outermost shell (the valence shell), leading to similar chemical properties.
Atomic numbers of common elements in life: H=\text{1}, C=\text{6}, N=\text{7}, O=\text{8}.
Isotopes:
Atoms of an element can vary in the number of neutrons in their nucleus.
These variants are called isotopes (e.g., carbon-12 has 6 neutrons, carbon-14 has 8 neutrons).
Isotopes do not change the element's identity (determined by protons) but affect its atomic mass.
Electron Shells and Orbitals:
The Bohr model simplifies electron location into electron shells at different distances from the nucleus.
Innermost shell: max 2 electrons.
Second shell: max 8 electrons.
Third shell: max 18 electrons.
Electrons actually move within defined areas of space called atomic orbitals.
Orbitals differ in shape, orientation, and energy level.
Each orbital can hold one or two electrons.
Every electron shell has a single, spherical s orbital.
Second and subsequent shells also have three dumbbell-shaped p orbitals.
Electrons in lower-shell orbitals generally have lower energy. Within a given shell, s-orbital electrons have lower energy than p-orbital electrons.
Chemical Properties and Valence Electrons:
Electrons in the outermost shell (valence electrons) determine an element's chemical properties.
Octet Rule: Elements tend to be most stable when their s and p orbitals in the outermost shell are full (i.e., when they have eight electrons: two in the s orbital and two in each of the three p orbitals).
Elements in the last column of the periodic table (noble gases) have this stable configuration and are chemically inert.
Electronegativity:
An atom's tendency to attract electrons from another atom. It depends on:
Number of electrons in outermost s and p orbitals (more electrons = higher electronegativity).
Distance between outermost electrons and the positively charged nucleus (closer electrons = greater pull).
Electronegativity increases from bottom left to top right of the periodic table.
H, C, N, and O (most abundant in living systems) are among the most electronegative atoms.
Unstable Elements and Isotopes:
Some elements (e.g., atomic numbers 95-118) and isotopes (e.g., uranium) are highly unstable and undergo radioactive decay.
Radioactive decay involves spontaneously losing atomic particles to become different isotopes or even different elements (e.g., $^{14}\text{C}$ decays to $^{14}\text{N}$). This property is used in radiometric dating.
KEY CONCEPT 2.2 Atoms Bond to Form Molecules
Chemical Bonds: An attraction between two atoms. Bonds can be strong or weak.
Molecule: Formed when two or more atoms join chemically.
Compound: A molecule formed by at least two different elements.
Covalent Bonds:
Form when two atoms share one or more pairs of electrons to achieve stable electron configurations in their outermost shells.
Each atom contributes one electron to each shared pair.
Examples: Hydrogen molecule (H\text{2} - single bond), Oxygen molecule (O\text{2} - double bond), Water (H\text{2}O), Methane (CH\text{4}).
Represented by lines in structural formulas.
Common elements in biological molecules (H, C, N, O) have similar electronegativities and readily form covalent bonds.
Properties of Covalent Bonds:
Orientation: Consistent bond length and angles for a given pair of elements/atoms (e.g., carbon forms a tetrahedral orientation with four single bonds).
Strength and Stability: Covalent bonds are strong, requiring significant energy to break, but strength varies.
Unequal Sharing of Electrons:
Nonpolar covalent bonds: Electrons are shared equally when atoms have similar electronegativities (e.g., C-C, or C-H where difference is small).
Polar covalent bonds: Electrons are unequally shared when atoms have different electronegativities. The more electronegative atom pulls electrons closer, becoming slightly negatively charged (\delta-) and the other slightly positively charged (\delta+).
Example: O-H bond in water; oxygen is more electronegative than hydrogen.
Polarity is crucial for hydrogen bonds and biological systems.
Ionic Bonds:
Form when a highly electronegative atom pulls one or more electrons completely away from an atom with low electronegativity (a transfer of electrons).
This creates electrically charged atoms called ions (positively charged cations and negatively charged anions).
The strong electrostatic attraction between these oppositely charged ions forms the ionic bond.
Example: Sodium (Na) losing an electron to Chlorine (Cl) forms Na^\text{+} and Cl^\text{-} ions, which attract to form NaCl (table salt).
Ionic compounds often form solid crystal lattices with regular arrangements of ions.
In biological systems, ionic substances often dissolve in water, where hydration shells form around ions, reducing the strength of the ionic bond.
Weaker Electrostatic Attractions:
Dipole-dipole interactions: Attractions between slightly positively and negatively charged atoms in nearby polar bonds. Generally weak.
Hydrogen bonds: A type of dipole-dipole interaction where a slightly positively charged hydrogen atom (covalently bonded to O or N) is attracted to a slightly negatively charged O or N atom in another molecule or within the same molecule.
Individually weak but collectively strong, influencing structure of large molecules (e.g., DNA, proteins) and water properties.
Van der Waals interactions: Very weak, transient electrostatic attractions between atoms in nonpolar molecules (or regions of polar molecules).
Caused by temporary asymmetrical distribution of electrons, creating brief dipoles that induce dipoles in neighboring close-proximity atoms.
Short-lived and very weak individually, but numerous interactions can create significant forces (e.g., gecko foot adhesion).
KEY CONCEPT 2.3 Chemical Transformations Involve Energy and Energy Transfers
Energy and Change: Energy is the capacity to produce a change. All changes in the universe involve energy transformation.
Categories of Energy:
Kinetic energy: Energy of movement (thermal, sound, electromagnetic). Examples: Earth orbiting, a rolling ball, moving molecules, electrons in orbitals, photons.
Potential energy: Stored energy (gravitational, elastic, chemical-bond, nuclear). Examples: a ball held above ground, stretched tendon, covalent bond.
Energy Transformations: Energy is converted from one form to another (e.g., light to chemical-bond energy) or transferred from one location to another (e.g., food energy to molecular building energy), or both.
Laws of Thermodynamics: Govern all energy transformations.
First Law of Thermodynamics (Conservation of Energy): Energy cannot be created or destroyed. The total amount of energy in a closed system remains constant during a transformation.
Energy is always conserved, it just changes form or location.
Breaking down complex molecules (e.g., starch) releases stored energy.
Building complex molecules requires energy input.
Bond strength and potential energy: Weak bonds require little energy to break (high potential energy). Strong bonds require much energy to break (low potential energy). Breaking weak bonds to make strong ones releases chemical-bond energy.
Second Law of Thermodynamics (Increase in Entropy): With each energy transformation, there is an increase in entropy (disorder, spread-out energy, unusable energy).
After an energy transformation, the amount of usable energy declines.
Energy becomes less concentrated and less usable.
Spontaneous processes increase entropy (e.g., a hot kettle cools down).
An input of energy is required to impose order (decrease entropy).
In living systems, transformations often occur in steps, with an overall increase in entropy at each step (e.g., food chains lose usable energy).
KEY CONCEPT 2.4 Chemical Reactions Transform Substances
Chemical Reaction: Occurs when atoms combine or change their bonding partners.
Reactants: Starting substances.
Products: Substances formed.
The total number of atoms of each element remains the same, but their arrangement changes.
Energy in Chemical Reactions:
The energy in covalent bonds and the entropy differ between reactants and products.
Free energy (\Delta G): The total energy change, combining differences in bond energy and entropy.
Exergonic reactions: Release energy (\Delta G < 0).
Chemical energy in products is less than in reactants, or entropy of products is greater than reactants, or both.
Example: Hydrolysis (Complex molecule + H\text{_2}O ightarrow simpler molecules).
Breaks larger molecules into smaller ones with stronger bonds (less potential energy).
Example: Sucrose + H\text{_2}O
ightarrow Glucose + Fructose.Releases energy and increases entropy.
Endergonic reactions: Require an input of energy (\Delta G > 0).
Example: Condensation reactions (Two molecules combine to form a larger one, releasing H\text{_2}O).
Produces larger molecules with weaker bonds (more potential energy).
Example: Glucose + Fructose
ightarrow Sucrose + H\text{_2}O.Living systems synthesize large molecules in endergonic reactions by coupling them with exergonic reactions.
Activation Energy (E_a):
The total amount of energy that must be supplied for a reaction to begin, even for exergonic reactions.
Required because covalent bonds in reactants must be broken, which needs energy input.
Once bonds are broken, new, lower-energy (stronger) bonds can form, releasing energy.
Main reason why even strongly exergonic reactions don't begin spontaneously (e.g., gasoline and oxygen need a spark).
Factors Affecting Reaction Rate:
Activation energy: Lower E_a increases the rate because more collisions will have sufficient energy.
Enzymes are proteins that lower E_a in biological systems.
Temperature: Higher temperature increases both the number and energy of collisions between molecules, thus increasing the rate.
Concentration: Higher concentration of reactants increases the number of collisions, thus increasing the rate.
Reversible Reactions and Equilibrium:
Many reactions (e.g., sucrose hydrolysis) are reversible, proceeding in both forward and reverse directions.
Equilibrium: The point at which the rates of the forward and reverse reactions become equal, and the relative concentrations of reactants and products no longer change.
For sucrose hydrolysis, equilibrium favors glucose and fructose (products) because the forward reaction is exergonic.
KEY CONCEPT 2.5 The Properties of Water Are Critical to the Chemistry of Life
Importance of Water: Essential for life; search for extraterrestrial life often begins with finding liquid water.
Water's critical properties stem from the polarity of its O-H covalent bonds and its ability to form hydrogen bonds.
Temperature Moderation:
High specific heat: The amount of heat absorbed or lost to change 1\text{ gram} of a substance by 1{^\circ}\text{C}. Water has high specific heat because hydrogen bonds must be broken to increase molecular movement (temperature).
High heat of vaporization: The amount of energy required to vaporize water. Due to hydrogen bonds, it takes a lot of heat to make water evaporate.
Consequences for life:
Water exists as a liquid over a broad temperature range (up to 100{^\circ}\text{C}).
Evaporative cooling: Organisms lose heat by evaporating water (e.g., sweating).
Large amounts of water in tissues minimize temperature fluctuations from environmental changes.
Ice Density:
When water freezes, it forms additional hydrogen bonds, creating a more organized structure.
Water molecules in ice are farther apart than in liquid water, making ice less dense than liquid water.
This is crucial for aquatic life, as ice floats and lakes/oceans freeze from the top down, preventing solid freezing and allowing survival.
Cohesion and Adhesion: Due to numerous hydrogen bonds.
Cohesion: Attraction of identical molecules for one another (water-water).
Generates surface tension (e.g., water droplets, insects walking on water).
Adhesion: Attraction of different molecules for one another (water-other polar surfaces).
Role in plants: Cohesion and adhesion allow narrow columns of water to move from roots to leaves, resisting gravity.
Water as a Solvent (Universal Solvent): Many substances dissolve in water.
Dissolving: Involves distributing solute molecules throughout the solvent, not forming/breaking covalent bonds.
Hydrophilic (water-loving): Substances that dissolve easily in water, typically polar or charged (ionic).
Water molecules surround ions in hydration shells, orienting their polar ends to the opposite charges of the ions, disrupting the ionic bonds in the crystal.
Hydrophobic (water-fearing): Substances that do not dissolve in water, typically nonpolar (e.g., oil).
Interaction is energetically unfavorable; nonpolar substances aggregate to minimize disruption of water's hydrogen bonds.
Amphipathic molecules: Have both hydrophobic and hydrophilic regions (e.g., soaps).
Acidity, Basicity, and pH:
Self-ionization of pure water: A very small fraction of water molecules spontaneously ionize: 2\text{ H}2\text{O} \rightleftharpoons \text{H}3\text{O}^+ + \text{OH}^-.
H^\text{+} (hydrogen ion/proton) and H\text{_3}O^\text{+} (hydronium ion) are often used interchangeably.
In pure water, [\text{H}_3\text{O}^+] = [\text{OH}^-] = 1 \times 10^{-7}\text{ moles/liter}. This is a neutral solution.
Acids: Compounds that raise the concentration of H\text{+} (H\text{_3}O^\text{+}) ions when dissolved in water ($\text{pH} < 7$).
Bases: Compounds that produce proton acceptors (like OH^\text{-} ions) or reduce H\text{+} concentration when dissolved in water ($\text{pH} > 7$).
pH: The negative logarithm of the H\text{3}O^\text{+} ion concentration (\text{pH} = -\log{10}[\text{H}_3\text{O}^+]).
Buffers: Solutions that resist changes in pH when acids or bases are added. They consist of a weak acid and its conjugate base. (e.g., the carbonic acid-bicarbonate system in blood maintains pH between 7.35-7.45).
Environmental impact: CO\text{_2} dissolving in water forms carbonic acid, contributing to acid rain and ocean acidification, which harms marine organisms with calcium carbonate skeletons.
KEY CONCEPT 2.6 Functional Groups Give Molecules Specific Properties
Biomolecules: Carbohydrates, lipids, proteins, and nucleic acids. These are carbon-based molecules (organic chemistry).
Carbon's Role: Carbon atoms form four covalent bonds, allowing for structurally complex molecules.
Functional Groups: Small clusters of atoms attached to larger organic molecules that confer specific chemical properties and reactivity.
Methyl (-CH\text{_3}): Nonpolar; important in protein modification and cytosine.
Hydroxyl (-OH): Polar; involved in hydrogen bonding, condensation reactions, and protein phosphorylation.
Sulfhydryl (-SH): Weakly polar; can form disulfide bridges to stabilize protein structure.
Aldehyde (-CHO): Polar; very reactive; important in energy-releasing reactions.
Keto (C=O): Polar; important in carbohydrates and energy reactions.
Carboxyl (-COOH): Charged, acidic; ionizes to -COO^\text{-} and H\text{+} in living tissues; reacts with amino groups to form peptide bonds.
Amino (-NH\text{2}): Charged, basic; accepts H\text{+} to form -NH\text{3}^\text{+}; reacts with carboxyl groups to form peptide bonds.
Phosphate (-OPO\text{3}H\text{2}): Charged, acidic; ionizes to -OPO\text{_3}}^\text{2-}; enters condensation reactions, and its hydrolysis when bonded to another phosphate is strongly exergonic.
Isomers (Molecular Diversity): Molecules with the same chemical formula but different arrangements of atoms.
Structural isomers: Differ in how atoms and functional groups are bonded to one another (e.g., glucose and fructose, both C\text{6}H\text{12}O\text{_6}).
Stereoisomers: Same connections but differ in three-dimensional geometry.
Properties of H, O, C, and N in Living Tissues
Relative Abundance: These four elements (H, O, C, N) are among the most common elements in the universe. (Helium and Neon, though abundant, are inert gases).
Chemical Versatility:
They can form both nonpolar and polar covalent bonds, which allow for a wide range of molecular structures and interactions.
They readily participate in hydrogen bonds, critical for water's properties and macromolecular structure.
They can form ions, especially within functional groups, enabling diverse chemical reactions.
Structural Complexity: Carbon, in particular, forms four covalent bonds, leading to large, structurally complex molecules (e.g., carbon backbones) essential for life.
Importance of Water: The commonness of Hydrogen and Oxygen is largely due to life evolving in and being substantially composed of water. Water's unique properties (high specific heat, solvent capabilities, etc.) make it the ideal medium for life's chemical reactions.