The Importance of Water for Life

• The importance of water for life is explored, including its role in bodily functions, digestion, and temperature regulation.

Properties of Water

• A human being can survive for more than two weeks without food, but only three or four days without water.
• Prolonged water deprivation causes intense thirst, dry skin and mucous membranes, fever, cardiac collapse, coma and death.
• Water accounts for 70% of the human body weight and 98% of algae, and is needed for breathing, eye lubrication, detoxification, and maintaining body temperature.
• Water participates in hydrolysis during digestion, aiding in substance transport, lubrication, and tissue flexibility.
• Respiration needs moisture, saliva aids in food ingestion, blood transports nutrients, and tears keep eyes lubricated; water is lost through transpiration, respiration, and urination.
• Plants need water to produce their food and to grow, regulating temperatures in living things and the biosphere because of its high heat capacity.
• This means its temperature does not change as rapidly as other liquids do, and that it is used as a source of hydrogen atoms in photosynthesis.

Chemical Composition of Water

• Each water molecule has two hydrogen atoms joined to one oxygen atom by covalent bonds involving shared electron pairs.
• The oxygen atom is more electronegative, rendering the electron pair closer and determining the polarity of the bond.
• Polarity represents the separation of electrical charges within the molecule, giving water an angular structure with a bond angle of 104.5°.
• The water molecule is a dipole as oxygen has a negative charge and the hydrogen has the positive

Polarity for universal dissolving

• Water dissolves more substances than other liquids as it can interact with polar molecules.
• This dissolving capability is important for metabolism, nutrient transport, and waste removal in living beings.
• Temperature increases the solubility of substances in water, which has a high dielectric constant and dissolves ionic and polar covalent compounds.
• Polar water molecules arrange around solute groups, separating ionic compounds into anions and cations surrounded by water molecules, a phenomenon known as ionic solvation.
• However, water does not dissolve nonpolar substances like fats, oils, cellulose, metals, or quartz.

Physical properties of water

• Pure water is odorless and tasteless with a bluish tint detectable in thick layers.
• At atmospheric pressure (760 mm of mercury), water has a freezing point of 0°C and a boiling point of 100°C.
• Water reaches maximum density at 4°C and expands when frozen. It also decreases in density as its temperature increases above 4°C.
• The lower density of ice and its low thermal conductivity mean ice will remain on the surface of bodies of water, allowing for aquatic life to thrive.
• Liquid water is denser than ice between both 0°C, reaching maximum density at 3.98°C. Otherwise water behaves like other liquids and diminishing in density as temperature rises, but expands again at sub zero temperatures.

Intermolecular Forces

• Intermolecular forces exist between molecules and determine the physical properties of liquids and molecular solids.
• These forces are weaker and shorter-range than ionic or covalent bonds.
• The intensity of intermolecular forces determines a liquid's boiling point and the melting point of a solid.

Types of Intermolecular Forces

• Dipole forces and dispersion forces (van der Waals forces) exist among molecules, and hydrogen bonds are a strong type of dipole-dipole interaction.
• Three types of attraction forces exist between molecules.
• Dipole forces occur in molecules that have dipoles, leading to attractions between oppositely charged poles.
• Dipole force attractions are more intense than the forces between nonpolar molecules of comparable molar mass.
• Examples of compounds with dipole molecules include hydrogen bromide (HBr), hydrogen chloride (HCl), and angular molecules like hydrogen sulfide (H₂S).
• Hydrogen bonds typically occur in polar molecules that have hydrogen atoms bonded to fluorine, oxygen, or nitrogen.
• A hydrogen bond forms from the attraction of a positive charge in one molecule to the negative charge on a nearby molecule, specifically to a non-shared electron pair on a fluorine, oxygen, or nitrogen atom.
• Hydrogen bonds are stronger than dipole-dipole or dispersion forces, but are weaker than standard chemical bonds, therefore compounds with hydrogen bonds have higher boiling points compared to other hydrogen compounds with elements in the same group.
• Compounds able to make hydrogen bridges are hydrogen fluoride (HF), water (H₂O), methanol (CH3OH), ethanol (C₂H₂OH), and ammonia.

Hydrogen Bridges

• The polarity of water molecules results in an interactive force between a positive pole of one molecule and the negative pole of another, thereby creating a hydrogen bridge.
• Hydrogen bridges cause water molecules to bond tightly, forming an almost incompressible liquid.
• Hydrogen bridges cause strong adhesion, especially where it joins other polar molecules like plant sap.
• Water has a high specific heat because it uses large quantities of heat to break the hydrogen bridges, releasing energy slowly as it cools

Dispersion Forces

• Fritz London proposed dispersion forces in 1930, describing interactions between particles that create momentary dipoles because electrons move within an atom or molecule.
• A temporary dipole in one atom can induce a dipole in a neighboring atom, causing attraction.
• Dispersion forces are stronger in nonpolar molecules of larger size because force intensity grows with molecular weight and subsequently increases boiling points.
• Dispersion forces are weaker than dipole attractions or ionic bonds.

Comparing Intermolecular Forces

• Dispersion forces exist in all substances, increasing with molecular weight.
• Dipole forces occur in polar molecules, adding to the effect of dispersion forces.
• Hydrogen bridges are the most intense kind of intermolecular force.
• Both ionic and covalent bonds are more intense than any intermolecular force.
• Nonpolar molecules usually display only dispersion forces, while polar molecules may display dispersion, dipole forces, or hydrogen bridges if the compounds contain hydrogen bonded to fluorine, oxygen, or nitrogen.

Relationship between Intermolecular Forces and Boiling Point

• Boiling point refers to the temperature at which a substance changes from a liquid to a gaseous state.
• Boiling point increases as intermolecular forces become stronger. If intermolecular forces are weak, boiling points are lower

Water Differences

• Water offers specific properties making it essential for life.
• Water molecules are strongly attracted because of hydrogen bridges forming a compact structure, leading to high adhesion as they adhere to other polar molecules.

Surface Tension

• Intermolecular attraction and repulsion produce properties like boiling point, melting point, vaporization heat and surface tension.
• A molecule will experience a net inward attraction at the surface, due to the surrounding water molecules.
• Surface tension supports resistance to surface penetration, helps form the spherical shape of liquid droplets, and also supports the floating of surface objects.

Water in the Hydrosphere

• Water has a substantial presence in the atmosphere, and is the main liquid in the hydrosphere.
• Life on Earth originally came from water.
• 68% of the world's water is salt water, 29% is in soils, 2.997% is fresh in icecaps, and only 0.003% is fresh for human use.
• Atmospheric pressure and temperature regulate water's physical state.
• Water vapor is lighter than air, so it can rise high in the atmosphere, forming clouds.
• Hydrolysis reactions neutralize acids and bases, water can behave as either, being an amphiprotic substance.
• Water has other traits such as the capacity to absorb upwards of 99% of solar radiation, and easily carry sound.
• Potable water conducts electricity, having dissolved air and ions.
• Water must have limited calcium and magnesium salts for multiple applications in refrigeration, steam boiler feeding, and cleaning textiles, therefore soft water lacks a high concentration of salts, whereas hard water contains high salt mineral content.

Dissolutions

• Mixtures either homogeneous or heterogeneous exist, but just certain types constitute a dissolution, or a solution.
• Sugar dissolved in water makes for a homogeneous mixture where the components form a single phase.
• Homogeneous mixtures exhibit a consistent appearance and the components are indistinguishable.
• In that case, a solution is a homogeneous system exhibiting identical compositional properties throughout.

Components of a Dissolution

• A dissolution is composed of a solvent, which exists in larger quantities, and a solute that exists in smaller quantities and dissolves to make a homogenous system.
• A dissolution can have one or multiple solutes.

Types of Dissolutions

• Classifying dissolutions is done either by physical state of the components or relative proportions.

Classifying by Component State

• The solvent dictates physical state, the solute exist in solid, liquid, or gas.
• Examples include liquid, solid, and gaseous dissolutions.

Classifying by relative component portions

• This classification involves gauging concentration.
• This factor impacts if a compound dissolves in another and defines if it is saturated, unsaturated, or supersaturated.
• A saturated solution has max solute concentration for stability given a temp, adding solute to saturation results in crystallization.
• A saturated solution occurs where solute and solvent exist in equilibrium alongside solution and crystallization.
• An unsaturated solution has less solute than it can need to form a saturated solution.
• For example, saline solution will have less than 36 g of salt per 100 g of water.
• Supersaturated solutions have more solute than needed for a saturated solution, but they are highly unstable.
• The types of dissolutions depend on crystallization vs dissolving.

Dilute vs. Concentrated

• A solution exhibits diluteness when the solute contents have negligible amounts relative to volume, but otherwise will be concentrated.
• How strong a cup of coffee is demonstrates that concentrated solutions require more dissolved solute and vice-versa.

Dissolution Process

• Mixing a solvent with a solute results in solute-solute, solvent-solvent and solvent-solute interactions.
• Dissolution happens when solute-solvent interaction is prominent, allowing the solute particles to abandon their structure.
• A solvent is considered a hydration in water and a solvatation otherwise.

Solubility

• Solubility reveals the amount of solute dissolved in a solvent at a temperature.
• Sodium chloride solubility (NaCl) measures 36 g per 100 g of water at 20°C.
• Substances either soluble or insoluble, and liquids either miscible or immiscible, depending on whether the sugar diffuses through a given fluid.
• Miscible liquids, as ethanol and water, readily mix. Liquids that do not mix, such as oil and water, are immiscible.

Impacting Factors

• The quantity of solute depends upon these factors
• The nature of compounds determine the solubility, as similar compounds dissolve each other through electrical properties, like water (polar) dissolving alcohol (polar).
• Hydrogen bridges among solutes and solvents raises solubility, like water with ethanol.

Temperature and pressure

• Increasing temperatures enhance the dissolving of solids in liquids while also establishing the solubility equilibrium.
• Increasing temperatures boost the dissolving of primarily endothermic solids due to molecular motion.
• The opposite can be said of gas dissolution, increasing pressure increases the likelihood of dissolution over temperature.

Dissolving Velocity

• Solute dissolving velocity determines its concentration within a solvent.
• Such can be modulated by the subdivision state, velocity, and temperature of a fluid.
• A solute dissolves faster when crushed due to the increase in the solvent-solute interactions.
• Agitation augments the velocity of dissolution by increasing direct contact between undissolved and dissolved particles.
• Temperature enhances the kinetic energy that enables the solute particles to split, increasing the velocity.

Colligative properties

• Colligative properties depend on the amount of solute particles present in a solution. That property relates to concentration and does not depend on the nature of its molecules.

Steam Pressure Diminution

• In closed containers which liquid molecules are traveling, the movement of such will produce collisions that exchange energy unevenly.
• Applying heat breaks down intermolecular forces to produce steam or a gaseous form.
• Pressure from vapour increases, and reaches equilibrium.
• Applying solute diminishes the speed of solvent particles (mostly water), reducing the concentration.
• Understand that volatile solutes easily reach gas or vapour state when air is present (such as gas).

Temperature

• Since steam pressure is lower for a fluid, one can infer that it possesses a higher boiling point.
• Decreasing steam pressure entails that higher temperatures are necessary.
• Decreasing vapour compression will increase the boiling temperature by any amount.

Freezing point and Osmosis

• Diminishing vapour compression has other effects, like reducing said water’s freezing point.
• Materials and membrane-based substances are semipermeable if they allow only some types of particle (such as steam from water, not solids or other types of ion).
• A fluid passes through a semipermeable by moving from low to high concentrates, creating an osmotic pressure until diffusion stops.
• Osmosis passes solvents through a permeable, but not solutes, moving increasingly dilute chemicals towards concentrated ones.

Importance of substances for beings

• Dissolutions are integral to existence.
• Roots intake water via osmosis, alongside intercellular exchange over cellular membranes.
• Marine composition contains dissolved gases, of which include nitrogen, oxygen, and carbon dioxide.
• Carbon reacts during photosynthesis and carbonates form structures in calcareous organisms.
• Sea-water promotes flotation because its density increases, reducing the need for skeletal components.
• Osmotic pressure increases the absorption of water from the roots to the leaves.
• Kidneys within fresh-watered fish can extract the salts to obtain pure water.
• Marine wildlife must extract and recycle water, losing it to osmosis.

Daily Uses

• A number of household cleaners are available as simple dissolutions, such as cleaning solutions and laundry detergent.
• Combining salt and ice makes for homemade sorbet.
• Antifreeze inhibits waters within a car radiator, especially in cold regions due to the need to decrease freezing point.
• Salinization prevents the formation of contaminants, therefore is important for food storage.
• Hypertension in marmalades and fruits keeps microorganisms out with concentrated sugar solutions.
• Altitude sickness reduces oxygen contents in the blood due to reduced air-pressure.
• Many materials like utensils and cosmetic products are possible through dissolutions.
• Modern medicine uses injections, serums, and pastes in saline solutions.
• Agriculture needs solutions/dissolutions like fertilizer, pesticides, and food stuff.
• Increased pressure at great depths causes nitrogen to resolve into the bodies of divers leading to euphoria.