Chapter 3–4 Study Notes (Properties of Water; Organic Chemistry)

Chapter 3: Niagara Falls and Water Resources

• Niagara Falls infrastructure and sites (as listed on the transcript image/map):

  • Grand Island Underground Conduit

  • Reservoir

  • Toronto Power Generating Station

  • Niagara River Intake

  • Goat Island

  • Canadian Niagara Generating Station

  • CANADIAN (Horseshoe) FALLS

  • AMERICAN FALLS (Niagara Falls, N.Y.)

  • Rainbow Niagara Falls, Ont.

  • Bridge

  • Whirlpool Rapids

  • Bridges

  • Whirlpool Rapids

  • Devil's Hole Rapids

  • ROBERT MOSES POWERHOUSE, Lewiston Pump-Generating Station No. 2

  • Lewiston-Queenston International Bridge

  • NIAGARA ESCARPMENT

  • SIR ADAM BECK STATIONS

  • Aerial Cableway

  • The Whirlpool Reservoir No. 1; Queenston–Lewiston

• Additional geographical/administrative notes within the page: references to the Niagara Escarpment and various power-related facilities; mentions of Canada and US (Niagara Falls, Ont. and N.Y.).

• Visual/map citation: Bridge/Reservoirs/Powerhouses indicated; © 2010 EB, Inc. (image source).

Properties of Water

• Water is composed of cohesive molecules with surface tension; they take the shape of the vessel (Fig. 3.4).

• Water is liquid at room temperature.

• Water’s temperature rises and falls slowly (high heat capacity).

• Water has a high heat of evaporation (evaporative cooling).

• Ice is less dense than liquid water (ice floats) (Fig. 3.5/3.6).

• Water is the universal solvent for chemical reactions.

• All of these properties are due to hydrogen bonding.

Why ice floats and life-supporting properties of water

• Water is a polar molecule: the O atom bears partial negative charges, while H atoms bear partial positive charges.

• Hydrogen bonds form between oppositely charged regions of water molecules.

• In liquid water, hydrogen bonds continuously break and reform, allowing molecules to slip closer together.

• In ice, hydrogen bonds are stable and spacing increases, making ice less dense than liquid water.

• Floating ice insulates the water below, helping aquatic life survive.

• Water has additional life-supporting properties discussed in subsequent sections.

Ecological implications of loss of floating ice (illustrative examples)

• Less ice reduces feeding opportunities for polar bears (hunt from ice) → bears struggle to find food.

• Black guillemots in Alaska cannot fly from land to their fishing grounds if ice edge is too far, starving young birds.

• Loss of floating ice as habitat contributes to declines in Pacific walrus populations due to overcrowding and deadly stampedes.

• Warmer water and more light lead to more phytoplankton, which are consumed by other organisms; harmful algal blooms are a threat.

• Bowhead whales and some fish species may increase due to greater plankton availability.

• Geographic context shown: Arctic regions (Russia, Canada, Alaska) with a Sept. 2019 ice-extent map; reference to North Pole median extent (1981–2010) and Greenland.

Salt dissolution and hydration shells (Fig. 3.8 reference)

• When table salt dissolves in water, ions are surrounded by a hydration shell of water molecules.

• Negative regions on oxygens of polar water molecules are attracted to Na+ (cation) ions.

• Positive regions on hydrogens of polar water molecules are attracted to Cl- (anion) ions.

• A specific example notes an interaction where an oxygen of water is attracted to a partial positive charge on a lysozyme molecule (illustrating hydration around a solute).

• Guiding question: WHAT IF this solution is heated for a long time? (Prompt for thinking about solute stability and hydrolysis under heating.)

Acids and bases

• In pure water, concentrations of H+ and OH− are equal at 10⁻⁷ M each under standard conditions.

• Changes in the concentrations of these ions affect cellular reactions.

Acids and Bases: definitions and relationships

• Acid: increases [H+]. Example: HCl.

• Base: increases [OH−]. Example: NaOH, NH3+.

• Strong acid/base: complete dissociation in water.

• Weak acid/base: reversible dissociation. Example: H2CO3 ⇌ HCO3− + H+.

• pH and ion concentrations: at room temperature, [H+] and [OH−] are related by the water autoionization constant.

pH scale and related concepts

• pH is defined as ext{pH} = -\, ext{log}[ ext{H}^+].

• A low pH corresponds to acidic solutions; a high pH corresponds to basic solutions.

• Relation: [ ext{H}^+][ ext{OH}^-] = 10^{-14} at 25°C (neutral water).

• Example order of pH values (illustrative):

  • Battery acid: around 0

  • Digestive juices, lemon juice: around 1–2

  • Vinegar, beer, wine, cola: around 2–3

  • Tomato juice: around 4–5

  • Black coffee: around 5–6

  • Rainwater: around 6–6.5

  • Pure water: 7

  • Seawater: around 8

  • Milk of magnesia: around 10

  • Household ammonia: around 11

  • Oven cleaner: around 13–14

Scientific Skills Exercise

• Reference to Fig. 3.12 (p. 53) and Exercise (p. 54).

• Instructions: Submit ON PAPER by the end of class TODAY (9/1).

Chapter 4: Organic Chemistry

• Historical context (19th century): Could synthesize simple compounds from inorganic salts—NOT from living organisms (rejection of vitalism).

• Jons Jakob Berzelius: distinguished living (organic molecules) from nonliving (inorganic); introduced the term “vitalism.”

• 1828– Wohler: synthesized urea in the lab from inorganic chemicals, challenging vitalism.

• 1830s– Herman Kolbe: synthesized acetic acid from purified elements.

• Stanley Miller (1953): simulated primitive Earth conditions (H2O, H2, NH3, CH4) and used electrical discharge to trigger reactions, forming organic molecules (Fig. 4.2).

Carbon as a Building Block

• Carbon is nonpolar in many contexts and tetravalent (can form four bonds).

• Variations in carbon skeleton:

  • Number of carbon atoms

  • Branching patterns

  • Double bonds

  • Ring structures

• Example representations (Fig. 4.3): straight-chain, branched, double-bond-containing, and cyclic hydrocarbons.

Isomers

• Isomers: same molecular formula, different structures, and thus different chemical properties.

• Example: ext{C}4 ext{H}{10} can be n-butane or isobutane.

• Enantiomers: mirror-image isomers (handedness) shown in examples (Fig. 4.7).

Pharmacological Importance of Enantiomers

• Enantiomers can have different biological effects.

• Example: Ibuprofen has active S-enantiomer and inactive R-enantiomer in terms anti-inflammatory action, though R can contribute other effects.

• Example: Albuterol vs isomers (R- vs S-): R-enantiomer more effective for bronchodilation in asthma; S-enantiomer less effective or different effects (Fig. 4.8).

Functional Groups

• Definition: groups of atoms that replace hydrogen on hydrocarbons (Fig. 4.9, pg. 63).

• List of functional groups covered:

  • Hydroxyl

  • Carbonyl

  • Carboxyl

  • Amino

  • Sulhydryl

  • Phosphate

  • Methyl

Figure 4.9a: Hydroxyl group

• Structure: –OH (often written HO- or -OH)

• Example: Ethanol

• Properties:

  • Polar due to electronegative oxygen; can form hydrogen bonds with water.

  • Helps dissolve organic compounds such as sugars (as in alcohols).

Figure 4.9b: Carbonyl group

• Two main types depending on placement:

  • Ketone: carbonyl within the carbon skeleton (R-CO-R')

  • Aldehyde: carbonyl at the end of the carbon skeleton (R-CHO)

• Example: Acetone (ketone), Propanal (aldehyde)

• Relevance: Carbonyl groups are found in sugars, giving rise to two major sugar groups:

  • Ketoses (contain ketone)

  • Aldoses (contain aldehyde)

Figure 4.9c: Carboxyl group

• Structure: –COOH (carboxyl)

• Examples: Acetic acid

• Properties:

  • Acts as an acid; donates H+ because the C–O bond is highly polar.

  • In cells, usually exists in ionized form as carboxylate (–COO⁻).

Figure 4.9d: Amino group

• Structure: –NH2 (amino)

• Example: Glycine

• Properties:

  • Acts as a base; can accept an H+ in solution (forming –NH3⁺).

  • In cells, often exists in ionized form with a 1+ charge.

Figure 4.9e: Sulfhydryl (thiol) group

• Structure: –SH

• Example: Cysteine

• Properties:

  • Two –SH groups can react to form a covalent bond, cross-linking helps stabilize protein structure.

  • Cross-linking of cysteines in hair proteins influences curliness/straightness of hair.

Figure 4.9f: Phosphate group

• Structure: –O–P(=O)(–O)(–O)

• Example: Glycerol phosphate

• Properties:

  • Contributes negative charge to molecules; can be present as a diester or monoester (2− when at the end, 1− when internal in a chain of phosphates).

  • Molecules containing phosphate groups can release energy upon hydrolysis.

Figure 4.9g: Methyl group

• Structure: –CH3 (methylated group)

• Example: 5-Methyl cytidine

• Function:

  • Addition of methyl groups can affect DNA expression or interaction with DNA-binding molecules.

  • Methylation patterns influence the expression of genes and, in hormones, affect their shape and function.

Homework: Chapter 4 assignments and essay prompt

• Chapter 4, page 65: Work with a group (2–3 students) on #11; write a short essay together, ensuring each group member’s contribution is included under the HPU honor code.

• Submission: ONE Word document to Blackboard at class end, listing all group members.

• Essay prompt (Question 11):

  • In 1918, sleeping sickness caused an unusual rigid paralysis similar to Parkinson-like symptoms. Later, L-dopa (the levorotatory enantiomer) relieved paralysis, while D-dopa showed no effect. In a 100–200 word essay, discuss how the effectiveness of one enantiomer and not the other illustrates the theme of structure and function.

11. Enantiomer prompt (content reminder)

• The prompt asks you to reflect on how stereochemistry (enantiomerism) leads to different biological activities and why two molecules that are mirror images can have different effects in living systems.