AP1 - Ch. 2 Chemistry

Chapter 2: Chemistry Comes Alive

2.1 Matter

• The "stuff" of the universe.

• Anything that has mass and takes up space.

States of Matter

1. Solid: Has a definite shape and volume.

2. Liquid: Has a definite volume but changeable shape.

3. Gas: Has both changeable shape and volume.

2.2 Energy

• The capacity to do work (put matter into motion).

Types of Energy

1. Kinetic: Energy in action.

2. Potential: Energy of position; stored (inactive) energy.

Forms of Energy

• Chemical: Stored in the bonds of chemical substances.

• Electrical: Results from the movement of charged particles.

• Mechanical: Directly involved in moving matter.

• Radiant or Electromagnetic: Energy traveling in waves (e.g., visible light, ultraviolet light, and $X$-rays).

Energy Form Conversions

• Energy is easily converted from one form to another; during conversion, some energy is "lost" as heat.

2.3 Composition of Matter

• Elements: Unique substances that cannot be broken down by ordinary chemical means.

• Atoms: More-or-less identical building blocks for each element.

• Atomic Symbol: One- or two-letter chemical shorthand for each element.

Properties of Elements

• Each element has unique physical and chemical properties:

  1. Physical properties: Detected with our senses.

  2. Chemical properties: Pertaining to the way atoms interact with one another.

Major Elements of the Human Body

1. Oxygen ($O$)

2. Carbon ($C$)

3. Hydrogen ($H$)

4. Nitrogen ($N$)

Lesser and Trace Elements of the Human Body

• Lesser elements: Make up $3.9\%$ of the body and include:

  • Calcium ($Ca$), Phosphorus ($P$), Potassium ($K$), Sulfur ($S$), Sodium ($Na$), Chlorine ($Cl$), Magnesium ($Mg$), Iodine ($I$), and Iron ($Fe$).

• Trace elements: Make up less than $0.01\%$ of the body; required in minute amounts and are found as part of enzymes.

2.4 Atomic Structure

• The nucleus consists of neutrons and protons.

  1. Neutrons: Have no charge and a mass of $1$ atomic mass unit (amu).

  2. Protons: Have a positive charge and a mass of $1$ amu.

  3. Electrons: Found orbiting the nucleus, having a negative charge and $1/2000$ the mass of a proton ($0$ amu).

Models of the Atom

1. Planetary Model: Electrons move around the nucleus in fixed, circular orbits.

2. Orbital Model: Regions around the nucleus where electrons are most likely to be found.

Identification of Elements

• Atomic Number: Equal to the number of protons.

• Mass Number: Equal to the mass of protons and neutrons.

• Atomic Weight: Average of the mass numbers of all isotopes.

• Isotope: Atoms with the same number of protons but a different number of neutrons.

• Radioisotopes: Atoms that undergo spontaneous decay called radioactivity.

Isotopes of Hydrogen

1. Hydrogen ($1H$): ($1p^+$; $0n$; $1e^-$)

2. Deuterium ($2H$): ($1p^+$; $1n^0$; $1e^-$)

3. Tritium ($3H$): ($1p^+$; $2n^0$; $1e^-$)

Radioisotopes

• Are unstable because they contain excess neutrons.

• Lose nuclear components in the form of high-energy radiation (alpha particles, beta particles, gamma rays).

• Biological Half-life: The time required for half of the radioactive material from a test to be eliminated from the body.

  • $T_{1/2} (C^{14}) = 5730$ years.

  • $T_{1/2} (H^3) = 12.2$ years.

  • $T_{1/2} (N^{16}) = 7.13$ seconds.

  • $T_{1/2} (Ba^{124}) = 10.6$ minutes.

2.5 Molecules and Compounds

• Molecule: Two or more atoms held together by chemical bonds.

  • Elements: Molecule of element—two or more of the same kind of atoms chemically bonded together.

  • Compound: Molecule of compound—two or more different kinds of atoms.

2.6 Mixtures

• Most matter exists as mixtures: two or more components that are physically intermixed.

Three Basic Types of Mixtures

1. Solutions: Homogeneous mixtures; particles are evenly distributed throughout.

   • Solvent: Substance present in greatest amount (usually a liquid, such as water).

   • Solute(s): Substances dissolved in solvent (present in smaller amounts).

   • Example: Blood sugar—glucose is solute, and blood (plasma) is solvent. True solutions are usually transparent.

2. Colloids: Also known as emulsions; heterogeneous mixtures; particles are not evenly distributed. Can see large solute particles in solution (cloudy appearance; do not settle out).

   • Examples: Jell-O; cytosol of the cell.

3. Suspensions: Heterogeneous mixtures with large, visible solutes that do settle out.

   • Example: Mixture of water and sand. Blood is considered a suspension because blood cells settle out when left in a tube.

Concentration of Solutions

• Percent (parts per $100$ parts) of solute in total solution.

  • Example: $10$ parts salt to $90$ parts water is a $10\%$ salt solution.

• Milligrams per deciliter ($mg/dl$) used for measuring conditions.

  • Example: Normal fasting blood glucose levels are around $80$ mg/dl.

2.7 Chemical Bonds

• Electron shells, or energy levels, surround the nucleus of an atom. Bonds are formed using the electrons in the outermost energy level.

• Valence Shell: Outermost energy level containing chemically active electrons.

• Octet Rule: Atoms interact to have $8$ electrons in their valence shell (except the first shell, which is full with $2$ electrons).

• Chemically Inert Elements: Outermost energy level fully occupied by electrons (e.g., Helium ($He$) and Neon ($Ne$)).

• Chemically Reactive Elements: Outermost energy level not fully occupied by electrons (e.g., Hydrogen ($H$), Carbon ($C$), Sodium ($Na$), and Oxygen ($O$)).

Types of Chemical Bonds

1. Ionic Bonds: Result from the gain or loss of electrons forming ions.

   • Anions: Have gained one or more electrons.

   • Cations: Have lost one or more electrons.

   • Formation: Formed by the transfer of electrons, forming crystals (e.g., Sodium chloride ($NaCl$)).

2. Covalent Bonds: Formed by the sharing of two or more electrons, producing molecules.

   • Examples: Single, Double, and Triple Covalent Bonds.

3. Hydrogen Bonds: Too weak to bind atoms together; common in dipoles such as water ($H_2O$), responsible for surface tension.

Polar and Nonpolar Molecules

• Nonpolar Molecules: Electrons shared equally between atoms.

• Polar Molecules: Unequal sharing of electrons, creating slight negative and positive charges at opposite ends (e.g., Water ($H_2O$)).

2.8 Chemical Reactions

• Occur when chemical bonds are formed, rearranged, or broken.

• Chemical Equations: Contain the number and type of reacting substances (reactants) and products produced.

Patterns of Chemical Reactions

1. Combination reactions: $A + B \rightarrow AB$ (bond formation).

2. Decomposition reactions: $AB \rightarrow A + B$ (breaking down molecules).

3. Exchange reactions: Bond formation and breaking (e.g., $AB + C \rightarrow AC + B$).

   • Oxidation-Reduction (Redox): Reactants losing electrons are electron donors and are oxidized; reactants taking up electrons are electron acceptors and become reduced.

Energy Flow and Rate

• Exergonic Reactions: Reactions that release energy.

• Endergonic Reactions: Products contain more potential energy than reactants.

• Factors Influencing Rate:

  • Temperature (higher is faster), Particle Size (smaller is faster), Concentration (higher is faster), and Catalysts (enzymes increase rates without being changed).

2.9 Biochemistry

• Organic Compounds: Contain carbon, covalently bonded, often large.

• Inorganic Compounds: Do not contain abundant carbon (e.g., water, salts, and many acids and bases).

Properties of Water

1. High Heat Capacity: Absorbs/releases large amounts of heat before changing temperature.

2. High Heat of Vaporization: Large amounts of heat needed to change from liquid to gas.

3. Polar Solvent Properties: Dissolves ionic substances; major transport medium.

4. Reactivity: Important in hydrolysis and dehydration synthesis reactions.

5. Cushioning: Acts as a resilient cushion around certain body organs.

Salts, Acids, and Bases

• Salts: Ionic compounds that dissociate in water; ions (electrolytes) conduct electrical currents (e.g., sodium, potassium, calcium, iron).

• Acids: Proton (hydrogen ion) donors; release $H^+$ in solution (e.g., $HCl \rightarrow H^+ + Cl^-$).

• Bases: Proton acceptors; take up $H^+$ from solution (e.g., $NaOH \rightarrow Na^+ + OH^-$; bicarbonate ion ($HCO<em>3^-$) and ammonia ($NH</em>3$)).

• Buffers: Systems like the Carbonic Acid-Bicarbonate System that resist abrupt large swings in pH of body fluids.

2.10 Organic Compounds

1. Carbohydrates: Carbon, hydrogen, oxygen ($C$, $H$, $O$); supply cellular food.

   • Monosaccharides (e.g., glucose), Disaccharides (e.g., sucrose), Polysaccharides (e.g., starch, glycogen).

2. Lipids: $C$, $H$, and $O$ (less oxygen). Includes Neutral Fats (triglycerides), Phospholipids (cell membranes), Steroids (cholesterol), and Eicosanoids.

3. Proteins: Polymers of amino acids ($20$ types) joined by peptide bonds. Structure includes Primary, Secondary, Tertiary, and Quaternary levels. Includes Fibrous (support) and Globular (functional) proteins.

   • Denaturation: Unfolding due to drops in pH or increased temperature.

4. Nucleic Acids: DNA and RNA. Structural unit is the nucleotide (nitrogenous base, pentose sugar, phosphate group).

   • DNA: Double-stranded genetic instructions; uses $A$, $G$, $C$, and $T$.

   • RNA: Single-stranded; uses $U$ instead of $T$.

5. Adenosine Triphosphate (ATP): Immediately usable energy for the cell; breaks down into $ADP$ and phosphate to release energy.

End of Chapter 2 Study Guide