Biology Notes Ch. 3 Water

What allows for water to have unique properties?

• Unique properties can be attributed to its structure and interactions of its molecules

• Water is shaped like a V and is held together through polar covalent bonds as O is more electronegative than hydrogen, thus meaning the overall charge is unevenly distributed.

• Properties arise from the attractions between the oppositely partially charged water molecules with each other

• In liquid form its hydrogen bonds are very fragile and is only abt 1/20th as strong as a covalent bonds, thus they break and reform with great frequency.

Cohesion & Surface Tension v. Adhesion

• Cohesion: when hydrogen bonds hold substances together, creating more linkages that makes water more structured

  • This allows for the transport of water and dissolved nutrients against gravity in plants

• Surface Tension: a measure of how difficult it is to stretch or break the surface of a liquid. At the air-water interface is an ordered arrangement of water molecules, hydrogen bonded to one another and the water below, but not the air above. This asymmetry gives water an unusually high surface tension

• Adhesion: the clinging of one substance to another

  • Adhesion of water by hydrogen bonds to the molecules of cell walls help counter the downward pull of gravity

Four emergent properties of water contributing to Earth’s suitability for life

• Cohesive behavior

• Ability to moderate temperature

• Expansion upon freezing

• Versatility as a solvent

How does water moderate air temperature?

• Absorbs heat from air that is warmer and releases stored heat to air that is cooler

• Effective as a heat bank because it can absorb or release a relatively large amount of heat with only a slight change to its own temperature

Kinetic energy

• Energy of motion

• Atoms and molecules have kinetic energy because they are always moving, although no necessarily in any particular direction

• The faster a molecule moves the more kinetic energy

• Kinetic energy associated with the random movement of atoms or molecules is called thermal energy

Temperature vs. Thermal energy

• T: represents the average kinetic energy of the molecules in a body of matter, regardless of volume

• TE: total kinetic energy, thus depends on the matter’s volume

  • Example: a pot of boiling coffee has a higher temperature than a pool, but it has less thermal energy than a swimming pool because a swimming pool is much bigger (with more volume)

Heat

• Thermal energy is transfer from one body of matter to another

• One convenient unit of heat is calorie (raise the temperature of 1 g of water by 1 degree C, kcal is 1000cals and is the quantity of heat required to raise the temperature of 1 kg of water by 1C)

• Another unit is Joule which is .239 cal or one calorie equals 4.184 J

Specific Heat

• Defined as the amount of heat that must be absorbed or lost for 1 g of the substance to change its temperature by 1 degree C

• Specific heat of water is defined as 1 calorie per gram and per degree Celsius aka 1 cal/ (g x C)

• Water has an unusually high specific heat compared to other substances

• Water has a higher specific heat so it will change its temperature less than other liquids when it absorbs or loses a given amount of heat.

Think of it as a measure of how well a substance resists changing its temperature when it absorbs or releases heat.

• When it does change its temperature, it absorbs or loses a relatively large quantity of heat for each degree of change

Why does water have high specific heat?

• Hydrogen bonding

  • Heat must be absorbed in order to break hydrogen bonds; but the same token, heat is released when hydrogen bonds form

  • 1 cal doesn’t really do much because its heat is used to disrupt hydrogen bonds before the water molecules can begin moving faster

What is the relevance of water’s high specific heat to life on Earth?

• A large body of water can absorb and store a huge amount of heat from the sun in the daytime and during summer while warming up only a few degrees

• At night and during winter, the gradually cooling water can warm the air

  • Moderates air temperature in coastal areas

• Stabilizes ocean temperature, creating a favorable environment for marine life

• Because organisms are primarily made of water, they are better able to resist changes in their own temperature than if they were made of a liquid with a lower specific heat.

Evaporation

• Transformation from a liquid to a gas

• Even at low temperatures, the speediest molecules can escape into the air

• At room temperature a glass of water will eventually evaporate, just slower compared to if the liquid is heated as the average KE of the molecules increases and they “leave” more quickly

Heat of Vaporization

• the quantity of heat a liquid must absorb for 1 g of it to be converted from the liquid to gas state

• water has a high heat of vaporization relative to most other liquids for the same reason that they have a high specific heat

  • when heat is absorbed much of that energy is used to break hydrogen bond before it can be used to increase the kinetic energy

• helps moderate Earth’s climate, as a considerable amount of solar heat absorbed by tropical seas is consumed during the evaporation of surface water

Evaporative cooling

• as liquid evaporates, the surface of the liquid that remains behind cools down because the “hottest” molecules, those with the greatest kinetic energy, are the most likely to leave as a gas.

• contributes to the stability of temperature in lakes and ponds and also provides a mechanism that prevents terrestrial organisms from overheating

  • Evaporation of sweat from human skin dissipates body heat and helps prevent overheating on a hot day or during exercise

Why is ice less dense than liquid water?

• the cause is hydrogen bonding, as when water temperature is about 4 degrees C, water behaves like other liquids, expanding as it warms and contracting as it cools. As the temperature falls from 4 degrees C to ) degrees C, water begins to freeze because more and more of its molecules are moving too slowly to break hydrogen bonds.

• At 0 degrees C the molecules become locked into a crystalline latice, each water molecule hydrogen bonded to four partners. The hydrogen bonds keep the molecules at arm’s length

  • Far enough apart to make ice bout 10% less dense than liquid water at 4 degrees

Thermal contraction (opposing force)

• as water cools from higher temperatures, the molecules start to slow down and move closer together

Hydrogen bonding as an expansion force

• water molecules form hydrogen bonds with each other and as water cools these bonds become more stable and create a more organized and open crystalline-like network that take sup more space.

Why is water most dense at 4 degrees C?

Due to a unique balance between two competing phenomena: thermal contraction and hydrogen bond expansion. As water cools, molecules slow down, drawing closer (thermal contraction), but also form stable hydrogen bond networks that take up more space. At 4°C, the structure-forming hydrogen bonds expand enough to counteract the thermal contraction, creating a maximum density.

Temperatures above 4 degrees C have less density because…

• As the crystal collapses, the ice melts and molecules have fewer hydrogen bonds, allowing them to slip closer together

Why is it good that ice is less dense than water?

• If ice sank, then eventually ponds, lakes, and even oceans could freeze solid, making life as we know it impossible on Earth

• When deep body of water cools, the ice floats, insulating the water below

Solution

• A liquid that is completely homogenous mixture of two or more substances

• an aqueous solution is one in which the solute is dissolved in water; water is the solvent

Solvent vs. Solute

• The dissolving agent is the solvent

• the substance that is dissolved is the solute

• Together they make a solution

Hydration Shell

• the sphere of water molecules around each dissolved ion

Why is water a versatile solvent and can only ionic compounds dissolve in water?

• A compound does not need to be ionic to dissolve in water; many compounds made up of nonionic polar molecules are water-soluble

• Such compounds dissolve when water molecules surround each of the solute molecules, forming hydrogen bonds with them.

• Even molecules as large as proteins can dissolve in water if they have ionic and polar regions on their surface

Hydrophilic vs. Hydrophobic

• Substances can be hydrophilic without actually dissolving

  • For example, some molecules in cells are so large that they do not dissolve

• Substances that are nonionic and non-polar or cannot form hydrogen bonds seem to repel water

Molecular mass

• The sum of all the masses of each element in a compound

Molarity

• the number of moles of solute per liter of solution — is the unit of concentration most often used by biologists for aqueous solutions

• Numbers of moles of solute/liter of solution

Water

• water is the biological medium on Earth

• All living organisms require water more than any other substance

• Most cells are surrounded by water, and cells themselves are 70-95% water

• The abundance of water is the main reason the Earth is habitable

Polar Molecule

the opposite ends have opposite charges which allows hydrogen bonds

What do chemical reactions depend on?

Collisions of molecules and therefore on the concentration of solutes in an aqueous solution

Avogadro’s number

6.02×10²³ and dalton were defined such that 6.02×10²³ daltons = 1 g

What ions can arise from participating in a hydrogen bond?

• A hydrogen atom can leave its electron behind which ill make it a hydrogen ion (H+)

• The water molecule that lost a proton is now a hydroxide ion (OH-)

• The proton binds to another water molecule which makes it H3O+

• Often H+ is just used to represent H3O+

How common are these ions in pure water? And are they important?

• This dissociation is a reversible reaction and is at a state of dynamic equilibrium when the water molecules dissociate at the same rate that they are being reformed

• In pure water there is only ten-millionth of a mole of hydrogen ions per liter (but this is still over 60000 trillion of each ion in a liter)

• They are very important because these ions are very reactive, and even a slight change int heir concentrations can drastically affect a cell’s proteins and other complex molecules

• Even though in pure water the concentrations of H+ and OH- are equal, this changes when you add different kinds of solutes, mainly acids and bases

Acid

• A substance that increases the hydrogen ion concentration of a solution

• For example if you add hydrochloric acid (HCL) to water you get HCL → H+ + Cl-

• Lowest numbers on the pH scale

Base

• a substance that reduces the hydrogen ion concentration of a solution

• some bases reduce the H+ concentration directly by accepting the hydrogen ions

• For example ammonia (NH3) acts as a base when NH3 + H+ → NH4+

• other bases reduce the H+ concentration indirectly by dissociating form hydroxide ions

• For example NaOH → Na+ + OH-

• Lower on the pH scale

Arrows Use

• A one way arrow for reactions is used when you are using a strong base or strong acid

• A two way arrow indicates a reaction is reversible and utilizes weak acids or bases

What is the concentration of H+ and OH- at 25 C in aqueous solutions?

• the concentration is a constant at 10^-14

• Is written as [H+][OH-] = 10^-14

• brackets indicate molar concentration

• in neutral solutions [H+]=10^-7

• in neutral solutions [OH-]=10^-7

• Example: if you add an acid to a solution which increases [H+] = 10^-5 then the [OH-]=10^-9 bc 10^-5 times 10^-9

pH

• defined as the negative logarithm (base 10) of the H+ concentration

• pH=-log[H+]

• so in a neutral aqueous solution, [H+] is 10 ^-7 so -log10^-7=-(-7)=7

• pH decreases as H+ concentration increases

Remember that each pH unit represents a ———— difference in H+ and OH- concentrations

• Tenfold

• It is this mathematical feature that makes the pH scale so compact. A solution of pH 3 is not twice as acidic as a solution of pH 6, but 1,000 times (10 x 10 x 10) more acidic. When the pH of a solution changes slightly, the actual concentrations of H+ and OH- in the solution change substantially

Buffers

• the presence of substances called buffers allows biological fluids to maintain a relatively constant pH despite the addition of acids or bases

• a substance that minimizes changes in the concentrations of H+ and OH- in a solution

• IT does so by accepting hydrogen ions from the solution when they are in excess and donating hydrogen ions to the solution when they have been depleted

• Most buffer solutions contain a weak acid and its corresponding base, which combine reversibly with hydrogen ions

What is the buffer in blood?

• Carbonic acid-bicarbonate

Ocean adifiction

• when carbon dioxide dissolves in seawater, it reacts with water to form carbonic acid, which lowers ocean pH.