Water Properties and Polarity

• The human body is 60 to 70% water
• Water is vital for life and its properties are important for biology at all levels

Properties of Water Articles

• Solvent properties of water: Explains how water dissolves polar and charged molecules
• Cohesion and adhesion of water: Describes how water sticks to itself (cohesion) and other molecules (adhesion)
• Specific heat, heat of vaporization, and density of water: Describes the high heat capacity/vaporization, and the lower density of solid versus liquid water

Polarity of Water Molecules

• Water's chemical behavior is tied to its molecular structure
• A molecule of water has two hydrogen atoms bonded to an oxygen atom, resulting in a bent shape
• Oxygen also has two pairs of unshared electrons
• The shared and unshared electron pairs repel each other
• The most stable arrangement of electron pairs is a tetrahedron, where the O-H bonds form two "legs"
• The angle between the O-H bonds is about 104.5°, less than a perfect tetrahedron (109°) due to repulsion from lone pairs
• Oxygen is more electronegative than hydrogen, so it hogs electrons

Charges in Water Molecules

• Oxygen has a partial negative charge, and hydrogen has a partial positive charge
• Water is a polar molecule because of its shape and polar covalent bonds

Hydrogen Bonding of Water Molecules

• Polarity causes water molecules to attract each other
• Positive hydrogen atoms connect with negative oxygen atoms
• Hydrogen bonds are weak interactions between a partially positive hydrogen and electronegative atom
• Hydrogen atoms in hydrogen bonds are attached to electronegative atoms like oxygen, nitrogen, or fluorine
• Water is also attracted to other polar molecules and ions

Hydrophilic and Hydrophobic Substances

• Hydrophilic substances interact with and dissolve in water
• Hydro means "water," and philic means "loving."
• Hydrophobic substances, like oils and fats, do not interact with water
• They separate from water and are called "phobic," meaning "fearing"
• Vinegar is essentially water with a little bit of acid

Universal Solvent

• Water is know as a "universal solvent" as it can dissolve a wide range of solutes
• A solvent is a substance that dissolves other molecules/compounds, known as solutes
• A homogenous mixture of solvent and solute is known as a solution

Aqueous

• Life's chemistry takes place in aqueous solutions

Polarity and Hydrogen Bonding

• Water makes an excellent solvent, dissolving many molecules
• Most life reactions occur in a watery environment inside of cells
• Water dissolves ions and polar molecules well, but not nonpolar molecules like oils
• Polar molecules have asymmetric internal charge distribution, leading to partial positive and negative regions

Water, Charges and Dissolvability

• Water interacts differently with charged and polar substances because of its polarity
• Water is polar; hydrogen is partially positive, oxygen is partially negative, has a bent shape
• Unequal charge distribution shows oxygen's electronegativity
• The shared electrons in O-H bonds spend more time with the O atom
• Partial positive and negative charges are represented by δ+ and δ− respectively

Electrostatic Interactions & Hydration Shells

• Water can form electrostatic interactions with polar molecules and ions, where positive attracts negative charges
• These interactions lead to hydration shell formation, where water molecules surround the solute
• Hydration shells allow particles to be dispersed evenly in water

Dissociation

• Dissociation allows the dissolution of ionic compounds, like table salt (NaCl) in water
• The crystal lattice of NaCl dissociates into Na+ and Cl− ions when stirred into water
• Dissociation is when a compound/molecule breaks apart to form ions
• Water molecules form hydration shells around the Na+ and Cl− ions
• Na+ ions are surrounded by partial negative charges from the oxygen ends of water molecules
• Cl− ions are surrounded by partial positive charges from the hydrogen ends, resulting in hydration shells around all ions

Nonpolar Dissolvability

• Nonpolar molecules like fats and oils have no partial charges, so they don't interact with water or form hydration shells
• Nonpolar substances stay separate, forming layers/droplets when added to water

Cohesion

• Cohesion refers to the attracion of molecules for other molecules of the same kind
• Water molecules have strong cohesive forces from the ability to form hydrogen bonds with each other
• Responsible for surface tension, where a liquid's surface resists rupture under tension or stress

Cohesive Forces

• Water molecules at the surface (water-air interface) will form hydrogen bonds with their neighbors
• Surface water molecules only bond with neighbors, stronger bonds form
• Surface tension causes water to form spherical droplets, supporting small objects

Adhesion of Water

• Adhesion attracts molecules of one kind for molecules of a different kind
• Adhesion enables water to "climb" upwards; known as capillary action
• Capillary action depends on the attraction between water molecules and the glass walls of the tube (adhesion), as well as on interactions between water molecules (cohesion)
• Water molecules are more attracted the glass than each other
• Glass is more polar than water molecules
• The curved surface formed by a liquid in a cylinder or tube is called a meniscus

Cohesive and Adhesive Forces

• Important for life
• Include the movement of water to the tops of trees, drainage of tears from tear ducts
• Cohesion allows water striders, insects that rely on surface tension, to stay afloat on water

Specific Heat

• Water has unique characteristics in all three states: solid, liquid, and gas
• Has unique characteristics due to ability to hydrogen bond
• Important key to biology because living things have high water content

Bonds forming and breaking

• In liquid water, hydrogen bonds constantly form and break as water molecules slide past each other
• Bond breaking is driven by kinetic energy of water molecules, which comes from system heat
• Hydrogen bonds break completely as heat rises
• Allows water molecules to escape into the air as gas (water vapor/steam)

Density of Water and Ice

• Water freezes -> hydrogen bonds maintain water molecules in a crystal structure
• Ice has lower density because hydrogen bonds cause water molecules to pushed less close together than in liquid water
• Most other liquids solidify, and have reduced kinetic (motion) energy, that allows molecules to pack more tightly than in liquid form

Heat Capacity of Water

• Liquid water needs a lot of heat to increase temperature because heat must break hydrogen bonds between molecules
• Water has a high specific heat capacity
• Specific heat capacity = heat amount neded to raise the temp of 1 gram of a substance by one degree Celcius
• Calorie = amoutn of heat needed to raise the temperature of 1 g water by 1°C

High Heat Capacity for Warm Blooded Animals

• Water minimizes changes in temperature due to its high heat capacity, which is about 5x great than sand
• Water also distributes heat in warm-blooded animals.

Heat of Vaporization of Water

• Changing water from liquid to gas (vaporization) requires a great amount of heat
• Hydrogen bonds must be broken for molecules to fly off as gas
• Water has a high heat of vaporization, the energy needed to change one gram of a liquid substance to a gas at constant temperature
• Water's heat of vaporization is ~ 540 cal/g at 100 °C
• Some water molecules with high kinetic energy can escape from surface at lower temperatures

Evaporative Cooling

• Evaporative cooling (the cooling of a surface) occurs as water molecules evaporate
• The molecules with the highest kinetic energy are lost
• Sweat evaporation, which is ~99% water, cools organisms to maintan steady temperature

Acids and Bases

• Acids/bases are commonly known as sour/slippery things
• An acidic solution has a high hydrogen ion (H+) concentration compared to pure water
• A basic solution has a low H+ concentration compared to pure water

Autoionization of Water

• Autoionization allows hydrogen ions to spontaneously generate in pure water by dissociating (ionizing) water molecules
• [H2O(l) ⇌ H+(aq) + OH−(aq)]
• Dissociation makes equal numbers of hydrogen (H+) ions and hydroxide (OH−) ions
• Hydrogen ions are transferred to a neighboring water molecule to form hydronium ions (H3O+)
• The concentration of hydrogen ions produced by dissociation in pure water is 1 × 10−7 M (moles per liter of water)
• Autoionized water molecules are a small fraction of total molecules in pure water

Solutions as acids or bases

• Solutions are divided into acids or based on their hydrogen ion concentration compared to pure water
• Acidic solutions have a higher H+ concentration than water, greater than 1 × 10−7 M
• Basic (alkaline) solutions have a lower H+ concentration, less than 1 × 10−7 M
• pH expresses a solution's hydrogen ion concentration
• Formual: pH = −log10​[H+]
• Pure water has a pH value of 7.0, also known as neutral pH
• Human blood has a pH value close to neutral
• An acid elevates H+ concentration
• Acids donate hydrogen

pH and Acids/Bases

• Stronger acids quickly dissociate, producing H+
• Hydrochloric acid (HCl) completely dissociates into hydrogen and chloride ions in water, therefore considered a strong acid
• Acids that don't readily dissociate in water are weak acids
• Strong bases like sodium hydroxide (NaOH) completely dissociate in water, releasing hydroxide ions
• pH scale ranks solutions in terms of acidity or basicity (alkalinity)

pH Scale

• A change of 1 pH unit corresponds to a ten-fold change in H+ ion concentration due to the pH scale's logarithmic nature
• The pH scale is often said to range from 0-14, where values below pH 7.0 is acidic and above pH 7.0 is alkaline, or basic
• Human cells are ~6.8 pH and blood is ~7.4 pH
• Internal stomach environment has extreme pH values are considered generally unfavorable for life

Buffers

• Organisms maintain pH within a narrow range to survive
• Human blood maintains a pH ~7.4
• Buffers resist changes in pH,maintaining stable H+ concentrations in biological systems
• More H+ ions, buffer -> absorb
• Less H+ ions, buffer -> donate
• A buffer consists of a conjugate acid-base pair that differ in the presence/absence of a proton
• Bicarbonate ions absorb H+ to form carbonic acid, pushed to the right
• Low H+ ion, carbonic acid turns into bicarbonate, donating H+ ions to the solution
• The body's pH would be to severe and variable without a certain buffer system