1 4. Acid Base Chemistry 4.1. Terminology: Bronsted / Lowry Acid: "An acid is a substance which can donate a hydrogen ion (H+) or a proton, while a base is a substance that accepts a proton. B + HA <> HB + A base 1 + acid 1 <> acid 2 + base 2 Where HB = the conjugate acid A = the conjugate base Examples: HCl + water; carbonate + water; H 2 S in Water Note that water can act as either an acid or a base. Such substances are called ampholytes, and are amphoteric Lewis Acid: This is a broader definition that is stated as a substance that accepts an electron, while a Lewis base donates a proton. All Bronsted/Lowry acids are Lewis acids, not all Lewis acids anr Bronsted acids Acids in Aqueous Solution: Acids are usually refered to as donating protons, or H +, while bases donate OH, or hydroxyls The proton is strongly bound to water forming the basic unit of H 3 O+, the + Hydronium ion. This species in turn binds to other waters forming H 9 O 4 A similar structure is formed with OH, H 7 O 4 the useage of H 3 O + is usually reserved to reactions involving water. The reaction for water: H 2 O <> H + + OH K = (H+ )(OH ) (H 2 O) K a = (H + )(OH ) = K w (Dilute solution) We derived the value for K eq from G f ; G r = kcal/mol, K eq 25 C The Dissociation (Equilibrium) Constant, K a : The K a refers to the K eq for the reaction where an acid donates a proton to water; The K b is the reaction in which a base accepts a proton from a water molecule. The value of K a or K b is the strength of the acid or base. Large numbers are strong acids, small numbers are weak acids A strong acid has a stronger tendency to donate a proton A strong base has a greater tendency to accept a proton. K a is usually expressed in log. pk a is the negative log of the equilibrium constant A strong acid has a pk a that is >0.8, a weak acid is <0.8. The K a of HI is A strong base has a pk b of greater than 1.4 A very strong acid dissociates completely at almost all concentrations A very weak acid does not dissociate completely, leaving HA in solution Borderline acids have HA present at high total A concentrations. Rule of Thumb: Strong acids produce weak conjugate bases; weak acids produce strong conjugate bases: e.g. HCl v. HCO 3
2 Relationship between K b and K a : CN + H 2 O <> HCN + OH K b = (HCN)(OH) (CN)(H 2 O) HCN + H 2 O <> H 3 O + + CN K a = (H 3 O+ )(OH ) (HCN)(H 2 O) By adding these two expressions, we get 2 H 2 O <> H 3 O+ + OH; K = K a K b = (H 3 O+)(CN)(HCN)(OH) (HCN)(H 2 O)(CN)(H 2 O) = (H 3 O+)(OH) (H 2 O) **This is a legal and useful approach. Remember, the K is an algebraic, numerical constant. Because the activity of water is 1: K a K b = (H 3 O + ) + (OH ) = K w = Where K w is the equilibrium constant water. This relationship allows the calculation of either K a or K b if the other is known by: ph + poh = pk w = 14 therefore, in pure water when the (H + ) = (OH ), the ph = 7, the poh = 7 Acid solution have a ph < 7, basic solutions have a ph > 7, or a poh < Neutrality, and the affect of temperature: Neutral solutions are ph 7 only at 25 C, where the K w is At other temperatures, the ph changes because the equilibrium changes. As temperature increases, since there is a positive H, dissociation increases. The H of the dissociation of water is kcal/mole ln K 50 C = ln K w, 25 C + H R So at 50 C, neutrality is at ph = = 5.73 x Muliprotic acids: A monoprotic acid has one proton that can be donated under normal circumstances (i.e. environmental conditions. A diprotic acid has 2, a triprotic 3, etc.) For groundwater systems, triprotic acids are the most complicated to deal with, and there are several of these, e.g. phosphoric acid, boric acid, and citric acid. Each proton has a unique Ka, or acid dissociation constant, with names of K a,1, K a,2, etc. A general rule for Bronsted acids is that the first proton to be donated from a multiprotic acid is much more acidic than the second, the second more acidic than the third, etc. In organic acids, the second and third may be similar. The conjugate base of the first deprotonation of a diprotic acid is amphoteric, and is both a proton donor and a proton acceptor. 2
3 4.2. Equilibrium Calculations This is a systematic approach to solving acidbase equilibrium that uses 4 equations. The goal is to use enough equations, and the right equations, to solve the specific problem being examined. An important consideration, and one of the most difficult to master, is the use of specific a assumptions, used to simplify the equations. This a trial and error approach Mass Balance: In acidbase reactions the reacting species are conserved, assuming a closed system. When HA is added to water, the acid ionizes (partly or completely): HA + H 2 O > A + H 2 O A mass balance equation would be: C T,A = [HA] + [A] C T,A is the analytical concentration of all species containing A, with units being moles per liter. Note that concentration is used. When C moles of the salt MA are added to a liter of solution: MA <> M + + A A + H 2 O <> HA + OH C T,M = [M+] + [MA], or if MA dissociates completely, C T,M = [M+] Equilibrium Relationships This utilizes the known equilibrium constant of the acids involved to help solve for concentration of an unknown species. The first equation is K w. K w = (H+)(OH) = at 25 o C Note that here activity must be used, and the difference between activity and concentration used in the mass balance equation must be reckoned. The dissociation of HA is described by : + ( H )( A ) ( HA)( OH ) Ka = Kb = ( HA) ( A ) Since K w = K a K b, both of these equations can not be used since they are not independent Charge Balance (electroneutrality) All solutions are electrically neutral. Unlike the analytical charge balance equation, the CB equation used here is an absolute equality The total molar concentration of cations is always equal to the total anions. [M + ] + 2[M 2+ ] + [H + ] = [A ] + + 2[A 2 ] + [OH ] NOTE: concentration is used, not activity. Also, note that [HA] is not used, as it is a neutral species. 3
4 Proton Condition: The proton condition is a mass balance on protons, and is essential part of solving for protons involved in equilibria. The proton mass balance is established with reference to a "zero level" or reference level for protons, called the PRL. The PRL is established as the species with which the solution was prepared, and species having protons in excess of the PRL are equated with species having less than the PRL. This means that protons are conserved, and if protons are transfered from the initial reactants, that the result is some species gain protons (>PRL), while others lose protons (<PRL) Strong AcidBase Calculations. Examples of strong acids in ground water are nitric and sulfuric acid from acid rain or industrial discharge, hydrochloric acid used to open up a well (dissolve calcite) Strong acids have a large K a, indicating that there are more products than reactants, i.e. that the reaction proceeds very far toward to products. The dissociation of strong acids in water is exothermic and has a large positive entropy. The release of heat is from the hydration of proton by water to make hydronium ion. + ( H )( Cl ) HCl Æ Cl + H+ Ka = K a = ~10 3, pk~ 3 ( HCl) So the product of proton and chloride concentration is 1000 times the concentration of free HCl. This leads to the basic assumption used to simplify acidbase calculations: STRONG ACIDS DISSOCIATE COMPLETELY in dilute solutions. This results simplification of the mass balance calculation as: [Cl]>>[HCl] C T,Cl = [HCL]+[Cl] Strong and intermediate acids important in natural systems: Acid Name Formula pk Perchloric HClO 4 7 Hydrogen Chloride HCl 3 Sulfuric H 2 SO 4 3 (pk 1 ) Nitric HNO 3 1 Hydronium H 3 O + 0 Bisulfate HSO (pk 2 ) Phosphoric H 3 PO (pk 1 ) Aquo ferric ion** 3+ Fe(H2O) (pk 1 ) Arsenic Acid H 3 AsO (pk 1 ) **The ferric iron ion in solution is coordinated to 6 waters of hydration, and you can think of these waters as having acidic protons. Another way of thinking of it is the exchange of a water for a hydroxyl ion that decreases the hydroxyl activity, thus decreasing ph i.e. an acid. So adding FeCl 3 to water decreases ph. Aluminum acts in the same way. 4
5 4.4. Weak AcidBase Calculations Weak acids and bases do not dissociate completely, so while the approach to solving the equations is similar to strongacid systems, the complication of the K a is added. The use of simplifying assumptions is even more important for this system. Acid Name Formula pk Hydrofluoric HF 3.45 Acetic CH 3 COOH 4.7 Arsenic Acid H 2 AsO Hydrogen Sulfide H 2 S 7.1 Dihydrogen Phosphate H 2 PO (pk 2 ) Hypochlorous Acid HOCl 7.6 Hydrogen Cyanide HCN 9.2 Boric Acid H 3 BO Ammonium Ion + NH Bicarbonate Ion HCO Arsenic Acid H 2 AsO (pk 3 ) Bisulfide HS 14 (pk 2 ) Water H 2 O 14 The small pk values mean that the concentration of the reactants is greater than the products, and the reaction does not proceed very far to the right. Example: Hydrogen sulfide is produced by sulfate reducing bacteria in some anaerobic groundwater, and is common in water wells east of Austin. Hydrogen sulfide is a very weak acid H 2 S Æ H + + HS + ( H )( HS ) K a,1 = = , pk = 7.1 ( H 2S) The fundamental relationship with weak acids is that the extent of of the reaction, or the amount of acid that dissociates, is directly related to the ph. When ph < pk, little acid dissociates. As ph increases, more acid must dissociate to maintain equilibrium. The relationship between ph and weak acids is also expressed as: salt ph = pka +log [ ] [ acid] Rule of Thumb: With weak acids, When the pk = ph, the activity of the acid is equal to the conjugate base. This is a critical relationship that you must remember!! i.e. For H 2 S, at ph 7.1, (H 2 S) = (HS ) 5
6 4.5. Strong AcidStrong Base Mixtures For this system we need to examine the changes that occur when acids and bases are mixed, rather then simply solving a static system. The general approach is to write equations that describe the solution after the acid and base are mixed, make simplifying assumptions, then solve as usual. Forget using the equilibrium equation there is no appreciable undissociated acid or base left. The proton condition equation is possible, but difficult. In this case, you assume that both the acid and base dissociate completely (with much heat evolved ), and you now have a salt solution at some new ph. The best equations to use are the mass balance and charge balance equations. If you know the anion concentration, and the metal cation concentration (e.g. add HCl to a NaOH solution), then solve for proton When the system is going to change across ph 7 (e.g. acid to base) try to use the mixture to first get to neutral, and then to the final ph. It is conceptually easier than going directly from acid to base conditions. 4. Strong acidweak base mixtures. These mixtures are more complicated because we lose the assumption of complete dissociation. Most natural waters, however, behave like a mixture of weak acids and bases. An understanding of these systems is essential to understanding ph buffers. number of equivalents of NaOH added Equivalent Fraction f = C TAc, OVERHEAD: Fig When f=0.5, ph = pk s. Any addition of base corresponding to f<0.5 will produce a solution ph between that of the original HAc solution and the pk a. 2. At f=0.5, the titration curve has an invlection point. At this ph (=pk a ) the change of ph for a given addition of strong base is small, and the solution is well buffered. 3. When f=1.0, the ph is that of a 0.001M NaAc solution, or the equiv of Na added is equal to the [Ac ] present. This ph is the equivalence point. At this ph value the curve has an inflection point in the opposite direction to that at f=0.5, and the resistance of the solution to ph change is small, or the system is poorly buffered. 4. The ph increases very slowly as it appraoches the ph of the solutin being added. 5. ph Buffers and Buffer Intensity: A buffer solution is a solution that in some way has the ability to maintain a stable composition when various components are added or removed. 6
7 A solution buffered at a particular ph value will contain an acid that reacts with any strong base added to the solution and a base that reacts with any strong acid added to the solution. The ph of a buffer solution can be predicted by the equation (assuming infinite dilution): salt ph = pka +log [ ] [ acid] Question: Which acidconjugate base pair should be used to buffer a solution at ph 4.7 and at ph 7? What ratio of acid to conjugate base should be used in each of these buffers? 1. Examine a table of K a values to choose an acid with a pk a near the desired ph. Acetic acid has a pk a of 4.7, and should be good at ph Simplifying assumptions: [H+] and [OH] << [HAC] and [NaAc] [ H ] 10 [ HAC] = = = Ka 10 [ NaAc] 3. The ph = pk therefore requires equal concentrations of salt and acid. While this gives the ratio, it does not give the concentration of salt and acid. The greater the concentration, the greater the buffer capacity. 4.. H 2 PO 4 has a pk near 7.0, and can be used to buffer a system: Buffer Intensity 72. [ Na2HPO4 ] 10 = = [ NaH PO ] Buffer capacity (β) or Buffer intensity is defined as the moles/l of strong base which when added to a solution causes a unit change in ph.: + doh [ ] dh [ ] β = = dph dph So the reciprocal of the slope of the titration curve is proportional to the buffer intensity. In general the buffer intensity is a maximum at ph = pk, and that β is at a minimum when f=1, at the equivalence point of the titration. Figure 415 OVERHEAD 7
8 4.6. Carbonic Acid/Carbonate system 1. Gas Exchange: Gaseous carbon dioxide, CO 2(g) dissolves into water producing CO 2(aq). The equilibrium of this processes is the Henry's Law Constant for carbon dioxide: K = K H = (CO 2(aq) ) (CO 2(g) ) = This species has not yet reacted with anything 2. Hydrolysis of Carbon Dioxide: CO 2(aq) + H 2 O <> H 2 CO 3 K = (H 2 CO 3 ) (CO 2(aq) ) = Based on the equilibrium constant, the extent of this reaction is small, so H 2 CO 3 is only about 0.16% of the CO 2(aq) at 25 This reaction is also kinetically slow, slower than acidbase reactions. It is nearly impossible to analytically separate the two species, and it is difficult to determine the acidity constant of H 2 CO 3 so the two are simply added to form the species H 2 CO 3 * 3. First Acidity Constant: K a,1 = (H+ )(HCO 3 ) (H 2 CO 3 *) ; K' 1 = (H+ )(HCO 3 ) (H 2 CO 3 ) = By substituting the summation for our fictional carbonic acid species, we obtain K a,1 = (H + )(HCO 3 ) (H 2 CO 3 ) + (CO 2(aq) ) Now multiplying the numerator and denominator by (H 2 CO 3 ), and knowing that: Since 1/Km >>1, (CO 2(aq) )/(H 2 CO 3 ) = 1/K m : (H + )(HCO 3 )/(H2 CO 3 ) (CO 2(aq) ) + (H 2 CO 3 )/(H 2 CO 3 ) = K' 1 (1/K m ) +1 = K a,1 K a,1 = K' 1 K m = = This means that while carbonic acid is a strong acid, the combination of the two species, and the slow kinetics, make it appear as a weak acid It actually is a fairly strong acid capable of dissolving silicates. 8
9 4.1.7 SYSTEMS APPROACH carbonic acid system: 1. Case 1. Closed system, no solid equations: K w Unknowns, (H + ), (OH) ph = 7 2. Case 2. Open System, No Solid Problem: Pure water in equilibrium with atmospheric carbon dioxide. No other solutes, assume infinite dilution, 25 o C. What are the concentrations of the unknown species. equations: K H, K 1, K 2, K w Unknowns: (H + ), (HCO 3 ), (OH ), (H2 CO 3 ), (CO 3 2 ) Constants: PCO 2 = Atm. Assumptions: (H + ) >> (OH ); (HCO 3 ) >> (CO 3 2 ) Mass Balance: C T,(CO3) = [CO 3 2 ] + [HCO3 ] + [H 2 CO 3 ] assumptions... Electroneutrality: [H + ] = [HCO 3 ] + 2[CO3 2 ] + [OH ] Solution: assumptions... 1) (H 2 CO 3 ) = K H PCO 2 = ( )( ) = 10 5 M 2) (H 2 CO 3 ) = (H+ )(HCO 3 ) K 1 ; K 1 (H 2 CO 3 ) = (H + )(HCO 3 ) 3) (10 5 )( ) = (H+) 2 4) (H + ) = , ph = 5.67 (HCO 3 ) = (OH ) = ) (CO 3 2 ) = K2 (HCO 3 )/(H + ) = In an open system, the [H 2 CO 3 ] remains constant, fixed by the PCO 2 The ph of distilled water open to the air is
10 Case 6: Open System, Fixed ph Problem: Water in equilibrium with atmospheric carbon dioxide. ph is fixed at 8.0, 25 o C using NaOH. What are the concentrations?. equations: K H, K 1, K 2, K w Unknowns: (HCO 3 ), (OH ), (H2 CO 3 ), (CO 3 2 ) Constants: PCO 2 = Atm. Assumptions: (H + ) < (OH ); (HCO 3 ) >> (H 2 CO 3 ) Mass Balance: C T,(CO3) = [CO 3 2 ] + [HCO3 ] + [H 2 CO 3 ] assumptions... Electroneutrality: [H + ] + [Na + ] = [HCO 3 ] + 2[CO3 2 ] + [OH ] Solution: assumptions... 1) (H 2 CO 3 ) = K H PCO 2 = ( )( ) = 10 5 M 2) (H 2 CO 3 ) = (H+ )(HCO 3 ) K 1 ; K 1 (H 2 CO 3 ) = (H + )(HCO 3 ) 3) (10 5 )( )/10 7 = (HCO3) 5) (CO 3 2 ) = K2 (HCO 3 )/(107) In an open system, even with fixed ph, the [H 2 CO 3 ] remains constant, fixed by the PCO 2 10
AP Chemistry Acids and Bases General Properties of Acids and Bases Acids Electrolyte Taste Litmus Phenolphthalein React with metals to give off H 2 gas H 2 SO 4 (aq) + Mg (s) MgSO 4 (aq) + H 2 (g) Ionize
AcidBase Chemistry Arrhenius acid: Substance that dissolves in water and provides H + ions Arrhenius base: Substance that dissolves in water and provides OH ions Examples: HCl H + and Cl Acid NaOH Na +
General Chemistry II Jasperse Acid-Base Chemistry. Extra Practice Problems 1 General Types/Groups of problems: Conceptual Questions. Acids, Bases, and p1 K b and pk b, Base Strength, and using K b or p7-10
Chapter 14 - Acids and Bases 14.1 The Nature of Acids and Bases A. Arrhenius Model 1. Acids produce hydrogen ions in aqueous solutions 2. Bases produce hydroxide ions in aqueous solutions B. Bronsted-Lowry
MULTIPLE CHOICE. Choose the one alternative that best completes the statement or answers the question. 1) Which statement concerning Arrhenius acid-base theory is not correct? A) Acid-base reactions must
1. For the equilibrium that exists in an aqueous solution of nitrous acid (HNO 2, a weak acid), the equilibrium constant expression is: a) K = [H+ ][NO 2 ] [HNO 2 ] b) K = [H+ ][N][O] 2 [HNO 2 ] c) K =
Acid-base A4 1 Acid-base theories ACIDS & BASES - IONIC EQUILIBRIA 1. LEWIS acid electron pair acceptor H, AlCl 3 base electron pair donor NH 3, H 2 O, C 2 H 5 OH, OH e.g. H 3 N: -> BF 3 > H 3 N BF 3 see
: General Chemistry Lecture 9 Acids and Bases I. Introduction A. In chemistry, and particularly biochemistry, water is the most common solvent 1. In studying acids and bases we are going to see that water
WEAK ACIDS AND BASES [MH5; Chapter 13] Recall that a strong acid or base is one which completely ionizes in water... In contrast a weak acid or base is only partially ionized in aqueous solution... The
Chapter 8 Acids and Bases Definitions Arrhenius definitions: An acid is a substance that produces H + (H 3 O + ) Ions in aqueous solution. A base is a substance that produces OH - ions in aqueous solution.
Acid-Base (Proton-Transfer) Reactions Chapter 17 An example of equilibrium: Acid base chemistry What are acids and bases? Every day descriptions Chemical description of acidic and basic solutions by Arrhenius
Chapter 16 Acid-Base Equilibria Learning goals and key skills: Understand the nature of the hydrated proton, represented as either H + (aq) or H 3 O + (aq) Define and identify Arrhenuis acids and bases.
AcidBaseEquil 1 Acid-Base Equilibrium See AqueousIons in Chemistry 1110 online notes for review of acid-base fundamentals! Acid- Base Reaction in Aqueous Salt Solutions Recall that use [ ] to mean concentration
Acids Identifying Acids and Bases Acid (anhydrides) contains H+ ions as the cation, with and other element as the anion Non-metal oxide H2SO4 HI P2O5 Bases Base (anhydrides) Contains OH- as the anion Combined
Bronsted Lowry Acid Base Chemistry Just a few reminders What makes an acid an acid? A BronstedLowry Acid is a compound that donates a proton (a hydrogen ion with a positive charge, H + ) o Think of the
Acid-Base Indicators and Titration Curves Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical
CHEM 102: Sample Test 5 CHAPTER 17 1. When H 2 SO 4 is dissolved in water, which species would be found in the water at equilibrium in measurable amounts? a. H 2 SO 4 b. H 3 SO + 4 c. HSO 4 d. SO 2 4 e.
Topic 5 5-1 Acid and Bases Acid and Bases 5-2 There are a number definitions for aicd and bases, depending on what is convenient to use in a particular situation: Arrhenius and Ostwald: Theory of electrolyte
Acids and Bases Know the definition of Arrhenius, Bronsted-Lowry, and Lewis acid and base. Autoionization of Water Since we will be dealing with aqueous acid and base solution, first we must examine the
Chapter 17 Acids and Bases How are acids different from bases? Acid Physical properties Base Physical properties Tastes sour Tastes bitter Feels slippery or slimy Chemical properties Chemical properties
ACIDS & BASES BRØNSTED ACIDS & BASES BRØNSTED ACIDS & BASES Brønsted acids are proton donors. Brønsted bases are proton acceptors. Amphoteric species can act as either an acid or a base, depending on the
CHEM 10123/10125, Exam 2 March 7, 2012 (50 minutes) Name (please print) Please box your answers, and remember that significant figures, phases (for chemical equations), and units do count! 1. (13 points)
Chem 1B Dr. White 1 Chapter 14 Acids and Bases 14.1 Nature of Acids and Bases A. Acids B. Bases Chem 1B Dr. White 2 C. Arrhenius Definition 1. acid 2. base 3. Acid-base reaction involving Arrhenius acids
ph: Measurement and Uses One of the most important properties of aqueous solutions is the concentration of hydrogen ion. The concentration of H + (or H 3 O + ) affects the solubility of inorganic and organic
Acids and Bases Chapter 16 The Arrhenius Model An acid is any substance that produces hydrogen ions, H +, in an aqueous solution. Example: when hydrogen chloride gas is dissolved in water, the following
Exercise #1 Brønsted-Lowry s and Bases 1. Identify the Bronsted-Lowry acids and bases and the conjugate acid base pairs in the following reactions. (a) HCl(aq) + H 2 O(l) H 3 O + (aq) + Cl (aq) (b) H 2
Chapter 16 AcidBase Equilibria Acids and bases are found in many common substances and are important in life processes. Group Work: Make a list of some common acids and bases. How do we know which is which?
ph. Weak acids. A. Introduction... 1 B. Weak acids: overview... 1 C. Weak acids: an example; finding K a... 2 D. Given K a, calculate ph... 3 E. A variety of weak acids... 5 F. So where do strong acids
1 Ch 8.5 Solution Concentration Units % (m/m or w/w) = mass of solute x 100 total mass of solution mass of solution = mass solute + mass solvent % (v/v) = volume of solute x 100 volume of solution filled
Lecture 6 Classes of Chemical Reactions Lecture 6 Outline 6.1 The Role of Water as a Solvent 6.2 Precipitation Reactions 6.3 Acid-Base Reactions 1 Electron distribution in molecules of H 2 and H 2 O The
CHAPTER 9 1. Which one of the following is the acid in vinegar? a. acetic acid b. citric acid c. muriatic acid d. ascorbic acid 2. Which is a basic or alkaline substance? a. gastric fluid b. black coffee
Acids and : A Brief Review Acids: taste sour and cause dyes to change color. : taste bitter and feel soapy. Arrhenius: acids increase [H ] bases increase [OH ] in solution. Arrhenius: acid base salt water.
UNIT (6) ACIDS AND BASES 6.1 Arrhenius Definition of Acids and Bases Definitions for acids and bases were proposed by the Swedish chemist Savante Arrhenius in 1884. Acids were defined as compounds that
CHEM 1105 ACIDS AND BASES 1. Early Definitions Taste: Effect on Indicators: Neutralization: acids - sour; bases - bitter acids turn blue litmus red; bases turn red litmus blue phenolphthalein is colourless
Name Chemistry Pre-AP Notes: Acids and Bases Period I. Describing Acids and Bases A. Properties of Acids taste ph 7 Acids change color of an (e.g. blue litmus paper turns in the presence of an acid) React
1. In which reaction is water acting only as a proton acceptor? 1) H 2 SO 4 (aq) H 2 O ( ) HSO 4 (aq) H 3 O (aq) NH 3 (g) H 2 O ( ) NH 4 (aq) OH (aq) CH 3 COO (aq) H 2 O ( ) CH 3 COOH(aq) OH (aq) H 2 O
ph of strong acid and base What does strong mean in terms of acids and bases? Solubility. Basically, if you say an acid or base is strong, then it dissociates 100% in water. These types of situations are
A.P. Chemistry Practice Test: Ch. 14, Acids and Bases Name MULTIPLE CHOICE. Choose the one alternative that best completes the statement or answers the question. 1) The conjugate base of HSO4 - is A) H2SO4
Chapter 16: Acid-Base and Solubility Equilibria: Reactions in Soil and Water Problems: 16.2-16.86 16.1 ACIDS AND BASES: THE BRØNSTED-LOWRY MODEL PROPERTIES OF ACIDS & BASES Acids produce hydrogen ions,
Chapter 15 Acids and Bases reading guide. Be active while reading the text. Take notes, think about what you ve read, and ask yourself questions while reading. Use this document as a guide for making your
EQUILIBRIUM PART 2 V. POLYPROTIC ACID IONIZATION A. Polyprotic acids are acids with two or more acidic hydrogens. monoprotic: HC 2 H 3 O 2, HCN, HNO 2, HNO 3 diprotic: H 2 SO 4, H 2 SO 3, H 2 S triprotic:
Acid/Base Reactions some covalent compounds have weakly bound H atoms and can lose them to water (acids) some compounds produce OH in water solutions when they dissolve (bases) acid/base reaction are very
SCH3U- R.H.KING ACADEMY SOLUTION & ACID/BASE WORKSHEET Name: The importance of water - MAKING CONNECTION READING 1. Read P. 368-375, P. 382-387 & P. 429-436; P. 375 # 1-11 & P. 389 # 1,7,9,12,15; P. 436
Ch 14 Page 1 Chapter 14: Acids and Bases Properties of Acids Sour taste React with some metals Turns blue litmus paper red React with bases Some Common Acids HCl, hydrochloric acid H 2 SO 4, sulfuric acid
Chapter 16: Acid-Base Equilibria In the 1 st half of this chapter we will focus on the equilibria that exist in aqueous solutions containing: weak acids polyprotic acids weak bases salts use equilibrium
Chemistry 201 Lecture 15 Practical aspects of buffers NC State University The everyday ph scale To review what ph means in practice, we consider the ph of everyday substances that we know from experience.
Acid-base Equilibria and Calculations A Chem1 Reference Text Stephen K. Lower Simon Fraser University Contents 1 Proton donor-acceptor equilibria 4 1.1 The ion product of water... 4 1.2 Acid and base strengths...
Titrations Titration - a titration is defined as the determination of the amount of an unknown reagent (analyte) through the use of a known amount of another reagent (titrant) in an essentially irreversible
Ionization of Water DEMONSTRATION OF CONDUCTIVITY OF TAP WATER AND DISTILLED WATER Pure distilled water still has a small conductivity. Why? There are a few ions present. Almost all the pure water is H
Talk in. Arrhenius Acid Base Definition and ph When we think of acids, we typically think of the Arrhenius definition. Svante Arrhenius (18591927) Arrhenius Acid = Any compound that increases the hydronium
Name: Class: Date: Chapter 13 & 14 Practice Exam Multiple Choice Identify the choice that best completes the statement or answers the question. 1. Acids generally release H 2 gas when they react with a.
Chapter 16 Acids and Bases Concept Check 16.1 Chemists in the seventeenth century discovered that the substance that gives red ants their irritating bite is an acid with the formula HCHO 2. They called
Chapter 9 Lecture Notes: Acids, Bases and Equilibrium Educational Goals 1. Given a chemical equation, write the law of mass action. 2. Given the equilibrium constant (K eq ) for a reaction, predict whether
Acids, Bases & Redox 1 Practice Problems for Assignment 8 1. A substance which produces OH ions in solution is a definition for which of the following? (a) an Arrhenius acid (b) an Arrhenius base (c) a
Acid-base Equilibria and Calculations A Chem1 Reference Text Stephen K. Lower Simon Fraser University Contents 1 Proton donor-acceptor equilibria 4 1.1 The ion product of water... 4 1.2 Acid and base strengths...
CHEM 101/105 Aqueous Solutions (continued) Lect-07 aqueous acid/base reactions a. a little bit more about water Water is a polar substance. This means water is able to "solvate" ions rather well. Another
Titration Curve of a Weak Acid Amina Khalifa El-Ashmawy, Ph.D. Collin College Department of Chemistry Introduction: Titration is an analytical process whereby two reactant solutions are carefully reacted
Lab Activity: Strong Acids & Strong Bases, Weak Acids & Weak Bases, and the Hydrolysis of Salts Strong acids and strong bases completely ionize in water to make hydronium or hydroxide ions. Because the
Dr Mike Lyons School of Chemistry Trinity College Dublin. email@example.com ACID-BASE REACTIONS/ THE PH CONCEPT. Chemistry Preliminary Course 2011 1 Lecture topics. 2 lectures dealing with some core chemistry
Worksheet 42 Bronsted Acids and Equilibria Worksheet 42 Bronsted Acids and Equilibria Name Date Due Hand In With Corrections by 1. Write the formula for a proton 2. Write the formula for a hydrated proton
1 Aqueous Equilibria: Buffers & Titrations AMHS AP Chemistry Multiple Choice questions Name Common-Ion Effect 1) The ph of a solution that contains 0.818 M acetic acid (K a = 1.77x10-5 ) and 0.172 M sodium
Lecture Presentation Chapter 9 Acids, Bases and Buffers in the Body Julie Klare Fortis College Smyrna, GA Outline 9.1 Acids and Bases Definitions 9.2 Strong Acids and Bases 9.3 Chemical Equilibrium 9.4
Name: Period: Date: I. Definitions: 1. Arrhenius (1887)- produce in aqueous solutions produce in aqueous solutions 2. Bronsted-Lowry (1923) 3. Lewis (1923), so only allows for aqueous solutions only protic
15 Ch a pt e r Aqueous Equilibria: Acids and Bases Chemistry th Edition McMurry/Fay Dr. Paul Charlesworth Michigan Technological University AcidBase Concepts 01 Arrhenius Acid: A substance which dissociates
AP*Chemistry The Chemistry of Acids and Bases "ACID"--Latin word acidus, meaning sour. (lemon) "ALKALI"--Arabic word for the ashes that come from burning certain plants; water solutions feel slippery and
Your consent to our cookies if you continue to use this website.