10/16/17 ACIDS AND BASES, DEFINED WATER IS AMPHOTERIC OUTLINE. 9.1 Properties of Acids and Bases. 9.2 ph. 9.3 Buffers

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1 ACIDS AND BASES, DEFINED A hydrogen atom contains a proton and an electron, thus a hydrogen ion (H + ) is a proton: Acids: Proton (H + ) transfer between molecules is the basis of acid/base chemistry Ø Are proton donors Ø Produce protons (H + ) when dissolved in water Bases: Ø Are proton acceptors Ø Produce hydroxide ions (OH - ) when dissolved in water WATER IS AMPHOTERIC Molecules that can act as an acid or a base are called amphoteric Solutions containing amphoteric molecules have spontaneous exchange of protons between their molecules, some acting as acids and some as bases Pure water is an important example of an amphoteric molecule: H 2 O + H 2 O H 3 O + + OH - Acts as base Acts as acid Ø This happens spontaneously, but very rarely in pure water! OUTLINE 9.1 Properties of Acids and Bases 9.2 ph 9.3 Buffers 1

2 ACIDS AND HYDRONIUM IONS In aqueous solution, acids donate a proton to water to produce a hydronium ion, H 3 O + : H + + H 2 O Example: Hydrochloric acid ( gastric juice ) à hydrochloric acid water hydronium ion chloride ion CHAPTER 7: Organic Functional Groups IONIZATION OF CARBOXYLIC ACIDS The carboxylic acid group is called an acid because it can lose a proton in aqueous solution Ø This forms a carboxylate ion Ø The name of the ion ends in ate H 3C H Acetic acid H 3C Acetate + H + A carboxylate ion is an example of a polyatomic ion COMMON BIOLOGICAL ACIDS The hydrogen atoms highlighted in pink are released as protons (H + ) in solution Note the relevant functional groups carbonic acid acetic acid phosphoric acid 2

3 BASES AND HYDROXIDE IONS In aqueous solution, bases produce hydroxide ion Base + H 2 O Hydroxides can form in two ways in water: 1. The base may steal a proton from H 2 O: 2. The base may dissociate, releasing a hydroxide ion: AMINES AS COMMON BIOLOGICAL BASES The nitrogen atoms highlighted in blue can all accept a hydrogen atom Note the relevant functional group H H ammonia ammonium ion A base must contain at least one nonbonding pair of electrons to accept an H + from water dopamine CHAPTER 7: Organic Functional Groups AMINO ACIDS: AN AMINE + CARBOXYLIC ACID Amino acids contain both the amino group and the carboxylic acid group: Ø In neutral (ph 7) solution, both groups are charged Amino acids are the building blocks of proteins. Carboxylates are (-) charged Amines are (+) charged Would you expect amino acids to be soluble in aqueous solution? 3

4 STRENGTH OF ACIDS AND BASES Acids or bases may be classified as strong or weak, depending on how they behave in water: A strong acid or base is one that fully dissociates in water: Ø Similar to an ionic compound Ø Produces electrolytes in water: H + + anion A weak acid or base is one that only partially dissociates in water: Ø The resulting solution contains a mixture of both the acid and its conjugate base STRONG ACIDS HCl is an example of a strong acid as it completely dissociates when dissolved in water: Strong acids are uncommon in biological systems STRONG BASES Sodium hydroxide (NaOH) is a good example of a strong base that completely dissociates in water: NaOH (s) + H 2 O à Na + (aq) + OH - (aq) Strong bases are also uncommon in biological systems Strong bases that utilize a metal cation are poorly soluble in water 4

5 ACID-BASE NEUTRALIZATION REACTIONS In an acid-base neutralization reaction, a proton transfers from an acid to a base, producing a salt (ionic compound) and water. The basis of all neutralization reactions is: H + + OH - à H 2 O acid base neutral The type of salt formed depends on what specific acid & base are used in the reaction: HI (aq) + KOH (aq) à H 2 O (l) + KI (aq) acid base water salt ANTACID NEUTRALIZATION REACTIONS Antacids are able to neutralize stomach acids because they are bases All of the antacids below are poorly soluble strong bases Ø They are poorly soluble so are safe due to limited release of their hydroxide ions. WEAK ACIDS Weak acids only partially dissociate in water: Ø A large portion of undissociated acid molecules remain intact in aqueous solution Ø A small fraction of molecules donate a hydrogen atom Carboxylic acids are generally weak acids: Ø A common example is acetic acid (vinegar) Strong = full dissociation Weak = partial dissociation 5

6 WEAK BASES Weak bases in solution also undergo reversible reactions with water: Ø Nearly all of the base form is present, along with a small amount of its conjugate acid as well as hydroxide ion. :NH 3 + H 2 O NH OH - ammonia water ammonium ion hydroxide ion Amines are generally weak bases: Ø The nitrogen on amines has a lone electron pair capable of binding a proton (hydrogen atom) CONJUGATE ACID-BASE PAIRS When an acid dissolves in water, hydronium ions form, and the dissociated acid remains as an anion. Ø The anion portion of the acid is its conjugate base An acid and its conjugate base are together called a conjugate acid-base pair PRACTICE PROBLEM For each of the following equations, label the conjugate acid-base pairs. Indicate whether water acts as an acid or a base in the reaction. a. NH 4+ + H 2 O H 3 O + + NH 3 b. HCO 3- + H 2 O OH - + H 2 CO 3 c. CH 3 COO - + H 2 O OH - + CH 3 COOH 6

7 REVERSIBLE REACTIONS Reactions involving a weak acid or base contain a bidirectional arrow, indicating it is reversible: acetic acid water acetate ion hydronium ion In reversible reactions, both the forward reaction and the reverse reaction are going on at the same time. Ø All chemical species the hydronium ion, the acetate ion, and acetic acid are present in solution ACID-BASE EQUILIBRIA Reversible reactions have two separate reaction rates both a forward AND a reverse rate. CH 3 COOH + H 2 O CH 3 COO - + H 3 O + acetic acid water acetate ion hydronium ion A reaction is at equilibrium when the forward and reverse reaction rates are equivalent Ø If both rates are the same, then there is no net change in reactant or product concentration over time. Chemical equilibrium is a dynamic equilibrium, and not a static equilibrium LE CHÂTELIER S PRINCIPLE An equilibrium may be disturbed by changes in conditions, such as concentration or temperature Le Châtelier s principle holds that when an equilibrium is disturbed, the reaction responds by shifting in the direction needed to restore equilibrium 7

8 EQUILIBRIUM CONSTANTS The extent to which an acid or base react with water at equilibrium is defined by its equilibrium constant (K) concentration of products K = = concentration of reactants [products] [reactants] Ø Equilibrium constants from reactions in aqueous solution do NOT account for the concentration of water CH 3 COOH + H 2 O CH 3 COO - + H 3 O + acetic acid water acetate ion hydronium ion K = [CH 3 COO - ] [H 3 O + ] [CH 3 COOH] ION-PRODUCT CONSTANT, K W In pure water, two H 2 O molecules can ionize to produce exactly one hydronium ion & one hydroxide ion: H 2 O + H 2 O H 3 O + + OH - Ø The product of their concentration is a constant value, termed the ion-product constant, or K W. Ø In pure water, concentration of each is equal at 10-7 M K w = [H 3 O + ] [OH - ] = ( M) ( M) = Because this relationship is constant, it allows us to calculate the amount of hydronium ion or hydroxide ion present, even if the equilibrium is disturbed MEASURING ACIDITY WITH ph Because the relationship between H 3 O + and OH - is constant, we can simply measure one of them to determine the acidity and basicity of a solution: Ø In practice, we only measure the concentration of hydronium ions (H 3 O + ) to determine the ph balance of a solution. The ph of a solution is defined as the negative log of the hydronium ion concentration: ph = -log [H 3 O + Use of brackets ] around a chemical formula indicate its The inverse it true as well: molar concentration [H 3 O + ] = 10 -ph (M = moles/liter) 8

9 THE ph SCALE Aqueous solutions are measured by the ph scale in units that range between 0 and 14: Ø Acidic solutions are at the lower end of this ph range. Ø Basic solutions are at the higher end of this ph range. Ø A ph of 7 is considered neutral This is a logarithmic scale! Ø Each change in one ph unit corresponds to a 10-fold difference in hydronium ion concentration. Ø If the ph of a solution changes from 4 to 2, it becomes 100 times more acidic (not twice as acidic) THE ph SCALE ph ~7.2 ph ~3.0 What is the fold change in [H 3 O + ] between your blood and a can of Coke? PRACTICE PROBLEMS 1. What is the ph of a urine sample with an [H 3 O + ] of M? 2. Calculate the [H 3 O + ] of grapefruit juice, which has a ph of

10 ph IN CELLS & LIVING TISSUES ph values can vary widely in different tissue compartments of the body: Ø The physiological ph of blood (and inside cells) has a range of ph of Ø ph values for blood are critical for good health Many medical conditions can alter blood ph & cause health complications. Ø Diabetes Ø Drug abuse Ø Kidney disease Ø Many other conditions FUNCTIONAL GROUPS AT PHYSIOLOGICAL ph Notice that certain functional groups are ionized.they carry a charge at physiological ph The form at physiological ph represents the biological useful structure. The change in chemical structure at other ph values results in loss of biological functionality. BUFFERS A buffer is a solution that resists changes in ph upon addition of small amounts of an acid or base: Ø Buffers ensure that the ph environment of an aqueous solution remains relatively constant Ø Our bodies use several different buffering systems to maintain constant ph in our blood, cells & tissues A buffer contains a weak acid and its conjugate base in roughly equal concentrations. Weak acid H 3C H Conjugate base H 3C Acetic acid Acetate 10

11 THE ACETATE BUFFER SYSTEM Sodium acetate is an ionic compound that completely dissociates into sodium & acetate ions: Ø Acetate (CH 3 COO - ) is the conjugate base for acetic acid Ø In aqueous solution, some of the acetate ions will reform the acid by reacting with water H 3C + H2O + OH - H 3C H Acetate Acetic acid Adding acid (H + ) or base to the solution will simply shift the equilibria between acetate:acetic acid: Ø Remember Le Châtelier s principle! LIMITS OF BUFFERING At some point, a buffer will lose its effectiveness: Ø ph will begin to change Ø One of the components will eventually be used up. The buffering capacity of a buffer depends upon: 1. Concentration of the weak acid and conjugate base 2. Specific properties of the acid/base pair A titration curve demonstrates the effective buffering range of acetic acid & the acetate ion PRACTICE PROBLEM When the weak base ammonia (NH 3 ) is mixed with ammonium chloride (NH 4 Cl) in water, it forms a buffered solution: a. Write the equilibrium equation representing this buffer. b. What is the purpose of NH 4 Cl in the buffer? c. How would this buffer react if H 3 O + were added? Show the equation. d. How would this buffer react if OH - were added? Show the equation. 11

12 BLOOD BUFFER Physiological buffering systems maintain a constant ph in living tissues and fluids: 1. Phosphate (PO 4-3 ), concentration in the mm range inside cells 2. Histidine, efficient buffer inside cells at neutral ph 3. Bicarbonate (HCO 3- ), important for ph buffering outside cells in blood plasma The bicarbonate buffer helps maintain blood ph: Ø The weak acid is carbonic acid (H 2 CO 3 ). Ø Its conjugate base is bicarbonate ion (HCO 3- ) H 2 CO 3 + H 2 O HCO 3- + H 3 O + weak acid conjugate base BREATHING & BLOOD ph BALANCE The source of bicarbonate in the bloodstream is carbon dioxide (CO 2 ), a metabolic waste product: Breathing (respiration) regulates ph by changing the concentration of CO 2 in the blood: Ø Ø Respiration (increases ph) Buffering system CO 2 + H 2 O H 2 CO 3 + H 2 O HCO 3- + H 3 O + Not breathing (decreases ph) Respiration removes CO 2, shifting equilibrium to the left Not breathing keeps CO 2 in the blood, shifting equilibrium to the right MEDICAL CHANGES IN BLOOD ph Acidosis results when blood ph drops below acceptable limits (too acidic) Ø Respiratory acidosis occurs when the lungs are not able to remove CO 2 efficiently Ø Metabolic acidosis occurs when there is an increase in acidity in the blood Alkalosis occurs when blood ph rises above acceptable limits (too basic = alkaline ) Ø Respiratory alkalosis results from too-rapid gas exchange (altitude sickness or hyperventilation) Ø Metabolic alkalosis may come about through vomiting or consuming too much antacid 12

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