Lec 1.4 & 1.5 - solutions & pH


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Lec 1.4 & 1.5 - solutions & pH

  1. 1. Biol 121, K. O'Neil, Instructor 1/10/2010 Water is an extremely important molecule in biology • As we complete our discussion of bonding, it is Biology 121 appropriate to look specifically at the bonding Lectures 1.4 & 1.5 pattern of water, the most vital molecule for life and Water · Solutions · pH 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 H2O • Water is a polar • Polar covalent molecule covalent molecule between 1 O and 2 H atoms • POLAR • Draw the Lewis dot H O H diagram for water 1
  2. 2. Biol 121, K. O'Neil, Instructor 1/10/2010 3D structure of water 3D structure of water • When we draw the structure of • ‘Bent’ molecule water, it appears to be a rather • 2 H atoms on one end planar molecule that lies flat. • O atom on opposite end • NOT TRUE! • Electronegativity? • Atoms and molecules actually – O is very electronegative, H is much less occupy space in 3 dimensions • δ- at O end, δ+ at H end • Water is actually a ‘bent’ • Very POLAR molecule in 3D • Results in extensive H- bonding Water Molecules Interact Two Bonds in Water • So, there are an enormous • An individual water number of partial charges present in a water solution. molecule is held together • To obtain neutrality and by polar covalent bonds increased stability, the between one O and two H. molecules of water spontaneously arrange These bonds are themselves so that partial intramolecular, strong and charges line up toward their opposite charges and neutralize relatively permanent. one another. • The intermolecular interactions that occur between separate water molecules are referred to as hydrogen bonds (H-bonds). Water – extremely polar, extremely Two Bonds in Water imporant • A collection of water • Water’s polarity and H- molecules are held together bonding tendencies give by H bonds between it its extremely important molecules. These bonds are intermolecular, weak and properties transient, but large in 1. Cohesion number. 2. Adhesion 3. High specific heat 4. Excellent solvent 2
  3. 3. Biol 121, K. O'Neil, Instructor 1/10/2010 1. Cohesion 1. Cohesion • Attraction between water • Surface tension – molecules molecules are more • H-bonds hold water molecules attracted to each other together so they don’t than to the air disperse • Water molecules form a – Drop of water is held in shape by H-bonds, cohesion layer on the top surface by – Drop on table is ‘tall’ crowding together and – Fill a cup past the top, rounded being pulled down 1. Cohesion 2. Adhesion • Transpiration is the process vascular • Water’s interaction with other plants use to pull water up from the surfaces soil to their stems and leaves. – Meniscus in a graduated • These plants have ‘pipelines’ of cylinder – water interacts with vascular tissue that carry water the glass. All liquids do not (and solutes) upwards, against behave this way. gravity, from the roots to the – Paper towel ‘wicks’ water leaves. upward • Due to water’s H-bonds, as water – Capillary action pulls water up molecules in the leaves evaporate, narrow passages because water cohesion pulls the next available interacts with inner surfaces water molecule to the leaf surface. 3. High Specific Heat What is temperature anyway? • Water maintains relatively stable temperature, • Temperature is a measure of requires a lot of energy to increase temperature how fast molecules are moving • Specific heat = amount of energy needed to raise 1 g – kinetic energy of substance 1 °C • Adding energy makes – specific heat of water = 1 cal/g/1 °C molecules move faster = higher – specific heat of alcohol = 0.59 cal/g/1 °C temperature – specific heat of sand = 0.2 cal/g/1 °C – Sun, burner on the stove • Can this explain why the sand burns your feet at the • Reducing energy makes beach but the water still feels cool? molecules move more slowly = lower temperature – Shade, refrigerator 3
  4. 4. Biol 121, K. O'Neil, Instructor 1/10/2010 3. High Specific Heat 4. Water is an excellent solvent • Heating water – must break H- • Many solutes (salt, sugar, alcohol, etc.) bonds to let water move faster dissolve in water – Breaking bonds takes energy, less energy left to make molecules move • Makes a solution that holds many kinds of – Metabolism makes a lot of heat – molecules water keeps us from overheating – Different things dissolved in water can interact • Cooling water – slowing down with each other – reactions, life functions molecules allows more H-bonding • Molecules in ice are held further apart  less dense, floats in water 4. Water is an excellent solvent 4. Water is an excellent solvent • Solution – a mix of 2 or more components • water is polar - dissolves polar sugar and ionic – base component = solvent, a liquid, usually water (NaCl), not non-polar (oil) – other component(s) = solute that dissolves in the – polar solute = hydrophilic (water-loving) solvent, usually a solid – non-polar solutes = hydrophobic (water-fearing) – solvent dissolves the solute by surrounding each – polar covalent molecules – dissolve in water molecule to keep them from interacting • δ- water end is attracted to δ+ solute end, δ+ • polar solvents dissolve polar solutes, water end is attracted to δ– solute end • non-polar solvents dissolve non-polar solutes 4. Water is an excellent solvent 4. Water is an excellent solvent • Water forms a ‘hydration • Water forms a ‘hydration sphere’ around each solute sphere’ around each solute molecule molecule • Covalent solute molecules • Ionic molecules dissolve and are separated by water, but dissociate in water – do not each molecule stays intact stay intact – Soln does not conduct – Solution will conduct electricity  non-electrolyte electricity  electrolyte 4
  5. 5. Biol 121, K. O'Neil, Instructor 1/10/2010 Summary of Water Properties • Learning Goal: Be able to list and describe the characteristics and properties of water that we have discussed. Biology 121 One O and two H covalently Week 3 bonded Tetrahedral atoms make it bent Polarity of covalent bonds causes H-bonds to form Water · Solutions · pH High specific heat Cohesion causes surface stabilizes temperatures tension and permits transpiration Solvent properties of Water Making solutions • Another unique and useful • Solutions are specific property of water……..it is an types of mixtures of excellent solvent. • Many many substances will two or more dissolve in water. molecules • Chemical processes occur best – Saline solution is one in solutions, so having 60-70% example, a mixture of or our bodies composed of the salt and water world’s best solvent is very useful….because we need our – Dextrose IV solutions cells to perform chemical are another example, processes. mixtures of sugar and water Making solutions Making solutions • The component in larger proportion • In a solution, each solute is called the solvent. Most of the molecule actually becomes solutions we will discuss will be completely surrounded by aqueous solutions, meaning the solvent molecules solvent is H2O. • In H2O solutions, we call this a hydration shell around the • The component in smaller quantity solute. is called the solute. Most of the • This ball and stick model solutes we will discuss are solids, represents a red and white which seem to ‘disappear’ in water organic molecule (like when dissolved. dextrose) completely • Gases and other liquids may also be surrounded by blue water dissolved in water. molecules 5
  6. 6. Biol 121, K. O'Neil, Instructor 1/10/2010 Like Dissolves Like Like Dissolves Like • Solubility of substances is • Although water is often termed the completely predictable “universal” solvent, it really is not. • Water – polar covalent molecule, Only certain substances will dissolve dissolves other polar covalently in water. bonded substances (like glucose) • Examples of water-soluble substances • Ionic compounds (like NaCl) are also are sugar and salt. “like” water – they are polar. • An example of a non-water soluble • Nonpolar covalently bonded substance is oil (think vinegar and oil). molecules are NOT like water, and • Because water is so vital to life, we will not dissolve in water characterize all molecules with respect to their ability to dissolve in water. Electrolytes Electrolytes • So polar covalent AND ionic • When sugar is added to water it compounds both dissolve in dissolves. Each individual sugar molecule, carrying partial water. Do they behave charges, becomes completely similarly in solution? Let’s surrounded by water. The take a close-up look at how partial charges on sugar interact hydration shells form. with the partial charges in water. • Sugar is hydrophilic (a “water lover”). Electrolytes Electrolytes • When salt (NaCl) is added to • Ionic compounds are water it also dissolves. But ionic bonds are different than covalent also hydrophilic (“water bonds -- the electrons are lovers”), but in a donated, not shared. different way than sugar. • When ionic compounds dissolve, the ions separate from each – Polar covalent molecules other, and each individual ion, dissolve in water, but do carrying full charges, becomes not dissociate completely surrounded by water. • The full charges on the ions – Ionic compounds dissolve interact with the partial charges AND dissociate in water in water. 6
  7. 7. Biol 121, K. O'Neil, Instructor 1/10/2010 Electrolytes Electrolytes and Nonelectrolytes • In water, substances are either • Ionic compounds carry full charges that are neutralized – soluble - polar covalent or ionic when the compound is a solid. – not soluble - nonpolar covalent • When the ions dissociate into • Water-soluble substances can water, the charges are no longer be further subdivided into neutralized. That’s why water – Polar covalent nonelectrolytes - containing ions carry electricity. those that just dissolve • So – no swimming in – Polar ionic electrolytes - those thunderstorms! that dissolve and dissociate pH and water • Pure water dissociates slightly into H+ and OH-. – O is so electronegative that it can pull e- away from H+ Biology 121 • Each ion is present in equal and small concentrations – H2O ↔ *H+] + [OH-] Lecture 1.5 • In pure water, each ion is present in equal concentrations pH – [H+] = [OH-] = 1 x 10-7 M (neutral) • Kw = dissociation constant of water Buffers – Kw = [H+] · [OH-] = 1 x 10-14 • Water is an electrolyte, right? Three Kinds of Electrolytes 1. Acids are Electrolytes • It is informative to classify electrolytes into subsets that • HCl is an acid. When dissociate into dissolved in water, HCl 1.) a H+ ion and something else. dissociates into a H+ cation 2.) a OH- ion and something else and a Cl- anion. 3.) neither H+ ions nor OH-. • Acid ↔ H+ + anion • These 3 categories are referred • HCl ↔ H+ + Cl- to, respectively, as • Increases [H+] 1.) acids, • Kw is constant, so when [H+] 2.) bases, and goes up, [OH-] goes down 3.) salts. 7
  8. 8. Biol 121, K. O'Neil, Instructor 1/10/2010 1. Acids are Electrolytes 2. Bases are Another Type of Electrolyte • Vinegar (acetic acid) is another example of an acid. • NaOH is a base. A base is an • Unlike HCl, acetic acid does not dissociate completely. 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 Summary - Electrolytes • NaCl is a salt. A salt is • Substances are either an electrolyte that – soluble - polar covalent or ionic dissociates into a cation – insoluble - nonpolar covalent in water. and an anion. • Water-soluble substances are • Salt ↔ cation + anion either – nonelectrolytes - polar covalent • NaCl ↔ Na+ + Cl- – electrolytes - ionic • Does not affect [H+] or • Electrolytes are either [OH-] – acids (H+), • Neither an acid nor a – bases (OH-) or – salts (neither). base Relationship between solubility and Significance of Hydrogen Ions polarity All • Why did we devise an entire substances classification scheme around H+ and OH- ions? Water- Water Solubility soluble insoluble Polarity/Bond Polar Ionic Non-polar Type Covalent covalent Electrolytes Acid Base Salt 8
  9. 9. Biol 121, K. O'Neil, Instructor 1/10/2010 Significance of H+ Ions Low, but Predictable, Dissociation Rate • Because when you put a • In a beaker of water, a tiny number of H2O molecules are hydrogen H+ ion and a dissociated. hydroxide OH- ion • In pure water, we will find 1 x 10-7 together…you get moles of H+ per liter of water – [H+] = 1 x 10-7 M WATER. • 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 Kw Never, Ever Changes • [H+] and [OH-] is constant in pure • Kw is called a CONSTANT water, and is represented by Kw (dissociation constant of water) because it remains • Kw = [H+] x [OH-] unchanged no matter what. • In pure water, we know those values. • Kw for any solution has a • Kw= (1 x 10-7) · (1 x 10-7) = 1 x 10-14 value of 1 x 10-14 Adding Acids to Water Adding Acids to Water • So, what happens if we have a beaker of water and • HCl dissociates into H+ and Cl-. So now we have MORE we add an acid, HCl, to that beaker? H+ 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. 9
  10. 10. Biol 121, K. O'Neil, Instructor 1/10/2010 Kw Never, Ever Changes Adding Acids to Water • So if [H+] goes up then [OH-] must go down. • When you add an acid to water, the [H+] goes up and • Specifically, Kw = 10-14, so if [H+] is 10-3, then [OH-] must the [OH-] goes down. be 10-11 M. Adding Bases to Water Adding Bases to Water • Let’s add a base, like NaOH, • Again, Kw is called a to that beaker. CONSTANT because it never changes. • NaOH dissociates into Na+ and OH-. So now we have • So if [OH-] goes up then [H+] must go down. more OH- than we started with. Now the • Kw = 10-14, so if [OH-] is 10-5, then [H+] must be 10-9. concentration of OH- is higher than it was. Let’s add enough NaOH to increase [OH-] to 10-5 M. pH - a convenient way to communicate Adding Salts to Water H+ concentration • Let’s try this with a salt, • pH = -log [H+]. (log base NaCl. 10 or log10) • NaCl dissociates into Na+ • pH = - exponent of the and Cl- ions. Neither H+ [H+] concentration. nor OH- are changed • Measured on a scale from the original water. from 0 to 14 • The [H+] is still 10-7, and – Neutral = 7, acidic <7, so is the [OH-]. basic >7 10
  11. 11. Biol 121, K. O'Neil, Instructor 1/10/2010 pH - a convenient way to communicate pH - a convenient way to communicate H+ concentration H+ concentration • In pure water, [H+] is 10-7. • If we add HCl, and the [H+] So pH = 7. On the pH goes up to 10-3 like in our previous example, the pH scale, that is exactly in the changes to 3. middle, and water is • So when [H+] goes up, pH considered neutral. goes down. pH - a convenient way to communicate pH - a convenient way to communicate H+ concentration H+ concentration • If we add NaOH, and the • pH 9 is on the high end of the [OH-] goes up to 10-5 and pH scale, and solutions that the [H+] goes down to 10-9, have high pH are considered keeping the Kw constant. basic (also called alkaline). The new pH is 9. • Things like ammonia and lye, which are caustic are basic. • So when [H+] goes down, pH • Blood is also very slightly basic. goes up. • Acids are on the low end of the pH scale. • Things like citrus juices, vinegar, stomach contents are acidic. pH - a convenient way to communicate Summary of pH H+ concentration • When you add NaCl, or • In summary, water dissociates into H+ and OH- at a small, but predictable, any other salt, to pure rate. water, the [H+] and [OH-] • Kw is a constant 10-14 for water. remain the same, and • If you add H+ to water, the OH- will go down. If you add OH- , the H+ will go the pH is still neutral. 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+] 11
  12. 12. Biol 121, K. O'Neil, Instructor 1/10/2010 Problems Substance [H+] pH [OH-] Pure water 10-7 mol/L neutral Biology 121 lemon juice 2 acidic Lecture 1.5 ammonia 10-11 mol/L basic bleach 9 pH coffee 5 Buffers vinegar 10-3 mol/L Strong acids v. weak acids Strong bases v. weak bases • Strong acids dissociate completely (100%) in • Strong bases dissociate completely (100%) in water – sodium hydroxide NaOH → Na+ + OH- water – potassium hydroxide KOH → K+ + OH- – hydrochloric acid HCl → H+ + Cl- • some bases increase [OH-] by accepting H+ from water – hydrobromic acid HBr → H+ + Br- • H+ acceptors are considered bases, even if they don’t • Weak acids only partially dissociate in water directly donate OH- – ammonia NH3 + H2O ↔ NH4+ + OH- – carbonic acid H2CO3 ↔ H+ + HCO3- – H+ acceptors remove H+ from solution like a ‘sponge’ (bicarbonate) • weak bases only partially dissociate (or only partially – phosphoric acid H3PO4 ↔ H+ + H2PO4- take H+ from water) Salts do not affect pH Buffers – important in biology • A salt dissociates into ions other than H+ or • pH is very important – too high or too low will OH-, no change in [H+] or [OH-] prevent cells from functioning – NaCl ↔ Na+ + Cl- – Most cells are pH 7.2-7.4 – KCl ↔ K+ + Cl- – Some compartments of cells need higher or lower • Can make a salt by combining an acid and a pH to complete their functions base – Buffers allow cells to maintain this narrow pH range – HCl + NaOH ↔ NaCl + H2O – HCl + KOH ↔ KCl + H2O 12
  13. 13. Biol 121, K. O'Neil, Instructor 1/10/2010 Buffers Buffers • Buffers interact with H+ and OH- to prevent acids • add acid (H+) to buffer: buffer accepts H+ to or bases from changing the pH of a solution prevent a drop in pH • A buffer is a weak acid or weak base that does – Example: carbonic acid H2CO3 ↔ H+ + HCO3- not dissociate completely • Add base OH- + H2CO3 ↔ H2O + HCO3- – creates a mix of acid (H+ donor) and base (H+ – (no change in [OH-] or [H+]) acceptor) • Add acid H+ + HCO3- ↔ H2CO3 – example: carbonic acid: H2CO3 ↔ H+ + HCO3- – (no change in [OH-] or [H+]) – H2CO3 = acid, H+ donor HCO3- = base, H+ acceptor – Example: ammonia NH3 + H2O ↔ NH4+ + OH- • Add base OH- + NH4+ ↔ NH3 + H2O • add base (OH-) to buffer: buffer donates an H+ to • Add acid H+ + NH3 ↔ NH4+ balance OH-, makes water Buffer has at least one Ka Buffer systems in our bodies • Ka of the buffer 1. Some proteins can donate/accept H+ to indicates the pH that it buffer cytoplasm of cells tries to maintain 2. Phosphate buffers also work in cells • Ka = pH where [H+ – H2PO4- ↔ H+ + HPO42- donor] = [H+ acceptor] – 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- Buffer systems Study Guide • • Buffers are limited – Give the equation for the dissociation of pure water, the concentration of H+ ([H+]) in pure water, and the dissociation constant of water (Kw). they can only buffer a • 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 small amount of H+ or • 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 OH- • 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 • Once buffer capacity is • Given the [H+] or pH or a solution, determine the [H+], [OH-], pH, and whether it is neutral, acidic, or basic exceeded, pH changes • Define and give an example of a weak acid and a weak base quickly • 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 13
  14. 14. Biol 121, K. O'Neil, Instructor 1/10/2010 End of Section 1 • Exam #1 will cover from the beginning of the semester to this point… 14