Chemical Basis of Life: Chapter 2

• Chapter 2 explores water, carbon, and the chemical basis of life.
• Learning objectives include understanding how and why atoms interact to form molecules, listing water's unique properties, defining energy, and explaining carbon's importance.

Chemical Evolution

• Life began when simple chemical compounds combined to form larger, complex substances.
• Examples of these compounds include CO2, H2, and N2 which formed H2CO (formaldehyde) and HCN (cyanide).
• Kinetic energy converted from sunlight and heat became chemical energy in bonds.
• A complex, self-replicating compound developed.
• The self-replicating molecule became surrounded by a membrane.
• The leading hypothesis for the origin of life centres on chemical evolution and the combination of atoms.

Atoms and Elements

• Four atoms (hydrogen, carbon, nitrogen, and oxygen) make up 96% of matter in organisms.
• It's important to understand the physical structure of these atoms to understand how simple substances evolve into complex structures.
• The structure of water, carbon dioxide, and other simple molecules served as the building blocks of chemical evolution.
• Protons are positively charged (+1) particles, neutrons are neutral, and electrons are negatively charged (-1).
• Protons and neutrons reside in the nucleus.
• Electrons are located in orbitals surrounding the nucleus.
• The atomic number is the number of protons in an atom's nucleus, shown as a subscript to the left of the element symbol Atoms with the same atomic number have the same chemical properties, defining the element.
• The mass number is the total of protons and neutrons in an atom, written as a superscript to the left of the symbol.
• Protons and neutrons each have a mass of one Dalton (Da).
• Electrons have negligible mass.
• Isotopes contain different numbers of neutrons.
• Carbon, for example, can have 6, 7, or 8 neutrons; its mass number can therefore be 12, 13, or 14.
• Electrons move around the atomic nucleus in orbitals.
• Orbitals group into electron shells.
• Innermost shells are filled before the outer ones.
• The outermost shell is the valence shell.
• Electrons in the valence shell are valence electrons.
• The number of unpaired valence electrons determines the valence of an atom
• Different atoms exhibit different valences.

Covalent Bonding

• Atoms achieve stability by filling their valence shells through chemical bonds.
• Attractions that bind atoms together are chemical bonds.
• Covalent bonds form through the sharing of unpaired valence electrons between two atoms, effectively filling the outer shell of each.
• Molecules consist of substances held together by covalent bonds.
• Compounds are molecules where atoms of different elements are held together.
• Electrons aren't always shared evenly; the strength with which an atom pulls electrons toward itself is its electronegativity.
• Electronegativity is determined by the number of protons and the distance of the valence shell from the nucleus.
• Electronegativity generally increases moving up and to the right on the periodic table (O > N > S, C, H, P).
• Atoms with high electronegativity hold electrons more tightly, resulting in a partial negative charge, while the other atom has a partial positive charge.
• Oxygen is one of the most electronegative elements; it attracts electrons very strongly.
• Differences in electronegativity affect how electrons are distributed in covalent bonds.
• Nonpolar covalent bonds involve electrons shared evenly between two atoms.
• Example of a nonpolar covalent bond: C–H bond.
• Polar covalent bonds have electrons shared unevenly.
• Example of a polar covalent bond: O–H bond.
• Atoms can form multiple single, double, or triple bonds if they have more than one unpaired electron.

Ionic Bonding

• Ionic bonds result when electrons transfer from one atom to another, giving both atoms full valence shells.
• An ion is an atom or molecule with an electrical charge.
• Cations become positively charged from the loss of electrons.
• Anions become negatively charged from the gain of electrons.
• Sodium is a cation.
• Chloride is an anion.
• The degree of electron sharing in chemical bonds forms a continuum, ranging from equal sharing in nonpolar covalent bonds to the complete transfer of electrons in ionic bonds.
• Molecular formula indicates the numbers and types of atoms in a molecule.
• Examples: H2O, CH4.
• Structural formulas indicate which atoms are bonded and whether the bonds are single, double, or triple.
• The shape of a simple molecule depends on the geometry of its bonds.
• A molecule’s shape dictates its behaviour.
• Nitrogen (N2) and carbon dioxide (CO2) are linear.
• Methane (CH4) is a tetrahedron because electrons repel each other and try to occupy as much space as possible.
• Water (H2O) is planar and bent because of the two unshared electron pairs.

Water

• A great solvent, with NaCl dissolving in H2O.
• Exhibits cohesion, adhesion, and surface tension because of binding between water molecules.
• Density: Water becomes lighter upon becoming solid, resulting in ice floating.
• High capacity for absorbing energy.
• The ability to dissociate into H+ and OH-.
• Chemical evolution likely occurred in an aqueous environment
• Life is based on water due to it being an excellent solvent.
• A solute dissolves into a solvent, making a solution.
• Substances are more likely to react in a solvent.
• The oxygen atoms in water are partially negative and the hydrogen atoms are partially positive; charges are at opposite ends of the molecule, creating a polar molecule
• Water molecules interact with each other because partial negative charges on oxygen attract the partial positive charges on hydrogen.
• Weak electrical attractions between water molecules are called hydrogen bonds.
• Hydrogen bonds also form with other polar molecules.
• Hydrophilic atoms and molecules ("water-loving) are ions and polar molecules that stay in solution; they interact with water's partial charges.
• Hydrogen bonding allows almost any polar molecule to dissolve in water.
• Hydrophobic molecules ("water-fearing") are uncharged, nonpolar compounds that do not dissolve in water.
• Oil is an example of a hydrophobic molecule.
• Water has several remarkable properties, mainly due to its ability to form hydrogen bonds, including being cohesive and adhesive.
• Cohesion is the binding between like molecules (water binds to itself by hydrogen bonding).
• Adhesion is the binding between unlike molecules (water binds to plastic or glass)
• Most substances shrink when they solidify, but water expands as it freezes; it becomes less dense.
• Ice floats at the surface because water forms a relatively open crystal structure when it freezes.
• Ice forms an insulating "blanket" on water surfaces.
• break H-bonds upon heating, molecules move closer together
• High capacity to absorb heat is known as specific heat.
• A lot of energy is needed to raise the temperature of 1 gram of H2O by 1°C because H-bonds must be broken first.
• A lot of energy breaks H-bonds to convert liquid H2O to gas H2O.
• The heat of vaporisation is the energy required to change 1 gram of liquid H2O to gas.
• Sweat evaporates and takes heat from the body.
• Providing protection from the Sun's energy is a property of water, which would otherwise break apart complex molecules in oceans.
• Water molecules dissociate into a hydrogen ion (H+) and a hydroxide ion (OH-)
• The reaction, H2O ⇆ H+ + OH- , happens in both directions at the same rate, creating chemical equilibrium.
• Since protons (H+) don't exist by themselves, the reaction produces hydronium ions (H3O+): 2H2O ⇆ H3O+ + 2OH-.
• Acids release protons during chemical reactions and raise the hydronium ion concentration [H3O+].
• Adding an acid to a solution increases the proton concentration.
• Bases acquire protons during chemical reactions and lower [H3O+].
• Adding a base to a solution decreases the proton concentration.
• The pH scale expresses proton concentration [H+] in a solution on a logarithmic scale.
• The equation: pH = - log [H+].
• The pH of water is 7.
• Acids have a pH of less than 7.
• Bases have a pH of greater than 7.
• Buffers reduce pH changes, which helps maintain homeostasis, which is relatively constant conditions in living systems
• Chemical evolution theory states that simple molecules reacted to make larger, more complex ones.
• This may have been in the atmosphere or deep-sea vents.
• Chemical reactions occur when a substance is combined with or broken down into another substance.
• Chemical reactions feature reactants and products.
• Example: CO2(g) + H2O(l) → H2H2CO3(aq)
• Chemical equilibrium is when forward and reverse reactions occur at the same rate, leaving reactant and product quantities steady.
• Equilibrium is disturbed by changes to reactants, products, or heat.
• Endothermic reactions require heat to proceed.
• Exothermic reactions release heat.
• Energy is the capacity to do work or supply heat, existing as either potential or kinetic.
• Potential energy is stored.
• Kinetic energy is active energy of movement.
• Potential energy in molecules relates to the position of shared electrons in covalent bonds.
• Long, weak bonds have electrons far away from the atoms' nuclei.
• Short, strong bonds have electrons shifted closer to one or both nuclei.
• Electrons in outer shells contain more potential energy than electrons in inner shells.
• Chemical energy describes a molecule's potential to form stronger bonds.
• Thermal energy involves the kinetic energy through molecular motion.
• Molecules are constantly in motion.
• Temperature measures thermal energy in a molecule.
• Low temperatures mean molecules are moving slowly, which we perceive as "cold".
• High temperatures mean molecules are moving rapidly, which we perceive as "hot".
• Heat measures thermal energy being transferred between two objects.
• The first law of thermodynamics states that energy is conserved and cannot be created or destroyed, only transferred or transformed.

Chemical Reactions

• Chemical reaction spontaneity requires they proceed without external influence and without added energy.
• Spontaneity has two factors: Products are less ordered than reactants and entropy (disorder) increases.
• Reactions are spontaneous when the products have lower potential energy than the reactants when shared electrons are held more tightly in the reactants.
• The second law of thermodynamics states entropy always increases and reactions yield products with less ordered and less usable energy.
• The physical and chemical processes proceed toward lower potential energy and increased disorder.
• Breaking and forming bonds during chemical reaction depends on collisions between substances, allowing electrons to react.
• Reaction rate depends on the collision rate. The number of collisions depends on the temperature and concentration of the reactants.
• Higher temperature or concentration both lead to more collisions and a faster reaction.
• Carbon is the most versatile atom due to its valence, allowing for many covalent bonds.
• Organic compounds contain carbon bonded to other elements, resulting in limitless arrangements and bond combinations.
• Carbon–carbon bond formation was an important milestone in chemical evolution.
• A molecule’s chemical behaviour comes from H, N, and O, relating to the functional groups and their reactions.
• Six key functional groups include: Amino, Carbonyl, Carboxyl, Hydroxyl, Phosphate, and Sulphydryl. -Amino acts as a base, attracting protons. -Carbonyl (aldehydes, ketones) link molecules. -Carboxyl (carboxylic acid) tends to lose a proton. -Hydroxyl (alcohols) is very polar, acting as a weak acid. -Phosphate releases much energy when breaking bonds. -Sulphydryl, in proteins, forms disulphide bridges.
• Stanley Miller recreated chemical evolution, simulating early-Earth to form complex molecules with heat and electrical charges.
• These created precursors to life.

Chapter 2 explores water, carbon, and the chemical basis of life. Simple chemical compounds combined to form larger substances. A leading hypothesis says life began with chemical evolution and atom combinations. Four atoms make up 96% of matter in organisms.