2.0 Introduction

1. Enormous Explosion Marked the Beginning of the Universe
1. Beginning of the Universe Began the Process of Evolution
2. Evolution of Life and of Molecules Utilized Same Processes fig

1. Atoms
1. Universe Composed of Matter
1. All matter made of atoms fig 2.2

2. Difficult to study due to size

2. The Structure of Atoms
   1. Composed of smaller subatomic particles fig 2.3
      1. Protons (+ charge) and neutrons (0 charge) in central nucleus
      2. Electrons (–charge) in circular orbits around nucleus
      1. Same number as protons to balance charge
      2. Dictates chemical activity

   2. Atomic number = number of protons
   1. Neutrons and protons have same mass
   2. Only protons have electrical charge

3. Atomic Mass
1. Mass versus weight
1. Mass is the amount of a substance
2. Weight is the force of gravity exerted on it
3. Atomic mass = mass of protons + mass of neutrons

2. Mass measured in daltons
1. Proton or neutron is roughly 1 dalton
2. Electron is 1/1840 dalton, practically mass-less

3. Subatomic particles seen indirectly via collisions

4. Isotopes
1. All atoms of an element have the same atomic number (proton number)
2. An element cannot be broken into other substances by chemical means
3. Isotopes of an element have:
1. Same number of protons, different number of neutrons
2. Same number of electrons, thus same chemical properties

4. Example: Carbon-12 versus carbon-13 and carbon-14 fig 2.4
5. Unstable forms, like carbon-14, decay
1. Emit radioactive energy
2. Half-life = time for half of a sample's atoms to decay
3. Potential harmful side effects, exposure must be limited

5. Electrons
1. Electrically neutral atom has same number of electrons and protons
2. Electron orbit maintained by electrical attraction
3. In ions the number of electrons and protons are different
1. Element that possesses a net electrical charge
2. Positive charge if electron lost, a cation
3. Negative charge if electron gained, an anion

 

2. Electrons Determine the Chemical Behavior of Atoms
1. Arrangement Determines Chemical Properties of Element
1. Orbital describes probable, not actual location
2. Shapes differ fig 2.5
1. Inner s orbitals are spherical
2. More distant p orbitals are dumbbell-shaped
3. Maximum number of two electrons per orbital

3. Orbitals extremely far away from nucleus, atom mostly empty space
4. Nuclei of different atoms rarely contact one another
5. Electrons interact, determine chemical behavior

2. Energy Within the Atom
1. Electrons (–) are attracted to (+) protons
2. Energy required to keep electrons in orbit
3. Electron energy of position is potential energy fig 2.6
1. Moving electron away from nucleus
1. Requires energy
2. Electron then has more potential energy

2. Moving electron toward nucleus
1. Releases energy
2. Electron then has less potential energy

4. Exchange of electrons between molecules fig 2.7
1. Oxidation is a loss of electrons
2. Reduction is a gain of electrons
3. Chemical energy stored in electrons by oxidation-reduction reactions

5. Energy level schematics fig 2.8
1. Electrons represented as concentric rings called energy levels
2. Electrons in outer most rings hold more energy
3. Don't confuse energy levels and electron orbitals

 

1. Kinds of Atoms
1. 92 Naturally Occurring Elements
1. Have different number of protons, different arrangement of electrons
2. Mendeleev discovered pattern of chemical properties fig 2.9

2. The Periodic Table
1. Eight groups of repeating chemical properties
2. Based on interactions of valence electrons in outer shell
3. Maximum of eight electrons in outer shell of elements important to life
1. Elements at maximum are inert, not reactive
2. Elements with one less than maximum are highly reactive

4. Octet rule (rule of eight) states that atoms want their outer shell full

3. Distribution of Elements in Living Organisms tbl 2.1
1. Only eleven elements found in greater than trace amounts
2. Elements are generally light, atomic mass less than 21

1. Ionic Bonds Form Crystals
1. Molecule Is a Stable Group of Atoms
1. Compounds are molecules containing more than one kind of element
2. A chemical bond is the holding force
3. Atoms attracted by opposite electrical charges: Ionic bonds

2. A Closer Look at Table Salt
1. Atoms donate or receive electrons from other atoms fig 2.10
2. Example: Sodium chloride, common table salt
1. Sodium atom, loses electron = Na⁺
2. Chlorine atom, accepts electron = Cl^–

3. Resulting atoms become charged ions, an ionic compound
4. Bond forms by attraction of ions of opposite charges
1. Not between two individual atoms
2. Between one ion and all oppositely charged ions in vicinity

2. Covalent Bonds Build Stable Molecules
1. Covalent Bonds
1. Two atoms share one or more pairs of valence electrons
2. Example: Single bonded diatomic hydrogen (H₂) fig 2.11
1. Hydrogen has unpaired electron and unfilled outer level
2. Two atoms combine, each nucleus shares two electrons

3. Bond requires close proximity of atoms to one another
4. Bond is very stable
1. Has no net charge
2. Octet rule satisfied
3. Has no free electrons

2. Covalent Bonds Can be Very Strong
1. Strength of bond depends on number of shared electrons
1. Double bond shares two pairs of electrons, stronger than a single bond
2. Triple bond strongest covalent bond

2. Structural formulas: H – H or O = O
3. Molecular formulas: H₂ or O₂

3. Molecules with Several Covalent Bonds
1. Atoms can share electrons with more than one other atom
2. Example: Carbon, has six electrons, four in the outer level
1. To satisfy octet rule must gain four electrons
2. Thus can form four chemical bonds

4. Chemical Reactions
1. Formation and breaking of chemical bonds
2. Involve shifting atoms without change in number or identity
1. Reactants: Original, pre-reaction molecules
2. Products: Molecules resulting from a reaction

3. Influenced by several factors
1. Temperature: Heat increases rate
2. Concentration: Reactant versus product have opposite effect
3. Catalyst: Special substance increases rate

1. Chemistry of Water
1. Unique Properties of Water Necessary for Living Organisms
1. Exists as liquid at temperature of earth's surface fig 2.12
2. Provides a medium in which other molecules can interact
3. Composes two-thirds of most organisms

2. The Atomic Structure of Water
1. Forms weak chemical associations
2. Simple atomic structure, H₂O fig 2.13a
3. Forms chemical bonds much weaker than covalent bonds
1. Property is derived from structure of water
2. Responsible for organization of living chemistry

 

2. Water Atoms Act Like Tiny Magnets
1. Electronegativity Attracts Electrons of Water Molecules
1. Has distinct ends, each with a partial charge
2. Polar molecule results from magnet like poles
3. Polarity is crux of chemistry of water and life

2. Charge Separation Results in Polar Nature
1. Most stable configuration is tetrahedron, bond angle 104.5ðo fig 2.13b
1. Partial (ðd+) charges at apexes opposite hydrogens
2. Partial (ðd–) charge at oxygen

2. Polar molecules interact with one another
1. Opposite charges attract, form hydrogen bonds fig 2.14
2. Bonds are transient, cumulative effects important
3. Hydrogen bonds affect physical properties of water tbl 2.2

3. Water Clings to Polar Molecules
1. Polarity of Water Attracts It to Other Polar Molecules
1. Cohesion is attraction of water to water
1. Results in surface tension of water fig 2.15
2. Causes things to get wet in water

2. Adhesion is attraction of water to another molecule
1. Attraction is electrostatic
2. Results in capillary action, water rises in thin tube fig 2.16
3. Height inversely proportional to tube diameter

2. Water Stores Heat
1. Exhibits high specific heat
1. Amount of heat to change temperature of a substance
2. Associated with and proportional to polarity
3. Thermal energy must first disrupt hydrogen bonds
4. Heats up slowly, retains heat longer than surroundings

2. High heat of vaporization
1. Amount of heat required to change water to vapor
2. Evaporation of water produces cooling effect

3. Forms ice with decrease in temperature
1. Crystal-like lattice of hydrogen bonds fig 2.17
2. Less dense than liquid water

3. Water Is a Powerful Solvent
1. Water molecules gather around charged molecules
2. Example: Table sugar (sucrose)
1. Water forms hydrogen bonds with OH^– groups of sucrose
2. Each sugar molecule surrounded by cloud of water molecules
3. Cloud is called the hydration shell

3. Hydration shells also form around ions fig 2.18

4. Water Organizes Nonpolar Molecules
1. Water excludes nonpolar molecules
2. Preferentially forms hydrogen bonds with itself
3. Minimizes disruption of hydrogen bonding
1. Hydrophobic: Not soluble in water, nonpolar
2. Hydrophilic: Soluble in water, polar

4. Hydrophobic exclusion
1. Forces nonpolar molecules to associate together
2. Shapes molecules with nonpolar regions

 

4. Water Ionizes
1. Water Covalent Bonds May Break Spontaneously
1. Proton dissociates from molecule
1. Becomes positively charged ion (H⁺)
2. Remainder of molecule is OH^–

2. Ionization: Process of spontaneous ion formation
3. Mole of a substance is its molecular mass
1. Corresponds to combined atomic mass of all molecules
2. Molar concentration of H⁺ ions in water is 10^-7 mole/liter

2. pH
1. pH scale quantifies H⁺ concentration fig 2.19
1. pH = negative log of H⁺ ion concentration = –log[ H⁺]
2. pH of 7 indicates neutrality H⁺ ions = OH^– ions

2. Scale is logarithmic, change of one on scale is really tenfold

3. Acids
1. Substance that dissociates to increase concentration of H⁺
2. Has low pH value, below 7
3. Stronger acids have more H⁺ ions

4. Bases
1. Substance that combines with H⁺ ions when dissolved in water
2. Lower concentration of H⁺
3. Also called alkaline solutions, have pH value above 7

5. Buffers
1. pH of body fluids is about 7
2. Minimize changes in H⁺ and OH– concentration fig 2.20
3. Act as reservoirs for H⁺
1. Donate to solutions when concentration falls
2. Take from solutions when concentration increases

4. Example: Carbonic acid/bicarbonate in blood fig 2.21

6. Acid Rain fig 2.22