Water & Carbon: The Chemical Basis of Life Chapter 2 Biology 11

Overview Basic Definitions Radioactive isotopes – not in chapter 2 text Understanding the four types of bonding Ionic Covalent Polar covalent hydrogen Water’s special properties Solvent Adhesion/cohesion Surface tension Specific heat

Overview Acid-Base Reactions and pH Chemical Energy Kinetic energy Potential energy Gibbs free-energy Chemical Evolution Functional Groups

Basic Definitions Elements – metals vs. nonmetals Atomic number – number of protons Mass number – protons and neutrons Isotopes Atomic mass unit – amu Orbitals – 3D shapes which holds electrons s, p, d, & f Valence electrons – outermost electrons Number of valence electrons determines chemical properties and chemical reactivity

Isotopes Isotopes are atoms that have the same number of protons but differ in the number of neutrons Not all isotopes are radioactive Radioactive means that the atom is trying to decay or reach a more stable state Three types of energy emission Alpha, beta, & gamma

Radioactive Isotopes

Four Types of Chemical Bonds Ionic Covalent Polar Covalent Hydrogen

Ionic Bonding Ionic Between a metal and a non metal Greatly different values of electronegativity Involves a complete transfer of electrons Held together by an electrostatic interaction between a + charge and a – charge Metal is always giving electrons Non metal is always accepting electrons In solution become “ions”

Ionic Bonding

Covalent Bonding Covalent Always between two non-metals No difference in electronegativity No charges present Strong bond Held together by the attraction of an electron of one atom to the nucleus of the other atom Equal sharing of electrons

Covalent Bonding

Polar Covalent Two non metals Differences in electronegativity Electronegativity is the ability of an atom to pull an electron to itself in a chemical bond One atom is more greedy than the other and therefore the electron of the less greedy atom spends more time around the nucleus of the greedy atom Creates a dipole moment Partial positive charge δ+ and δ-

Bonding and Solubility Like dissolves like Solubility is the ability of water to “coat” another molecule or to interact chemically with that molecule Molecules with a great deal of covalent bonding and not much polar covalent bonding are hydrophobic Waxes, oils, fats Molecules with many polar covalent bonds are easily soluble in water glucose

Bonding Summary

Hydrogen Bonding All of the bonds this far have been intramolecular Hydrogen bonding is an intermolecular force Week interaction always involving a hydrogen atom on one molecule and either an oxygen, or nitrogen on another atom Responsible for water’s special properties Holds together DNA double helix tRNA structure 3D protein structure (alpha helix & beta sheets)

Hydrogen Bonding in DNA

Water Universal Solvent Water easily dissolves ionic and polar molecules To be dissolved in water is to be surrounded and coated by water molecules

Water Universal Solvent Hydrophobic molecules repel water

Cohesion Binding between like molecules Transpiration in trees Meniscus Surface tension

Adhesion Binding between unlike molecules Usually between a liquid and solid surface

Density of Ice and Water When water freezes each water molecule must form four hydrogen bonds This forms a regular and repeating structure which has air space between the molecules This is why ice is less dense than liquid water

Specific Heat Water has a high capacity for absorbing heat Specific heat Amount of energy required to raise the temperature of 1 gram of a substance by one degree C. Before heat can be transferred so that the water molecules can move faster (increased kinetic energy increased heat) the hydrogen bonds must be broken

Heat of Vaporization Energy required to change one gram of liquid water to water vapor (gas) Why is water such an efficient coolant Water molecules have to absorb a great deal of energy from your body in order to evaporate You loose heat

Acids and Bases In chemical reality protons do not exist by themselves Protons associate with water to form hydronium ions H 2 O + H 2 O H 3 O + + OH - H 2 O H + + OH -

Acids and Bases Substances that give up protons during chemical reaction and raise the hydrogen ion concentration are acids Substances that acquire protons during chemical reactions and lower the hydrogen ion concentration are bases Acid base reactions require a proton donor and a proton acceptor HCl + H 2 O H 3 O + + Cl -

pH Calculations & pH Scale pH = -log[H + ]

Basic Terms of Chemical Reactions Reactants Products Chemical Equilibrium Forward and reverse reactions occur at the same rate The amount of reactant and product are not necessarily the same Exothermic – energy given out to system Endothermic – energy consumed

Energy Dynamics Potential Energy Kinetic energy Thermal energy kinetic energy of molecular motion 1 st law of thermodynamics 2 nd law of thermodynamics

Spontaneous Reactions ∆G = ∆H – T∆S ∆G negative = spontaneous Exergonic energy releasing ∆G positive = not spontaneous Endergonic energy consuming Reactions are spontaneous when ∆H is negative and ∆S is positive We have to use the combined contributions of changes in heat and disorder to determine spontaneity

Understanding ∆H Enthalpy ∆H is the difference in potential bond energy between the products and reactants ∆H reflects the number and kinds of chemical bonds in reactants and products When heat content of the product is less than the reactant ∆H is negative and exothermic Gives off heat to surroundings - ∆H When heat content of the reactants is less than the products ∆H is positive and endothermic Takes heat in from surroundings + ∆H

Bond Enthalpy You can also think of this as the bonds in methane hold more energy than the bonds of CO 2 or it takes more energy to form methane bonds than CO 2 bonds

Understanding ∆S Entropy Measurement of disorder Reactions are spontaneous when the products molecules are less ordered than the reactant molecules

Chemical Evolution First molecules on a hot earth CH 4, NH 3, H 2 O, CO 2, N 2 Spontaneous generation must have occurred at some point in earth’s history Chemical evolution Early in earth’s history simple inorganic molecules in the atmosphere and oceans combined to form larger more complex molecules

Chemical Evolution Kinetic energy and heat from sunlight was converted into chemical bonds Larger molecules accumulated and reacted with one another to produced more complex molecules One of these complex molecules was able to self replicate The big shift As the molecule multiplied evolution by natural selection replaced chemical evolution

Formation of Early Complex Molecules Using only the chemical precursors of the early atmosphere could these molecules form Formaldehyde H 2 CO Hydrogen cyanide HCN Reaction between CO 2 and H 2 is endergonic Formaldehyde and water have more potential energy and are more ordered

Energy Inputs and Chemical Evolution When earth’s early inorganic substances are placed in a test tube nothing happens But what happens when these molecules are struck by sunlight or lightening? In the early earth’s atmosphere many high energy photons would have reached the planet? Why

Energy Inputs and Chemical Evolution Energy from photons can break molecules apart by knocking electrons off Free radicals form which are highly reactive

Temperature and Early Chemical Reactions For the complex molecules to form from the inorganic molecules one chemical bond must break and one chemical bond must form Reactants must collide When temperature are high reactants move faster (increased kinetic energy) and collide more frequently

Chemical Evolution Sunlight was converted into chemical energy Potential energy now held in chemical bonds Why was HCN and H 2 CO so important The formation of C – C bonds was possible Heat alone can link to formaldehyde molecules into acetaldehyde Reactions between acetaldehyde and formaldehyde produce sugars Crucial step towards production of the types of molecules found in living organisms

Step 1 Chemical Evolution

Step 2 Chemical Evolution

Step 3 Chemical Evolution

Water’s Specific Heat & Chemical Evolution Water’s high specific heat insulated dissolved substances from sources of energy like intense sunlight which could have broken the chemical bonds apart Water’s heat of vaporization would have kept land masses near water cool for further chemical evolution

Importance of Carbon Because carbon can form 4 bonds it can form has a limitless array of molecular shapes The carbon atoms in an organic molecule furnish a skeleton that gives the molecule its overall shape However, the type of macromolecule and the types of reactions that a molecule can participate in is dictated by functional groups Review Table 2.3 of your text