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CONCEPT: ACID IDENTIFICATION The most common feature of an acid is that many possess an H + ion called the. When it comes to acids there are 2 MAJOR TYPES that exist: are acids where the H + ion is attached to an electronegative element. These types of acids lack the element and usually possess no. The most common type of these particular acids are the haloacids:,, &. are acids that contain the, &. They are created by the hydration of nonmetal oxides. PRACTICE: Which of the following compound(s) cannot be classified as an acid? a) H2S b) HCN c) H2 d) C6H6 e) All are acids. Page 2
CONCEPT: BINARY ACID STRENGTH STRONG ACIDS are considered Electrolytes so they ionize completely in water. HCl (aq) H 2 O H + (aq) + Cl (aq) WEAK ACIDS are considered Electrolytes so they don t completely ionize in water. HF + H2O F (aq) + H3O + (aq) The strength of a BINARY ACID is based on the or of the nonmetal. For elements in the same period then look at their. The, the acidic. For elements in the same group then look at their. The, the acidic. BINARY ACID STRENGTH PRACTICE 1: Which is the weakest acid from the following? a) H2S b) H2Se c) H2Te d) All would have the same acid strength. PRACTICE 2: Which of the following acids would be classified as the strongest? a) CH4 b) NH3 c) H2O d) HF e) PH3 Page 3
CONCEPT: OXYACID STRENGTH The strength of OXYACIDS is based on the number of or the of the nonmetal. RULE: If my oxyacid has 2 or More than then my oxyacid is a ACID. HNO 3 Oxygens Hydrogens C 6 H 5 OH Oxygens Hydrogens HBrO 4 Oxygens Hydrogens When comparing the strengths of different oxyacids remember: If they have different number of oxygens then the oxygen the acidic If they have the same number of oxygens then the electronegative the nonmetal the acidic. 3 Exceptions Electronegativity H 2 C 2 O 4 Oxygens Hydrogens HIO 3 Oxygens Hydrogens HSO 4 Oxygens Hydrogens PRACTICE: Rank the following oxyacids in terms of increasing acidity. a) HClO3 b) HBrO4 c) HBrO3 d) HClO4 Page 4
CONCEPT: BASE STRENGTHS STRONG BASES are considered Electrolytes so they ionize completely in water. NaOH (aq) H 2 O Na + (aq) + OH (aq) WEAK BASES are considered Electrolytes so they don t completely ionize in water. NH3 + H2O NH4 + (aq) + OH (aq) Bases possess THREE major features: or or. Group : Any Group metal when combined with OH, H, O 2 or NH2 makes a STRONG BASE. Group : Any Group metal, from to, when combined with OH, H, O 2 or NH2 makes a STRONG BASE. : are considered WEAK BASES. Ex: are considered WEAK ACIDS. Ex: Page 5
PRACTICE: ACID & BASE IDENTIFICATION EXAMPLE: Classify each of the following as a strong acid, weak acid, strong base or weak base. a) HCHO2 c) H2NNH2 b) (CH3CH2)3NH + d) HBrO3 PRACTICE 1: Classify each of the following as a strong acid, weak acid, strong base or weak base. a) KOCH3 b) CH3OH PRACTICE 2: Classify each of the following as a strong acid, weak acid, strong base or weak base. a) HOCN b) H5IO6 PRACTICE 3: Classify each of the following as a strong acid, weak acid, strong base or weak base. a) NaN3 b) SrH2 Page 6
CONCEPT: ARRHENIUS ACIDS & BASES The most general definition for acids and bases was developed by Svante Arrhenius near the end of the 19 th century. According to him, the cation and the anion are fundamental to the concept of acids and bases. His definition however failed to describe acidic and basic behavior in nonaqueous media. The Arrhenius definition states an acid is a compound that increases when dissolved in a solvent. The Arrhenius definition states a base is a compound that increases when dissolved in a solvent. PRACTICE 1: Which ions are formed from the dissociation of the following compound? a) Sr(OH)2 (s) Dissolves in H 2 O PRACTICE 2: Which ions are formed from the dissociation of the following compound? a) H2SO4 (l) Dissolves in H 2 O PRACTICE 3: Which ions are formed from the dissociation of the following compound? a) HBO3 2- Dissolves in H 2 O Page 7
CONCEPT: BRONSTED LOWRY ACIDS & BASES In 1923, Johannes Brønsted and Thomas Lowry developed a new set of definitions for acids and bases. According to the Bronsted-Lowry definition, acids are considered and bases are considered. Unlike Arrhenius acids and bases, they are not limited to aqueous solutions. Every Arrhenius acid is a Brønsted-Lowry acid (and likewise for the bases). Brønsted-Lowry acids and bases always occur in pairs called. EXAMPLE 1: Write the formula of the conjugate base for the following compound: HSO4 EXAMPLE 2: Write the formula of the conjugate acid for the following compound: V2O5 2- PRACTICE 1: Write the formula of the conjugate base for the following compound: H2Se PRACTICE 2: Write the formula of the conjugate for the following compound: NH2NH2 Page 8
PRACTICE: BRONSTED LOWRY ACIDS & BASES (CALCULATIONS) EXAMPLE 1: Identify the acid, base, conjugate acid and conjugate base in the following reactions: a) HF (aq) + H2O (aq) F (aq) + H3O + (aq) EXAMPLE 2: Identify the acid, base, conjugate acid and conjugate base in the following reactions: a) CN (aq) + H2O (aq) HCN (aq) + OH (aq) PRACTICE 1: Which of the following is a Bronsted-Lowry acid? a) CH4 b) HCN c) NH3 d) Br2 PRACTICE 2: Determine the chemical equation that would result when carbonate, CO3 2-, reacts with water. Page 9
CONCEPT: AMPHOTERIC SPECIES An amphoteric, or, is a species that can act as a(n) ACID or BASE. Water is prime example of an amphoteric species. Partially dissociated conjugate bases of polyprotic acids are also amphoteric. These compounds possess and a. Ex: PRACTICE: Which of the following species is/are amphoteric? a) CO3 2 b) HF c) NH4 + d) HPO3 2- e) CH3O Page 10
CONCEPT: LEWIS THE FINAL TYPE OF ACID & BASE In the 1920s, Gilbert Lewis proposed a new set of definitions for acids and bases. A Lewis acid is a(n). acts as a Lewis acid when connected to an electronegative element:,,,, or charged hydrogen or metals. If your central element has 8 valence electrons. A Lewis base is a(n). Compounds with. NH 3 H 2 O CH 3 OH CH 3 OCH 3 Compounds with a. CN OH CH 3 O N 3 Page 11
PRACTICE: LEWIS.THE FINAL TYPE OF ACID & BASE (CALCULATIONS) EXAMPLE: Identify each of the compounds in the following chemical equation. Br Br H 3 C H 2 C H 3 C O CH 2 Br Al Br H 3 C H 2 C H 3 C O CH 2 Al Br Br PRACTICE 1: Identify the Lewis acids and bases in the following reactions. a) H + + OH H2O b) Cl + BCl3 BCl4 c) SO3 + H2O H2SO4 PRACTICE 2: Identify each of the following compounds as either a Lewis acid, a Lewis base or neither. a) ZnCl2 b) CN c) NH4 + d) Co 3+ Page 12
CONCEPT: ph and poh To deal with incredibly small concentration values of [H + ] and [OH - ] we can use the ph scale. Under normal conditions, the ph scale operates within the range of to. By taking the log of [H + ] and [OH - ] we can find ph and poh. ph = log[h + ] poh = log[oh ] p = log By recognizing the relationship between [H + ] and [OH - ] with ph and poh we can create new formula relationships. ph = log[h + ] poh = log[oh ] A species with a ph greater than 7 is classified as and the [H + ] is than the [OH - ]. The the base then the the ph. A species with a ph less than 7 is classified as and the [H + ] is than the [OH - ]. The the acid then the the ph. A species with a ph equal to 7 is classified as and the [H + ] is than the [OH - ]. By using log with the equilibrium expression for water a relationship between ph and poh can be created. ph + poh =14 Page 13
PRACTICE: ph and poh (CALCULATIONS 1) EXAMPLE: What is the hydroxide ion and hydronium ion concentration of an aqueous solution that has a ph equal to 6.12? PRACTICE 1: Which of the following solutions will have the lowest concentration of hydronium ions? a) 0.100 moles C6H5NH2 b) 0.100 moles Be(OH)2 c) 0.100 moles SrH2 d) 0.100 moles (CH3)2NH PRACTICE 2: Which of the following statements about aqueous solutions is/are true? a) For an basic solution the concentration of H3O + is greater than the concentration of OH. b) The ph of a neutral aqueous solution is 7.00 at all temperatures. c) An acidic solution under normal conditions has a ph value less than 7.00. d) If the concentration of H3O + decreases then the concentration of OH will also decrease. e) The ph of aqueous solutions is less than 7. Page 14
PRACTICE: ph and poh (CALCULATIONS 2) EXAMPLE: A solution is prepared by dissolving 0.235 mol Sr(OH)2 in water to produce a solution with a volume of 750 ml. a) What is the [OH - ]? b) What is the [H + ]? PRACTICE: What is the Kw of pure water at 20.0 C, if the ph is 7.083? a) 8.26 10-8 b) 6.82 10-15 c) 7.23 10-14 d) 1.00 10-14 Page 15
CONCEPT: AUTO IONIZATION OF WATER Water can react with itself in a reaction called self ionization where and are produced. H 2 O (l) + H 2 O (l) This reaction is usually written more simply as: H 2 O (l) The equilibrium equation for water is called the (KW) for water and is given by the following: K W = [H + ][OH ] At 25 o C, KW =, but remember KW, like all other constants K, is temperature dependent. Increasing the temperature will KW. Constant 0 o C 10 o C 50 o C 100 o C K W 1.14 x 10-14 2.93 x 10-14 5.476 x 10-14 5.13 x 10-13 EXAMPLE: Determine the concentration of hydronium ions for a neutral solution at 25 o C and at 50 o C. Page 16
CONCEPT: CALCULATING ph and poh OF STRONG SPECIES STRONG ACIDS & BASES are considered Electrolytes so they ionize completely in water. HCl (aq) NaOH (aq) H 2 O H 2 O H + (aq) + Cl (aq) Na + (aq) + OH (aq) EXAMPLE 1: Calculate the ph of a 0.0782 M solution of CaH2. EXAMPLE 2: Calculate the ph of a 0.000550 M HBr solution to the correct number of significant figures. a) 3.3 b) 3.26 c) 3.260 d) 3.2596 e) All are correct PRACTICE: Calculate the ph of 50.00 ml of 4.3 x 10-7 M H2SO4. Page 17
CONCEPT: CALCULATING ph and poh OF WEAK SPECIES WEAK ACIDS & BASES are considered Electrolytes so they don t completely ionize in water. HF + H2O NH3 + H2O F (aq) + H3O + (aq) NH4 + (aq) + OH (aq) EXAMPLE 1: Pryridine, an organic molecule, is a very common weak base. C5H5N (aq) + H2O (l) C5H5NH + (aq) + OH - (g) Assume you have a 0.0225 M aqueous solution of pyridine, C5H5N, determine its ph. The Kb value for the compound is 1.5 x 10-9. Page 18
PRACTICE: CALCULATING ph and poh OF WEAK SPECIES (CALCULATIONS 1) EXAMPLE: An unknown weak base has an initial concentration of 0.750 M with a ph of 8.03. Calculate its equilibrium base constant. PRACTICE: Determine the ph of a solution made by dissolving 6.1 g of sodium cyanide, NaCN, in enough water to make a g 500.0 ml of solution. (MW of NaCN = 49.01 mol ). The Ka value of HCN is 4.9 x 10-10. Page 19
CONCEPT: ACID & BASE CONSTANTS As you might already realize, there are relatively few strong acids. The great majority of acids are weak acids. Consider a weak monoprotic acid, HA, and its ionization in water: HA (aq) + H2O (l) A (aq) + H3O + (aq) The equilibrium expression for this ionization would be: K a = Pr oducts Reactan ts = Where Ka represents the and it measures the strength of weak acids. When looking at weak bases we don t use Ka, but instead, which represents the. The relationship between Ka and Kb can be expressed with the following equation: K W = K a K b In general, the the Ka the stronger the acid and the the concentration of H +. In general, the the pka the stronger the acid and the the concentration of H +. PRACTICE: If the Kb of NH3 is 1.76 x 10-5, determine the acid dissociation constant of its conjugate acid. Page 20
PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 1) EXAMPLE 1: Knowing that HF has a higher Ka value than CH3COOH, determine, if possible, in which direction the following equilibrium lies. HF (aq) + CH3COO (aq) F (aq) + CH3COOH (aq) a) Equilibrium lies to the left. b) Equilibrium lies to the right. c) Equilibrium is equal and balanced. d) Not enough information given. EXAMPLE 2: What is the equilibrium constant for the following reaction and determine if reactants or products are favored. HCN (aq) + ClO2 (aq) CN (aq) + HClO2 (aq) The acid dissociation constant of HCN is 4.9 x 10-10 and the acid dissociation of HClO2 is 1.1 x 10-2. HCN (aq) + H2O (aq) CN (aq) + H3O + (aq) HClO2 (aq) + H2O (aq) ClO2 (aq) + H3O + (aq) Page 21
PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 2) EXAMPLE: Which of the following solutions will have the lowest ph? a) 0.25 M HC2F3O2 b) 0.25 M HIO4 c) 0.25 M HC3H5O3 d) 0.25 M H2CO3 e) 0.25 M HSeO4 PRACTICE 1: Which Bronsted-Lowry base has the greatest concentration of hydroxide ions? a) C2H8N2 (Kb = 8.3 x 10-5 ) b) C5H5N (Kb = 1.7 x 10-9 ) c) (CH3)3N (Kb = 1.0 x 10-6 ) d) C3H7NH2 (Kb = 3.5 x 10-4 ) e) C6H5NH2 (Kb = 3.9 x 10-10 ) PRACTICE 2: Which Bronsted-Lowry acid has the weakest conjugate base? a) HCNO (Ka = 2.0 x 10-4 ) b) HF (Ka = 3.5 x 10-4 ) c) HN3 (Ka = 2.5 x 10-5 ) d) H2CO3 (Ka = 4.3 x 10-7 ) Page 22
CONCEPT: ACID & BASE NEUTRALIZATION When an acid neutralizes a base an ionic compound called a is formed. These solutions can be neutral, acidic or basic, based on acid-base properties of the cations and anions formed. RULES FOR IDENTIFYING YOUR IONS CATIONS (POSITIVE IONS) 1) Transition Metals: If your transition metal has a charge of +2 or higher it is acidic. If the charge is less than +2 then it is neutral. EX: 2) Main-Group Metals: If your main-group metal has a charge of +3 or higher it is acidic. If the charge is less than +3 then it is neutral. EX: 3) Positive Amines are acidic. EX: ANIONS (NEGATIVE IONS) 1) NEGATIVE ION: If you have a negative ion then add an H + to it. If you create a weak acid then your negative ion is basic. EX: Page 23
PRACTICE: ACID & BASE NEUTRALIZATION 1 EXAMPLE: Determine if each of the following compounds will create an acidic, basic or neutral solution. a) NaOCl b) PbCl4 PRACTICE 1: Determine if each of the following compounds will create an acidic, basic or neutral solution. a) LiC2H3O2 b) C6H5NH3Br PRACTICE 2: Determine if each of the following compounds will create an acidic, basic or neutral solution. a) Co(HSO4)2 b) Sr(HSO3)2 PRACTICE 3: Determine if each of the following compounds will create an acidic, basic or neutral solution. a) C3H7NH3F Page 24
PRACTICE: ACID & BASE NEUTRALIZATION 2 EXAMPLE 1: Determine whether each compound will become more soluble in an acidic solution. a) NaBr b) LiCl c) KIO EXAMPLE 2: Determine the ph of a 0.50 M NH4Cl solution. The Kb of NH3 is 1.75 x 10-5. PRACTICE: Determine the ph of a 0.55 M NaCN solution. The Ka of hydrocyanic acid, HCN, is 4.9 x 10-10. Page 25
PRACTICE: ACID & BASE NEUTRALIZATION 2 EXAMPLE 1: Determine whether each compound will become more soluble in an acidic solution. a) NaBr b) LiCl c) KIO EXAMPLE 2: Determine the ph of a 0.50 M NH4Cl solution. The Kb of NH3 is 1.75 x 10-5. PRACTICE: Determine the ph of a 0.55 M NaCN solution. The Ka of hydrocyanic acid, HCN, is 4.9 x 10-10. Page 26
CONCEPT: DIPROTIC ACIDS Diprotic acids and bases are compounds that can donate or accept H + ion. For diprotic acids their equations can be illustrated by: H2A (aq) + H2O (l) HA (aq) + H3O + (aq) K a1 = Pr oducts Reactan ts = HA (aq) + H2O (l) A 2 (aq) + H3O + (aq) K a2 = Pr oducts Reactan ts = For diprotic bases their equations can be illustrated by: A 2 (aq) + H2O (aq) HA (aq) + OH (aq) K b1 = Pr oducts Reactan ts = HA (aq) + H2O (aq) H2A (aq) + OH (aq) K b2 = Pr oducts Reactan ts = Based on these equations the relationship between the different forms of diprotic species are: As a result of these equations for diprotic acids and bases the relationship between Ka and Kb will be: K a1 K b2 = K w K a2 K b1 = K w When dealing with diprotic acids: 1) H2A can be treated as a monoprotic acid and we use can be used to find ph. 2) HA represents the intermediate form and we use can be used to find ph. 3) A 2 represents the basic form and we use can be used to find ph. Page 27
PRACTICE: DIPROTIC ACIDS CALCULATIONS 1 EXAMPLE 1: Sulfurous acid, H2SO3, represents a diprotic acid with a Ka1 = 1.6 x 10-2 and Ka2 = 4.6 x 10-5. Calculate the ph and concentrations of H2SO3, HSO3 and SO3 2 when given 0.250 M H2SO3. EXAMPLE 2: Determine the ph of 0.115 M Na2S. Hydrosulfuric acid, H2S, contains Ka1 = 1.0 x 10-7 and Ka2 = 9.1 x 10-8. Page 28
CONCEPT: POLYPROTIC ACIDS Our understanding of diprotic acids and bases can be used to understand polyprotic acids and bases. For polyprotic acids their equations can be illustrated by: H 3A (aq) + H 2O (l) H 2A (aq) + H 3O + (aq) K a1 = Pr oducts Reactan ts = H 2A (aq) + H 2O (l) HA 2 (aq) + H 3O + (aq) K a2 = Pr oducts Reactan ts = HA 2 (aq) + H 2O (l) A 3 (aq) + H 3O + (aq) K a3 = Pr oducts Reactan ts = For polyprotic bases their equations can be illustrated by: A 3 (aq) + H 2O (l) HA 2 (aq) + OH (aq) K b1 = HA 2 (aq) + H 2O (l) H2A (aq) + OH (aq) K b2 = H2A (aq) + H 2O (l) H 3A (aq) + OH (aq) K b3 = Pr oducts Reactan ts = Pr oducts Reactan ts = Pr oducts Reactan ts = As a result of these equations for polyprotic acids and bases the relationship between Ka and Kb will be: K a1 K b3 = K w K a2 K b2 = K w K a3 K b1 = K w When dealing with polyprotic acids: H3A can be treated as a monoprotic acid and we use can be used to find ph. A 3 represents the basic form and we use can be used to find ph. H 2 A [H + ] K a1 K a2 [ ] 0 + K a1 K w K a1 + [ ] 0 HA 2 [H + ] K a2k a3 [ ] 0 + K a2 K w K a2 + [ ] 0 Page 29
PRACTICE: POLYPROTIC ACIDS CALCULATIONS EXAMPLE 1: Determine the ph of 0.300 M sodium hydrogen phosphate, Na2HPO4. Phosphoric acid, H3PO4, contains Ka1 = 7.5 x 10-3, Ka2 = 6.2 x 10-8 and Ka3 = 4.2 x 10-13. EXAMPLE 2: Determine the ph of 0.300 M citric acid, H3C6H5O7 it possesses Ka1 = 7.4 x 10-4, Ka2 = 1.7 x 10-5 and Ka3 = 4.0 x 10-7. Page 30