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Lecture 12, Chemistry of Solutions Notes

by: CatLover44

Lecture 12, Chemistry of Solutions Notes 202-NYB-05

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These notes cover 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.
Chemistry of Solutions
Nadia Schoonhoven
Class Notes
Oxyacids, electronegativity
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This 12 page Class Notes was uploaded by CatLover44 on Friday October 14, 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 7 views. For similar materials see Chemistry of Solutions in Chemistry at Dawson Community College.


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Date Created: 10/14/16
Chemistry of Solutions Lecture 12 Thursday, October 4, 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, A salt’s cation and anion both influence a solution’s pH. 1) Write an equation for an equilibrium reaction of each ion of H O, which can act 2 as an acid or a base. 2) Find an equilibrium constant for each reaction (a or 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 CO 3- + H 2 ⟷ OH + HCO- 3- + - [H ] vs [OH ], are we dealing with an acid or a base? + -10 K af NH 4 is 5.6x10 - -11 K bf HCO is 3.6x10 Since K > K , the basic reaction is dominant. b a -14 -11 K b1= K /wK = a.0x10 / 5.6x10 -4 K b1= 1.8x10 For K , just determine which is greater, K or K . This reaction has a greater K , so a2 a b b - [OH ] increases. - - (b) F + H O2⟷ HF + OH K = K x K , K / K = 1.0x10 -14/ 7.2x10 = 1.4x10 -11, so this solution is only w a b w a slightly acidic, almost neutral. Exercise: a 0.10 M aqueous solution of H PO will b3 ac4dic, basic, or neutral? 2 -3 H 3O ,4K a1= 7.5x10 - -8 H 2O ,4K a2 = 6.2x10 2- -13 HPO 4 , Ka3 = 4.8x10 Acidic Reaction: H PO + H O ⟷ H O + PO + 3-, K = 4.8x10 -13 3 4 2 3 4 a Basic Reaction: - - -8 H 3O +4H O ⟷ 2H + H PO , K = 6.2x104 b - K tells us whether more hydrogen or hydroxide ions are produced in a given reaction. - Which K doawe use to calculate K ? The K ob the secona dissociation. - -14 -8 -7 - Use K a2 from H P2 : 140x10 / 6.2x10 = 1.6x10 = K , which ib greater than K , so the solution is basic. a - 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 ). 3 - 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 3 CH ,4CH NO3are 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. 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. 4 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. - 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. 5 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. - Coordinated bonding where both electrons are supplied by one atom = abducts. 6 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. 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 (2n H ) is attracted to the partial negative charge on the O atom in H O. K 2ecreasas with decreasing electronegativity, and increases with increasing electronegativity, so electronegativity is directly related to acidity. 7 - The more polar a molecule is, the greater its K , and 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 of electron-withdrawing groups, which are the neighbouring electronegative atoms in a molecule that pull on electrons. - - o CCl C3O has a stronger inductive effect than CH COO ; electro3 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 . + 8 - 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 d onor increases. Acid-base Properties of Oxides - H – O – X can also act as bases. 9 - 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. - 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 10 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 11 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 electronegativit y 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. 12


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