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Lecture 13, Chemistry of Solutions notes

by: CatLover44

Lecture 13, Chemistry of Solutions notes 202-NYB-05

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Topics covered: localized and delocalized electron pairs, and equilibrium concentration / pH problems.
Chemistry of Solutions
Nadia Schoonhoven
Class Notes
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This 13 page Class Notes was uploaded by CatLover44 on Saturday October 15, 2016. The Class Notes belongs to 202-NYB-05 at Dawson Community College taught by Nadia Schoonhoven in Fall 2016. Since its upload, it has received 4 views. For similar materials see Chemistry of Solutions in Chemistry at Dawson Community College.


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Date Created: 10/15/16
Chemistry of Solutions Lecture 12 Thursday, October 6, 2016 Topics Covered: acid-base properties of salts, effect of structure on acidity, oxyacids, hydrated metal ions, electronegativity, acid-base properties of oxides, dissolution of oxides in water. Acid-Base Properties of Salts: a salt’s cation and anion can individually affect a solution’s pH. - In salts where both ions have basic or acidic properties, such as NH C H 4 ,2 3 2 their properties depend on whether K or Kais labger. If K < K ,athe sblt is basic. If K a K ,bthe salt is acidic. If K a K , bt's a neutral salt. - How to approach these types of problems: 1) Write an equation for an equilibrium reaction of each ion of H O, w2ich can act as an acid or a base. 2) Find an equilibrium constant for each reaction (K or K ). a b 3) Which reaction is more product-favoured? Compare the K’s. The reaction with the larger K will dominate. 4) Will the solution be acidic or basic when it's dissolved in water? (a) Ammonium carbonate: + + NH 4 ⟷ H + NH 3 2- - - CO 3 + H O2⟷ OH + HCO 3 + - [H ] vs [OH ], are we dealing with an acid or a base? Kaof NH 4 is 5.6x10 -10 - -11 Kbof HCO is35.6x10 Since K b K , ahe basic reaction is dominant. -14 -11 Kb1 = K w K =a1.0x10 / 5.6x10 Kb1 = 1.8x10 -4 2 For K ,a2ust determine which is greater, K or K . Thas reacbion has a greater K , so b - [OH ] increases. - - (b) F + H O 2 HF + OH -14 -4 -11 K w K x a , K b K =w1.0xa0 / 7.2x10 = 1.4x10 , so this solution is only slightly acidic, almost neutral. Exercise: a 0.10 M aqueous solution of H PO will be acidic, basic, or neutral? 3 4 -3 H 3O ,4K a1 = 7.5x10 H 2O ,4K a2 = 6.2x10 -8 2- -13 HPO 4 , Ka3 = 4.8x10 Acidic Reaction: + 3- -13 H 3O +4H O ⟷ 2 O + PO 3 4 , K a 4.8x10 Basic Reaction: - - -8 H 3O +4H O ⟷ 2H + H PO , K = 6.2x10 4 b 3 - K tells us whether more hydrogen or hydroxide ions are produced in a given reaction. - Which K doawe use to calculate K ? The Kbof the secand dissociation. - -14 -8 -7 - Use K a2from H P2 : 140x10 / 6.2x10 = 1.6x10 = K , which bs greater than K , ao the solution is basic. - Where does an acid-base reaction’s equilibrium lie? This depends on how strong an acid (or base) is. - In acid-base reactions, the anion (base) and water are competing for the + proton (H ). - At equilibrium, more molecules of the stronger base are bonded to the proton. - What makes a proton acidic or basic? o Bond polarity o Bond strength - CHCl ,3CH , 4H NO 3re 2ll basic (weak bases). C – H bonds are non-polar and very strong. - Bond strength (most to least polar): H – F >> H – Cl > H – Br > H – I. - Based on polarity alone, we would never expect HF to be the strongest acid. - HF’s bond strength and its ability to form hydrogen bonds with itself accounts for its relatively low acidity. 4 o HF is polar. o F is a small atom that overlaps a lot with H. o Cl is a much larger atom with less overlap, and forms weaker bonds. Since Cl bonds are easier to break, they form stronger acids. o In the periodic table, acid strength increases as you go down the columns. - Bond strength is important when we compare acidities of molecules having the + acidic H attached to elements in the same group (column) in the periodic table. For example, compare H – F to H – Br, and H – OH to H – SH. - Bond polarity is important when comparing the acidities of molecules having the acidic H attached to elements in the same period (row) of the periodic table. Oxyacids - Oxyacids have the following form: H – O – X. 5 - Their strength depends on the number of oxygen atoms attached to the central atom. The more oxygen atoms there are, the stronger the acid is. Lewis Acid-Base Theory - partial positive and partial negative charges interact, orbitals overlap with each other, and electrons are shared between atoms. This explains many types of reactions, includes the transfer of H+. - Lewis bases are electron pair acceptors (have a pair to share). Examples are water and alcohols, ammonia, and amines. Anything with a lone pair is a lone pair donor. + - Lewis acids are electron pair donors. H has an empty 1s orbital (need two from you), other examples are Be, B, and Al (valence shells aren't full). Metal cations have a positive charge (all metals lose electrons). Lewis acids are anything electron deficient. 6 - Coordinated bonding where both electrons are supplied by one atom = abducts. Hydrated Metal Ions - A strong positive charge on the metal central atom weakens the O – H bonds. - Solutions of highly-charged metal cations are acidic due to Lewis acid-base interactions of Mn and H O. 2 - Mn pulls on an oxygen atom’s lone pair (using its coordinate bond); this makes oxygen pull harder on the O – H bonding pair, making the bond more + polar; H has an affinity to incoming bases (partial positive charge). - Coordinated water molecule (pKa ~ 3-7) are more acidic than free water molecules, pKa = 14. - Hydrolysis occurs when solvent-water deprotonates a coordinated water molecule. 7 3+ 3+ 2+ 2+ - This is true for highly-charged cations, like Al , Fe , Cu , and Pb . Not all alkali earth ions are highly-charged. - Note: as electronegativity increases, acidity of the molecule increases . - Why are carboxyl acids weak? Because O – H bonds are extremely polar, and the partial positive charge of H O (on H ) is attracted to the partial 2 negative charge on the O atom in H O. K decreases with decreasing 2 a electronegativity, and increases with increasing electronegativity, so electronegativity is directly related to acidity. - The more polar a molecule is, the greater its K , ana the more electronegative an atom is, the greater its K . a - But they're relatively strong weak acids because their conjugate bases are weak bases. The lone pair causing the negative charge is on a highly electronegative atom, and it is delocalized by resonance. Every oxygen has an average charge of -½. - If a base has a resonance form, it's a weak base (it’s resonance stabilized). The base is less reactive, so it's weak. - When comparing two carboxylic acids, why do they have such differences in their acidities? The inductive effect explains this. The inductive effects 8 of electron-withdrawing groups, which are the neighbouring electronegative atoms in a molecule that pull on electrons. o CCl COO has a stronger inductive effect than CH COO ; electron - 3 3 density is pulled from the rest of the molecule towards the Cl atoms; this makes the O atoms pull harder on the extra lone pair of electrons; atoms hold on tightly to their lone pairs, so they are less reactive; this makes the base weaker, since the lone pair isn't being + donated, so it can't compete as well for H . + - If the conjugate base is too weak to get the H back, the equilibrium often favours deprotonation, so the acid is relatively strong . - How strong are common bases that have a lone pair on their O or N atoms? o Basicity increases as the average charge of the proton donor increases. 9 Acid-base Properties of Oxides - H – O – X can also act as bases. + - Acidic oxides: the H – O bond from H – O – X can be broken to form H and O – X. In acidic oxides, O – X is held together with a covalent bond. This occurs when X has a similar electronegativity to O, or a high electronegativity. Other examples are C, N, O, S, and Cl. 10 - Basic oxides: occur when the O – X bond is ionic; when X has a very low electronegativity like in Na, Li, or Ca atoms. - Note: electronegativity increases as you read the periodic table from left to right. - When determining which base is stronger, the first step is always to look + at what H is attached to. Dissolution of Oxides in Water Summary 11 Useful Chart to Help You Compare Base Strengths Exercise: which of the following bases is the weakest: hydroxide (OH ), carbonate - 2- - - (CO 3 ), acetate (CH CO3) o2 amide (NH )? Use t2e chart! 1) OH - Electronegativity trend: atoms become more electronegative as you read the periodic table from left to right. O is more electronegative than H, whose electronegativity value equals that of P. Since acidity increases with electronegativity, this is a strong base. Its conjugate acid is water, a very weak acid (or base depending on what reaction you're dealing with). 2- 2) CO 3 12 C and O are in the same row on the periodic table, so we can't compare bond strength, but we can determine its strength by using the electronegativity trend. O 2- is more electronegative than C, so O wants to keep its electrons. This means CO 3 is a weak base. The carbonate ion is also the conjugate base of the strong acid H 2O ,3so this is a very weak base. It also has a resonance form, so it's resonance stabilized which also tells us that it's a weak base (no base strength). 3) CH CO - 3 2 The acetate ion isn't the conjugate base of a strong acid, so it's a stronger base than the carbonate ion. It's a weak base because its resonance is delocalized over many atoms. Each O atom has an average charge of -½ , and the weaker an average charge is, the weaker the base is. - 4) NH 2 N and H are different sizes (N is larger than H). The larger atom is less basic, so this is a weak base. The weakest base is the carbonate ion. 13


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