Acids and Bases Chemistry PDF

Acids and Bases Chapter 15: Acids and Bases Arrhenius Definitions: ◆ acids - compounds that produce an increase in...

Acids and Bases Chapter 15: Acids and Bases Arrhenius Definitions: ◆ acids - compounds that produce an increase in [H+] when dissolved in water ◆ bases - compounds that produce an increase in [OH–] when dissolved in water Lewis Definitions: ◆ acids - electron pair acceptors ◆ bases - electron pair donors Brønsted-Lowry Definitions: ◆ acids - H+ donors ◆ bases - H+ acceptors Lewis Acids & Lewis Bases Lewis Acids & Lewis Bases ◆ more broad way to define acids and bases ◆ some examples: ◆ Lewis acids – electron pair acceptors Al3+ + n H2O ! [Al(H2O)n]3+ metal cations (Mn+) and boron are common Lewis acids Cu2+ + n NH3 ! [Cu(NH3)n]2+ species that are electron deficient; electrophiles BF3 + NH3 ! F3B–NH3 ◆ Lewis bases – electron pair donors species with O, N, halogen frequently have lone ◆ acidic oxides (oxides of nonmetals): pairs of electrons to share ∴ Lewis bases SO3 + H2O ! H2SO4 species that are electron rich; nucleophiles ◆ product of a Lewis Acid + Lewis Base reaction is called ◆ basic oxides (oxides of metals): a Lewis Acid-Base adduct CaO + H2O ! Ca2+ (aq) + 2 OH– (aq)  Brønsted-Lowry Acids & Bases Brønsted-Lowry Acids & Bases ◆ Brønsted-Lowry acids – H+ donors ◆ writing ionization (or dissociation) equations to describe Brønsted-Lowry acid/base behavior in ◆ Brønsted-Lowry bases – H+ acceptors aqueous solutions: ◆ reaction of a Brøsted-Lowry acid + base is a ◆ acid ionization (or dissociation) equation: neutralization reaction characterized by H+ transfer HA (aq) + H2O (l) ! A– (aq) + H3O+ (aq) acid base conjugate hydronium example of neutralization reaction: base ion HCl (aq) + NaOH (aq) ! NaCl (aq) + H2O (l) acid base salt water ◆ base ionization (or dissociation) equation: net ionic equation: B (aq) + H2O (l) ! BH+ (aq) + OH– (aq) base acid conjugate hyroxide H+ (aq) + OH– (aq) ! H2O (l) acid ion  Brønsted-Lowry Acids & Bases Brønsted-Lowry Acids & Bases HA (aq) + H2O (l) ! A– (aq) + H3O+ (aq) some specific examples: acid base conjugate hydronium acid ionization equation for nitrous acid: base ion HNO2 (aq) + H2O (l) ! NO2– (aq) + H3O+ (aq) notes: HNO2 acid H2O base B (aq) + H2O (l) ! BH+ (aq) + OH– (aq) * H3O conjugate acid * + base acid conjugate hyroxide NO2– conjugate base acid ion * H3O+ & H+ are used some terminology: interchangeably amphoteric: a species that can act as an acid or a base water is an example of an amphoteric species base ionization equation for ammonia: conjugate base: species that remains after an acid NH3 (aq) + H2O (l) ! NH4+ (aq) + OH– (aq) donates its H+ notes: H2O acid NH3 base NH4 conjugate acid + OH– conjugate base conjugate acid: species that forms after a base accepts a H+ Strong vs. Weak Acids and Bases Identify each species in the following equation as Acid and base strength is based on the extent of etiher the Brønsted-Lowry acid, the Brønsted-Lowry ionization that occurs when the substance is base, the conjugate acid, or the conjugate base. dissolved in water. Strong Acids: Identify the conjugate acid-base pairs in the reaction. ◆ strong electrolytes - completely ionized in solution H2SO4 (aq) + HPO42– (aq) ! HSO4– (aq) + H2PO4– (aq) ◆ there are 6 strong acids - KNOW THEM! HCl, HBr, HI, HNO3, HClO4, H2SO4 (diprotic) Weak Acids: ◆ weak electrolytes - partially ionized (typically < 5%) in aqueous solution ◆ any acid that is not a strong acid is a weak acid some examples: HF, H2CO3, H3PO4, HNO2, HBrO4  Strong vs. Weak Acids and Bases Strong vs. Weak Acids and Bases Strong Bases: ◆ strong electrolytes - completely ionized in solution ◆ the strong bases are the hydroxides of the alkali metals & hydroxides of most alkaline earth metals; KNOW THEM! LiOH, NaOH, KOH, RbOH, CsOH Ca(OH)2, Sr(OH)2, Ba(OH)2 Weak Bases: ◆ weak electrolytes - partially ionized (typically < 5%) in aqueous solution ◆ weak bases tend to be organic compounds that contain nitrogen; ammonia and substituted amines some examples: NH3, (CH3)NH2, (CH3)3N C5H5N, N2H4, NH2OH Weak Acids and Weak Bases: Weak Acids and Acid Ionization Constant, Ka Reversible H+ Transfer Reactions HA (aq) + H2O (l) ⇄ A– (aq) + H3O+ (aq) ◆ In Chapter 4 we defined weak acids and weak bases as weak electrolytes (only partially ionized in [H3O+][A–] aqueous solution). Ka = –––––––––– [HA] ◆ Now we can talk about their behavior in terms of an equilibrium that exists in solution: ◆ Ka is the acid ionization constant HA (aq) + H2O (l) ⇄ A– (aq) + H3O+ (aq) ◆ the larger the value of Ka... the equilibrium position lies farther to the right B (aq) + H2O (l) ⇄ BH+ (aq) + OH– (aq) higher [H3O+] greater extent of ionization ◆ These are heterogeneous equilibria. stronger acid ◆ We will discuss/define equilibrium constants, Ka & Kb.  Weak Bases and Base Ionization Constant, Kb B (aq) + H2O (l) ⇄ BH+ (aq) + OH– (aq) [BH+][OH–] Kb = –––––––––– [B] ◆ Kb is the base ionization constant ◆ the larger the value of Kb... the equilibrium position lies farther to the right higher [OH–] greater extent of ionization stronger base Relationship Between Strengths in Conjugate Acid/Base Pairs ◆ the stronger an acid, the weaker its conjugate base ◆ the weaker an acid, the stronger its conjugate base ◆ the stronger a base, the weaker its conjugate acid ◆ the weaker a base, the stronger its conjugate acid Relationship Between Structure and Strengths of Acids Relationship Between Structure and Strengths of Acids: Binary Acids (HA) ◆ Brønsted-Lowry acids are H+ donors... so... ◆ For a set of binary acids in which A belongs to the acid strength is dependent on how readily same group of the periodic table, H–A bond strength donated the acidic H+ is is the determining factor in acid strength. ◆ the weaker the interaction between A–H (in the stronger the H–A bond, the weaker the acid binary acids) or O–H (in oxoacids), the stronger the acid ◆ H–A bond strength is related to atomic size: bond strength decreases as atomic radius ◆ the stronger the interaction between A–H (in increases binary acids) or O–H (in oxoacids), the weaker the acid atomic radius increases moving down the periodic table Relationship Between Structure and Strengths of Acids: Relationship Between Structure and Strengths of Acids: Binary Acids (HA) Binary Acids (HA) ◆ For a set of binary acids in which A is in the same group/period of A group VIA group VIIA period of the periodic table, H–A bond polarity is the H2O HF determining factor in acid strength. 2nd period Ka = 1 x 10–14 Ka = 6.8 x 10–4 HA bond strength the more polar the H–A bond, the stronger the acid 3rd period H2S HCl decreases Ka = 9 x 10–8 Ka very large ◆ H–A bond polarity depends on the electronegativity H2Se HBr HA acid 4th period of A: Ka = 1.3 x 10–4 Ka very large strength increases bond polarity increases as the H2Te HI 5th period electronegativity of A increases Ka = 2.3 x 10–3 Ka very large electronegativity increases moving left to Electronegativity of A increases HA bond polarity increases right across the periodic table HA acid strength increases Relationship Between Structure and Strengths of Acids: Carboxylic Acids: Oxoacids (HnAOm) O–H Bond Polarization and Acid Strength ◆ For a set of oxoacids with the same number of O’s, the acid strength increases as the electronegativity of A increases. ◆ if A is more electronegative, it pulls electron density toward itself resulting in a more polarized O–H bond the more polar the O–H bond, the stronger the acid ◆ acetic acid (CH3COOH) has Ka = 1.8 x 10–5 ◆ How will the acid strength change as 1, 2 or 3 H’s are replaced with F? With Cl? Relationship Between Structure and Strengths of Acids: Relationship Between Structure and Strengths of Acids: Oxoacids (HnAOm) Oxoacids (HnAOm) acetic acid CH3COOH ◆ For a set of oxoacids with the same atom A, the acid Ka = 1.8 x 10–5 strength increases as the number of O’s increases. HOI monochloroacetic monofluoroacetic acid ◆ As the number of electronegative O’s in the acid !I = 2.5 CH2FCOOH molecule increases, the net effect is that electron CH2ClCOOH Ka = 2.3 x 10–11 Ka = 2.5 x 10–3 Ka = 1.4 x 10–3 density is pulled away from H resulting in a more HOBr dichloroacetic acid polarized O–H bond. !Br = 2.8 CHCl2COOH the more polar the O–H bond, the stronger the acid Ka = 2.0 x 10–9 Ka = 5.5 x 10–2 HOCl trifluoroacetic acid trichloroacetic acid acid: HClO HClO2 HClO3 HClO4 !Cl = 3.0 CF3COOH CCl3COOH Ka = 10 Ka = 0.23 Ka = 3.5 x 10–8 1.2 x 10–2 ∼1 v. large Ka = 3.5 x 10–8 Auto-Ionization of Water and KW Acidic, Basic & Neutral Aqueous Solutions ◆ recall that water is amphoteric - can act as an acid or a base ◆ distinguish between acidic, basic and neutral solutions based on the relative [H3O+] & [OH–] ◆ now consider a reaction between 2 water molecules: if [H3O+] > [OH–], solution is acidic H2O (l) + H2O (l) ⇄ H3O+ (aq) + OH– (aq) if [OH–] > [H3O+], solution is basic ◆ this is called the auto-ionization of water if [H3O+] = [OH–], solution is neutral heterogeneous equilibrium KW = [H3O+][OH–] ◆ for a neutral solution at 25°C: at 25°C, KW = 1.0 x 10–14 [H3O+] = [OH–] = 1.0 x 10–7 M In any aqueous solution at 25°C: [H3O+][OH–] = KW = 1.0 x 10–14  The pH Scale example: In a sample of lemon juice, [H3O+] = 2.5 x 10–3 M. logarithmic scale of [H3O+] in solution Calculate the [OH–], and classify lemon juice as an pH = !log[H3O+]; [H3O+] = 10–pH acidic, basic or neutral solution. example: At 50°C, KW = 5.5 x 10–14. Determine [H3O+] and [OH–] in a neutral solution at 50°C. pH Calculations: Relationship Between [H3O+] and pH Relative Acidity and Basicity of Solutions ◆ as [H3O+] changes by a factor of 10, the pH of the recall: solution changes by 1 unit ◆ in any aqueous solution at 25°C: ◆ higher [H3O+] corresponds to lower pH [H3O+][OH–] = 1 x 10–14 ◆ higher [H3O+] corresponds to more acidic solution ◆ pH = –log[H3O+]; [H3O+] = 10–pH ◆ higher [H+] " more acidic solution " lower pH ◆ higher [OH–] " lower [H+] " more basic solution " higher pH  pH Calculations: Other Logarithmic Quantities Relative Acidity and Basicity of Solutions ◆ pOH = – log [OH–] [OH–] = 10–pOH ex: Calculate the pH of 0.00283 M HNO3 (aq). the higher the [OH–], the lower the pOH as [OH–] changes by factor of 10, the pOH ex: Will the pH of 0.00283 M HNO2 (aq) be less changes by 1 unit than, greater than, or equal to the pH of 0.00283 M HNO3 (aq)? Why? ◆ pKa = – log Ka Ka = 10–pKa the larger the Ka of an acid, the smaller the pKa ex: Calculate the [H3O+] in a sol’n with pH = 3.61. ◆ pKb = – log Kb Kb = 10–pKb ex: Calculate the pH of 0.20 M Ba(OH)2 (aq) and the larger the Kb of a base, the smaller the pKb 0.20 M NaOH (aq). Should they be the same? Why or why not? ◆ pKW = – log KW at 25°C, pKW = 14.00