The Chemical Basis of Life Chapter 2 The Chemical Basis of Life 1

I. Elements: Chemical Symbols: Abbreviations for the name of each element. Usually one or two letters of the English or Latin name of the element First letter upper case, second letter lower case. Example: Helium (He), sodium (Na), potassium (K), gold (Au).

II. Structure & Properties of Atoms Atoms: Smallest particle of an element that retains its chemical properties. Made up of three main subatomic particles. Particle Location Mass Charge Proton (p+) In nucleus 1 +1 Neutron (no) In nucleus 1 0 Electron (e-) Outside nucleus 0* -1 * Mass is negligible for our purposes.

Atomic Particles: Protons, Neutrons, and Electrons. Helium Atom Atomic Particles: Protons, Neutrons, and Electrons Helium Atom Carbon Atom

Structure and Properties of Atoms 1. Atomic number = # protons The number of protons is unique for each element Each element has a fixed number of protons in its nucleus. This number will never change for a given element. Written as a subscript to right of element symbol. Examples: C6 , O8 Because atoms are neutral (no charge), the number of electrons and protons are always the same. In the periodic table elements are organized by increasing atomic number.

Structure and Properties of Atoms: 2. Mass number = # protons + # neutrons Gives the mass of a specific atom. The number of neutrons may vary. The number of neutrons can be determined by: # neutrons = Mass number - Atomic number

III. How Atoms Form Molecules: Chemical Bonds Molecule: Two or more atoms combined chemically. Compound: A substance with two or more elements combined in a fixed ratio. Water (H2O) Hydrogen peroxide (H2O2) Carbon dioxide (CO2) Carbon monoxide (CO) Table salt (NaCl) Atoms are linked by chemical bonds. Chemical Formula: Describes the chemical composition of a molecule of a compound. Symbols indicate the type of atoms Subscripts indicate the number of atoms

Sodium has 11 electrons, 1 valence electron. How Atoms Form Molecules: Chemical Bonds “Octet Rule”: When the outer shell of an atom is not full, i.e.: contains fewer than 8 (or 2 for He) valence electrons, the atom tends to gain, lose, or share electrons to achieve a complete outer shell (8 or 2) electrons. Example: Sodium has 11 electrons, 1 valence electron. Sodium loses its electron, becoming an ion: Na -------> Na+ + 1 e- 1(2), 2(8), 3(1) 1(2), 2(8) Outer shell has 1 e- Outer shell is full Sodium atom Sodium ion

N = N -- There may be more than one covalent bond between atoms: 1. Single bond: One electron pair is shared between two atoms. Example: Chlorine (Cl2), water (H2O); methane (CH4) Cl --- Cl 2. Double bond: Two electron pairs share between atoms. Example: Oxygen gas (O2); carbon dioxide (CO2) O=O 3. Triple bond: Three electron pairs shared between two atoms. Example: Nitrogen gas (N2) N = N --

Single and Double Covalent Bonds

Water: A Unique Compound for Life

Water: The Ideal Compound for Life Living cells are 70-90% water Water covers 3/4 of earth’s surface Water is the ideal solvent for chemical reactions On earth, water exists as gas, liquid, and solid

I. Polarity of water causes hydrogen bonding Water molecules are held together by H-bonding Partially positive H attracted to partially negative O atom. Individual H bond are weak, but the cumulative effect of many H bonds is very strong. H bonds only last a fraction of a second, but at any moment most molecules are hydrogen bonded to others.

Hydrogen Bonds in Water are Responsible for Many of its Properties

Ice Forms Stable Hydrogen Bonds and is Less Dense than Water

Unique properties of water caused by H-bonds Stable Temperature: Water resists changes in temperature because it has a high specific heat. Specific Heat: Amount of heat energy needed to raise 1 g of substance 1 degree Celsius Specific Heat of Water: 1 calorie/gram/oC High heat of vaporization: Water must absorb large amounts of energy (heat) to evaporate. Heat of Vaporization of Water: 540 calorie/gram. Evaporative cooling is used by many organisms to regulate body temperature. Sweating Panting

Unique properties of water caused by H-bonds Universal Solvent: Dissolves many (but not all) substances to form solutions. Solutions are homogeneous mixtures of two or more substances (salt water, air, tap water). All solutions have at least two components: Solvent: Dissolving substance (water, alcohol, oil). Aqueous solution: If solvent is water. Solute: Substance that is dissolved (salt, sugar, CO2). Water dissolves polar and ionic solutes well. Water does not dissolve nonpolar solvents well.

A Salt Crystal Dissolving in Water

Solubility of a Solute Depends on its Chemical Nature Solubility: Ability of substance to dissolve in a given solvent. Two Types of Solutes: A. Hydrophilic: “Water loving” dissolve easily in water. Ionic compounds (e.g. salts) Polar compounds (molecules with polar regions) Examples: Compounds with -OH groups (alcohols). “Like dissolves in like”

Solubility of a Solute Depends on its Chemical Nature Two Types of Solutes: B. Hydrophobic: “Water fearing” do not dissolve in water Non-polar compounds (lack polar regions) Examples: Hydrocarbons with only C-H non-polar bonds, oils, gasoline, waxes, fats, etc.

ACIDS, BASES, pH AND BUFFERS A. Acid: A substance that donates protons (H+). Separate into one or more protons and an anion: HCl (into H2O ) -------> H+ + Cl- H2SO4 (into H2O ) --------> H+ + HSO4- Acids INCREASE the relative [H+] of a solution. Water can also dissociate into ions, at low levels: H2O <======> H+ + OH-

B. Base: A substance that accepts protons (H+). Many bases separate into one or more positive ions (cations) and a hydroxyl group (OH- ). Bases DECREASE the relative [H+] of a solution ( and increases the relative [OH-] ). H2O <======> H+ + OH- Directly NH3 + H+ <=------> NH4+ Indirectly NaOH ---------> Na+ + OH- ( H+ + OH- <=====> H2O )

HCl (aq) -------------> H+ + Cl- Strong acids and bases: Dissociation is almost complete (99% or more of molecules). HCl (aq) -------------> H+ + Cl- NaOH (aq) -----------> Na+ + OH- (L.T. 1% in this form) (G.T. 99% in dissociated form) A relatively small amount of a strong acid or base will drastically affect the pH of solution. Weak acids and bases: A small percentage of molecules dissociate at a give time (1% or less) H2CO3 <=====> H+ + HCO3- carbonic acid Bicarbonate ion (G.T. 99% in this form) (L.T. 1% in dissociated form)

C. pH scale: [H+] and [OH-] pH scale is used to measure how basic or acidic a solution is. Range of pH scale: 0 through 14. Neutral solution: pH is 7. [H+ ] = [OH-] Acidic solution: pH is less than 7. [H+ ] > [OH-] Basic solution: pH is greater than 7. [H+ ] < [OH-] As [H+] increases pH decreases (inversely proportional). Logarithmic scale: Each unit on the pH scale represents a ten-fold change in [H+].

D. Buffers keep pH of solutions relatively constant Buffer: Substance which prevents sudden large changes in pH when acids or bases are added. Buffers are biologically important because most of the chemical reactions required for life can only take place within narrow pH ranges. Example: Normal blood pH 7.35-7.45. Serious health problems will arise if blood pH is not stable.

CHEMICAL REACTIONS A chemical change in which substances (reactants) are joined, broken down, or rearranged to form new substances (products). Involve the making and/or breaking of chemical bonds. Chemical equations are used to represent chemical reactions. Example: 2 H2 + O2 -----------> 2H2O 2 Hydrogen Oxygen 2 Water Molecules Molecule Molecules

Chemical Reactions Require Making and Breaking Bonds