Lec 1.4 & 1.5 - solutions & pH

Biol 121, K. O'Neil, Instructor 1/10/2010

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Biology 121Lectures 1.4 & 1.5

Water · Solutions · pH

Water is an extremely important molecule in biology

• As we complete our discussion of bonding, it is appropriate to look specifically at the bonding pattern of water, the most vital molecule for life and living systems.

Water is important

• Water makes up 70% of our body weight, fills all our cells

• Provides a solution for molecules to interact

• Water is used during photosynthesis to produce O2 that we breathe

• Allows our bodies to maintain optimum temperature range – prevents heating up or cooling down too far

Water is H2O

• Water is a polar covalent molecule between 1 O and 2 H atoms

• POLAR

• Draw the Lewis dot diagram for water

Water is H2O

• Polar covalent molecule

H O H


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3D structure of water

• When we draw the structure of water, it appears to be a rather planar molecule that lies flat.

• NOT TRUE!

• Atoms and molecules actually occupy space in 3 dimensions

• Water is actually a ‘bent’ molecule in 3D

3D structure of water

• ‘Bent’ molecule

• 2 H atoms on one end

• O atom on opposite end

• Electronegativity?– O is very electronegative,

H is much less

• δ- at O end, δ+ at H end

• Very POLAR

• Results in extensive H-bonding

Water Molecules Interact

• So, there are an enormous number of partial charges present in a water solution.

• To obtain neutrality and increased stability, the molecules of water spontaneously arrange themselves so that partial charges line up toward their opposite charges and neutralize one another.

• The intermolecular interactions that occur between separate water molecules are referred to as hydrogen bonds (H-bonds).

Two Bonds in Water

• An individual water molecule is held together by polar covalent bonds between one O and two H. These bonds are intramolecular, strong and relatively permanent.

Two Bonds in Water

• A collection of water molecules are held together by H bonds between molecules. These bonds are intermolecular, weak and transient, but large in number.

Water – extremely polar, extremely imporant

• Water’s polarity and H-bonding tendencies give it its extremely important properties

1. Cohesion

2. Adhesion

3. High specific heat

4. Excellent solvent


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1. Cohesion

• Attraction between water molecules

• H-bonds hold water molecules together so they don’t disperse– Drop of water is held in shape by

H-bonds, cohesion

– Drop on table is ‘tall’

– Fill a cup past the top, rounded

1. Cohesion

• Surface tension –molecules are more attracted to each other than to the air

• Water molecules form a layer on the top surface by crowding together and being pulled down

1. Cohesion

• Transpiration is the process vascular plants use to pull water up from the soil to their stems and leaves.

• These plants have ‘pipelines’ of vascular tissue that carry water (and solutes) upwards, against gravity, from the roots to the leaves.

• Due to water’s H-bonds, as water molecules in the leaves evaporate, cohesion pulls the next available water molecule to the leaf surface.

2. Adhesion

• Water’s interaction with other surfaces

– Meniscus in a graduated cylinder – water interacts with the glass. All liquids do not behave this way.

– Paper towel ‘wicks’ water upward

– Capillary action pulls water up narrow passages because water interacts with inner surfaces

3. High Specific Heat

• Water maintains relatively stable temperature, requires a lot of energy to increase temperature

• Specific heat = amount of energy needed to raise 1 g of substance 1 °C

– specific heat of water = 1 cal/g/1 °C

– specific heat of alcohol = 0.59 cal/g/1 °C

– specific heat of sand = 0.2 cal/g/1 °C

• Can this explain why the sand burns your feet at the beach but the water still feels cool?

What is temperature anyway?

• Temperature is a measure of how fast molecules are moving – kinetic energy

• Adding energy makes molecules move faster = higher temperature

– Sun, burner on the stove

• Reducing energy makes molecules move more slowly = lower temperature

– Shade, refrigerator

http://upload.wikimedia.org/wikipedia/commons/2/23/Thermally_Agitated_Molecule.gif


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3. High Specific Heat

• Heating water – must break H-bonds to let water move faster

– Breaking bonds takes energy, less energy left to make molecules move

– Metabolism makes a lot of heat –water keeps us from overheating

• Cooling water – slowing down molecules allows more H-bonding

• Molecules in ice are held further apart less dense, floats in water

4. Water is an excellent solvent

• Many solutes (salt, sugar, alcohol, etc.) dissolve in water

• Makes a solution that holds many kinds of molecules

– Different things dissolved in water can interact with each other – reactions, life functions


• Solution – a mix of 2 or more components

– base component = solvent, a liquid, usually water

– other component(s) = solute that dissolves in the solvent, usually a solid

– solvent dissolves the solute by surrounding each molecule to keep them from interacting

• polar solvents dissolve polar solutes,

• non-polar solvents dissolve non-polar solutes


• water is polar - dissolves polar sugar and ionic (NaCl), not non-polar (oil)

– polar solute = hydrophilic (water-loving)

– non-polar solutes = hydrophobic (water-fearing)

– polar covalent molecules – dissolve in water

• δ- water end is attracted to δ+ solute end, δ+ water end is attracted to δ– solute end


• Water forms a ‘hydration sphere’ around each solute molecule

• Ionic molecules dissolve and dissociate in water – do not stay intact

– Solution will conduct electricity electrolyte


• Water forms a ‘hydration sphere’ around each solute molecule

• Covalent solute molecules are separated by water, but each molecule stays intact

– Soln does not conduct electricity non-electrolyte


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Summary of Water Properties• Learning Goal: Be able to list and describe the characteristics

and properties of water that we have discussed.

One O and two H covalently bonded Tetrahedral atoms

make it bent Polarity of covalent bonds causes H-bonds

to form

High specific heat stabilizes temperatures

Cohesion causes surface tension and permits

transpiration

Biology 121Week 3

Water · Solutions · pH

Solvent properties of Water

• Another unique and useful property of water……..it is an excellent solvent.

• Many many substances will dissolve in water.

• Chemical processes occur best in solutions, so having 60-70% or our bodies composed of the world’s best solvent is very useful….because we need our cells to perform chemical processes.

Making solutions

• Solutions are specific types of mixtures of two or more molecules– Saline solution is one

example, a mixture of salt and water

– Dextrose IV solutions are another example, mixtures of sugar and water

Making solutions

• The component in larger proportion is called the solvent. Most of the solutions we will discuss will be aqueous solutions, meaning the solvent is H2O.

• The component in smaller quantity is called the solute. Most of the solutes we will discuss are solids, which seem to ‘disappear’ in water when dissolved.

• Gases and other liquids may also be dissolved in water.

Making solutions

• In a solution, each solutemolecule actually becomes completely surrounded by solvent molecules

• In H2O solutions, we call this a hydration shell around the solute.

• This ball and stick model represents a red and white organic molecule (like dextrose) completely surrounded by blue water molecules


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Like Dissolves Like

• Although water is often termed the “universal” solvent, it really is not. Only certain substances will dissolve in water.

• Examples of water-soluble substances are sugar and salt.

• An example of a non-water soluble substance is oil (think vinegar and oil).

• Because water is so vital to life, we characterize all molecules with respect to their ability to dissolve in water.

Like Dissolves Like• Solubility of substances is

completely predictable

• Water – polar covalent molecule, dissolves other polar covalently bonded substances (like glucose)

• Ionic compounds (like NaCl) are also “like” water – they are polar.

• Nonpolar covalently bonded molecules are NOT like water, and will not dissolve in water

Electrolytes

• So polar covalent AND ionic compounds both dissolve in water. Do they behave similarly in solution? Let’s take a close-up look at how hydration shells form.

Electrolytes

• When sugar is added to water it dissolves. Each individual sugar molecule, carrying partial charges, becomes completely surrounded by water. The partial charges on sugar interact with the partial charges in water.

• Sugar is hydrophilic (a “water lover”).

Electrolytes

• When salt (NaCl) is added to water it also dissolves. But ionic bonds are different than covalent bonds -- the electrons are donated, not shared.

• When ionic compounds dissolve, the ions separate from each other, and each individual ion, carrying full charges, becomes completely surrounded by water.

• The full charges on the ions interact with the partial charges in water.

Electrolytes

• Ionic compounds are also hydrophilic (“water lovers”), but in a different way than sugar.

– Polar covalent molecules dissolve in water, but do not dissociate

– Ionic compounds dissolve AND dissociate in water


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Electrolytes

• Ionic compounds carry full charges that are neutralized when the compound is a solid.

• When the ions dissociate into water, the charges are no longer neutralized. That’s why water containing ions carry electricity.

• So – no swimming in thunderstorms!

Electrolytes and Nonelectrolytes

• In water, substances are either

– soluble - polar covalent or ionic

– not soluble - nonpolar covalent

• Water-soluble substances can be further subdivided into

– Polar covalent nonelectrolytes -those that just dissolve

– Polar ionic electrolytes - those that dissolve and dissociate

Biology 121Lecture 1.5

pH

Buffers

pH and water

• Pure water dissociates slightly into H+ and OH-. – O is so electronegative that it can pull e- away from H+

• Each ion is present in equal and small concentrations– H2O ↔ *H+] + [OH-]

• In pure water, each ion is present in equal concentrations– [H+] = [OH-] = 1 x 10-7 M (neutral)

• Kw = dissociation constant of water– Kw = [H+] · [OH-] = 1 x 10-14

• Water is an electrolyte, right?

Three Kinds of Electrolytes• It is informative to classify

electrolytes into subsets that dissociate into

1.) a H+ ion and something else.

2.) a OH- ion and something else

3.) neither H+ ions nor OH-.

• These 3 categories are referred to, respectively, as

1.) acids,

2.) bases, and

3.) salts.

1. Acids are Electrolytes

• HCl is an acid. When dissolved in water, HCldissociates into a H+ cation and a Cl- anion.

• Acid ↔ H+ + anion

• HCl ↔ H+ + Cl-

• Increases [H+]

• Kw is constant, so when [H+] goes up, [OH-] goes down


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1. Acids are Electrolytes

• Vinegar (acetic acid) is another example of an acid.

• Unlike HCl, acetic acid does not dissociate completely.

2. Bases are Another Type of Electrolyte

• NaOH is a base. A base is an electrolyte that dissociates into OH- and a cation.

• OH- is a more complex ion called hydroxide.

• Base ↔ cation + OH-

• NaOH ↔ Na+ + OH-

• Increases [OH-]

• Kw is constant, so when[OH-] goes up, [H+] goes down

3. Salts are another type of Electrolyte

• NaCl is a salt. A salt is an electrolyte that dissociates into a cation and an anion.

• Salt ↔ cation + anion

• NaCl ↔ Na+ + Cl-

• Does not affect [H+] or [OH-]

• Neither an acid nor a base

Summary - Electrolytes

• Substances are either – soluble - polar covalent or ionic– insoluble - nonpolar covalent in

water.

• Water-soluble substances are either – nonelectrolytes - polar covalent– electrolytes - ionic

• Electrolytes are either – acids (H+), – bases (OH-) or – salts (neither).

Relationship between solubility and polarity

Electrolytes

Polarity/Bond Type

Solubility

All substances

Water-soluble

Polar Covalent

Ionic

Acid Base Salt

Water insoluble

Non-polar covalent

Significance of Hydrogen Ions

• Why did we devise an entire classification scheme around H+ and OH- ions?


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Significance of H+ Ions

• Because when you put a hydrogen H+ ion and a hydroxide OH- ion together…you get WATER.

Low, but Predictable, Dissociation Rate

• In a beaker of water, a tiny number of H2O molecules are dissociated.

• In pure water, we will find 1 x 10-7

moles of H+ per liter of water– [H+] = 1 x 10-7 M

• In pure water, there must be exactly the same amount of OH- as H+. For every H+ that “pops” off a water, there is an OH- left, a 1:1 ratio.

• So there are 1 x 10-7 moles per liter of OH- also. – [OH-] = 1 X 10-7 M

Low, but Predictable, Dissociation Rate

• [H+] and [OH-] is constant in pure water, and is represented by Kw(dissociation constant of water)

• Kw = [H+] x [OH-]• In pure water, we know those

values.• Kw= (1 x 10-7) · (1 x 10-7) = 1 x 10-14

Kw Never, Ever Changes

• Kw is called a CONSTANT because it remains unchanged no matter what.

• Kw for any solution has a value of 1 x 10-14

Adding Acids to Water

• So, what happens if we have a beaker of water and we add an acid, HCl, to that beaker?


• HCl dissociates into H+ and Cl-. So now we have MOREH+ than we started with. The concentration of H+ is higher than it was.

• Let’s add enough HCl to increase [H+] to 10-3 M.


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Kw Never, Ever Changes

• So if [H+] goes up then [OH-] must go down.

• Specifically, Kw = 10-14, so if [H+] is 10-3, then [OH-] must be 10-11 M.


• When you add an acid to water, the [H+] goes up and the [OH-] goes down.

Adding Bases to Water

• Let’s add a base, like NaOH, to that beaker.

• NaOH dissociates into Na+

and OH-. So now we have more OH- than we started with. Now the concentration of OH- is higher than it was. Let’s add enough NaOH to increase [OH-] to 10-5 M.

Adding Bases to Water

• Again, Kw is called a CONSTANT because it never changes.

• So if [OH-] goes up then [H+] must go down.

• Kw = 10-14, so if [OH-] is 10-5, then [H+] must be 10-9.

Adding Salts to Water

• Let’s try this with a salt, NaCl.

• NaCl dissociates into Na+

and Cl- ions. Neither H+

nor OH- are changed from the original water.

• The [H+] is still 10-7, and so is the [OH-].

pH - a convenient way to communicate H+ concentration

• pH = -log [H+]. (log base 10 or log10)

• pH = - exponent of the [H+] concentration.

• Measured on a scale from 0 to 14

– Neutral = 7, acidic <7, basic >7


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• In pure water, [H+] is 10-7. So pH = 7. On the pH scale, that is exactly in the middle, and water is considered neutral.


• If we add HCl, and the [H+] goes up to 10-3 like in our previous example, the pH changes to 3.

• So when [H+] goes up, pH goes down.


• If we add NaOH, and the [OH-] goes up to 10-5 and the [H+] goes down to 10-9, keeping the Kw constant. The new pH is 9.

• So when [H+] goes down, pH goes up.


• pH 9 is on the high end of the pH scale, and solutions that have high pH are considered basic (also called alkaline).

• Things like ammonia and lye, which are caustic are basic.

• Blood is also very slightly basic.

• Acids are on the low end of the pH scale.

• Things like citrus juices, vinegar, stomach contents are acidic.


• When you add NaCl, or any other salt, to pure water, the [H+] and [OH-] remain the same, and the pH is still neutral.

Summary of pH

• In summary, water dissociates into H+

and OH- at a small, but predictable, rate.

• Kw is a constant 10-14 for water.• If you add H+ to water, the OH- will go

down. If you add OH- , the H+ will go down.

• Salts don’t alter H+ or OH-.• For convenience [H+] is reported via

the pH scale, and when [H+] goes up, pH goes down.

• pH = -log [H+] • 1 unit pH change reflects a 10-fold

change in [H+]


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Problems

Substance [H+] pH [OH-]

Pure water 10-7 mol/L neutral

lemon juice 2 acidic

ammonia 10-11 mol/L basic

bleach 9

coffee 5

vinegar 10-3 mol/L

Biology 121Lecture 1.5

pH

Buffers

Strong acids v. weak acids

• Strong acids dissociate completely (100%) in water

– hydrochloric acid HCl → H+ + Cl-

– hydrobromic acid HBr → H+ + Br-

• Weak acids only partially dissociate in water

– carbonic acid H2CO3 ↔ H+ + HCO3-

(bicarbonate)

– phosphoric acid H3PO4 ↔ H+ + H2PO4-

Strong bases v. weak bases

• Strong bases dissociate completely (100%) in water– sodium hydroxide NaOH → Na+ + OH-

– potassium hydroxide KOH → K+ + OH-

• some bases increase [OH-] by accepting H+ from water

• H+ acceptors are considered bases, even if they don’t directly donate OH-– ammonia NH3 + H2O ↔ NH4

+ + OH-

– H+ acceptors remove H+ from solution like a ‘sponge’

• weak bases only partially dissociate (or only partially take H+ from water)

Salts do not affect pH

• A salt dissociates into ions other than H+ or OH-, no change in [H+] or [OH-]

– NaCl ↔ Na+ + Cl-

– KCl ↔ K+ + Cl-

• Can make a salt by combining an acid and a base

– HCl + NaOH ↔ NaCl + H2O

– HCl + KOH ↔ KCl + H2O

Buffers – important in biology

• pH is very important – too high or too low will prevent cells from functioning

– Most cells are pH 7.2-7.4

– Some compartments of cells need higher or lower pH to complete their functions

– Buffers allow cells to maintain this narrow pH range


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Buffers

• Buffers interact with H+ and OH- to prevent acids or bases from changing the pH of a solution

• A buffer is a weak acid or weak base that does not dissociate completely– creates a mix of acid (H+ donor) and base (H+

acceptor)

– example: carbonic acid: H2CO3 ↔ H+ + HCO3-

– H2CO3 = acid, H+ donor HCO3- = base, H+ acceptor

• add base (OH-) to buffer: buffer donates an H+ to balance OH-, makes water

Buffers

• add acid (H+) to buffer: buffer accepts H+ to prevent a drop in pH– Example: carbonic acid H2CO3 ↔ H+ + HCO3

-

• Add base OH- + H2CO3 ↔ H2O + HCO3-

– (no change in [OH-] or [H+])

• Add acid H+ + HCO3- ↔ H2CO3

– (no change in [OH-] or [H+])

– Example: ammonia NH3 + H2O ↔ NH4+ + OH-

• Add base OH- + NH4+ ↔ NH3 + H2O

• Add acid H+ + NH3 ↔ NH4+

Buffer has at least one Ka

• Ka of the buffer indicates the pH that it tries to maintain

• Ka = pH where [H+

donor] = [H+ acceptor]

Buffer systems in our bodies

1. Some proteins can donate/accept H+ to buffer cytoplasm of cells

2. Phosphate buffers also work in cells– H2PO4

- ↔ H+ + HPO42-

– keeps pH near Ka = 6.86, good for cell compartments with pH 6.9-7.4

3. Bicarbonate buffers the blood by exchanging CO2 in the lungs – CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3

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Buffer systems

• Buffers are limited –they can only buffer a small amount of H+ or OH-

• Once buffer capacity is exceeded, pH changes quickly

Study Guide

• Give the equation for the dissociation of pure water, the concentration of H+ ([H+]) in pure water, and the dissociation constant of water (Kw).

• Define and give an example of an acid and know what happens to the [H+], [OH-], and pH when an acid is added to water

• Define and give an example of a base and know what happens to the [H+], [OH-], and pH when a base is added to water

• Define and give an example of a salt and know what happens to the [H+], [OH-], and pH when a salt is added to water

• Given the [H+] or pH or a solution, determine the [H+], [OH-], pH, and whether it is neutral, acidic, or basic

• Define and give an example of a weak acid and a weak base

• Understand that the partial dissociation of a weak acid of weak base produces a mixture of H+ donors and H+ acceptors

• Explain what happens when a base is added to a weak acid (use equations) and how this affects the pH

• Explain what happens when an acid is added to a weak acid (use equations) and how this affects the pH


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End of Section 1

• Exam #1 will cover from the beginning of the semester to this point…

Lec 1.4 & 1.5 - solutions & pH

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