1 / 31

Arrhenius Acids and Bases

Arrhenius Acids and Bases. In 1884, Svante Arrhenius proposed these definitions. Acid: A substance that dissolves in water to produce H 3 O + ions. Base: A substance that dissolves in water to produce OH - ions.

jody
Download Presentation

Arrhenius Acids and Bases

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Arrhenius Acids and Bases • In 1884, Svante Arrhenius proposed these definitions. • Acid: A substance that dissolves in water to produce H3O+ ions. • Base: A substance that dissolves in water to produce OH- ions. • This definition of an acid is a slight modification of the original Arrhenius definition, which was that an acid produces H+ in aqueous solution. • Today we know that H+ reacts immediately with a water molecule to give a hydronium ion H3O+.

  2. Arrhenius Acids and Bases • When HCl, for example, dissolves in water, it reacts with water to give a hydronium ion and a chloride ion. • We use curved arrows to show the change in position of electron pairs during this reaction.

  3. Arrhenius Acids and Bases • With bases, the situation is slightly different. • Many bases are metal hydroxides such as KOH, NaOH, Mg(OH)2, and Ca(OH)2. • These compounds are ionic solids and, when they dissolve in water, their ions merely separate. • Other bases are not hydroxides; these bases produce OH- by reacting with water molecules.

  4. Arrhenius Acids and Bases • We use curved arrows to show the transfer of a proton from water to ammonia.

  5. Brønsted-Lowry Acids & Bases • Acid: A proton donor. • Base: A proton acceptor.

  6. Conjugate Acids & Bases • Conjugate base: The species formed form on acid when an acid donates a proton to a base. • Conjugate acid: The species formed from a base when the base accepts a proton from an acid. • Acid-base reaction: A proton-transfer reaction. • Conjugate acid-base pair: Any pair of molecules or ions that can be interconverted by transfer of a proton.

  7. Brønsted-Lowry Acids & Bases • Brønsted-Lowry definitions do not require water as a reactant.

  8. Brønsted-Lowry Acids & Bases • We use curved arrows to show the flow of electrons that occurs in the transfer of a proton from acetic acid to ammonia.

  9. Brønsted-Lowry Acids & Bases • Note the following about the conjugate acid-base pairs in Table 2.1: • An acid can be positively charged, neutral, or negatively charged; examples of each type are H3O+, H2CO3, and H2PO4- . 2. A base can be negatively charged or neutral; examples are OH-, Cl-, and NH3. 3. Acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons that each may give up; examples are HCl, H2CO3, and H3PO4.

  10. Brønsted-Lowry Acids & Bases • Carbonic acid, for example, can give up one proton to become bicarbonate ion, and then the second proton to become carbonate ion. 4. Several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as both an acid and as a base.

  11. Brønsted-Lowry Acids & Bases 5. There is an inverse relationship between the strength of an acid and the strength of its conjugate base. • The stronger the acid, the weaker its conjugate base. • HI, for example, is the strongest acid in Table 2.1 and its conjugate base, I-, is the weakest base in the table. • CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3- (bicarbonate ion).

  12. Acid and Base Strength • Strong acid: One that reacts completely or almost completely with water to form H3O+ ions. • Strong base: One that reacts completely or almost completely with water to form OH- ions. • Here are the six most common strong acids and the four most common strong bases.

  13. Acid and Base Strength • Weak acid: A substance that only partially dissociates in water to produce H3O+ ions. • Acetic acid, for example, is a weak acid; in water, only 4 out every 1000 molecules are converted to acetate ions. • Weak base: A substance that only partially dissociates in water to produce OH- ions. • ammonia, for example, is a weak base

  14. Acid-Base Reactions • Acetic acid is incompletely ionized in aqueous solution. • The equation for the ionization of a weak acid, HA, is

  15. Acid-Base Equilibria • To determine the position of equilibrium in an acid-base reaction: • Identify the two acids in the equilibrium; one on the left and one on the right. • Use the information in Table 2.2 to determine which is the stronger acid and which is the weaker acid. • Remember that the stronger acid gives the weaker conjugate base, and the weaker acid gives the stronger conjugate base. • The stronger acid reacts with the stronger base to give the weaker acid and weaker base. • Equilibrium lies on the side of the weaker acid and the weaker base.

  16. Acid-Base Equilibrium • Equilibrium in the following acid-base reaction lies to the right, on the side of the weaker acid and the weaker base.

  17. Structure and Acidity • The most important factor in determining the relative acidity of an organic acid is the relative stability of the anion, A-, formed when the acid, HA, transfers a proton to a base. • We consider these four factors: 1. The electronegativity of the atom bonded to H in HA. 2. Resonance stabilization of A-. 3. The inductive effect. 4. The size and delocalization of charge in A- .

  18. Structure and Acidity • Electronegativity of the atom bearing the negative charge. • Within a period • The greater the electronegativity of the atom bearing the negative charge, the more strongly its electrons are held. • The more strongly the electrons are held, the more stable the anion A-. • The more stable the anion A- the greater the acidity of the acid HA.

  19. Structure and Acidity • Resonance delocalization of the charge on A- • Compare the acidity of a carboxylic acid and an alcohol, both of which contain an -OH group. • Carboxylic acids are weak acids. Values of pKa for most unsubstituted carboxylic acids fall within the range of 4 to 5. • Alcohols are very weak acids. Values of pKa for most alcohols fall within the range of 15 to 18. • How do we account for the fact that carboxylic acids are stronger acids than alcohols?

  20. Structure and Acidity • The greater the resonance stabilization of the anion, the more acidic the compound. • There is no resonance stabilization in an alkoxide anion. • We can write two equivalent contributing structures for the carboxylate anion; the negative charge is spread evenly over the two oxygen atoms.

  21. Structure and Acidity • Inductive polarization of electron density transmitted through covalent bonds by a nearby atom of higher electronegativity

  22. Structure and Acidity • The more stable the anion A-, the greater the acidity of the acid HA. • The larger the volume over which the charge on an anion (or cation) is delocalized, the greater the stability of the anion (or cation). • When considering the relative acidities of the hydrogen halides, (HI > HBr > HCl > HF), we need to consider the relative stabilities of the halide ions. • Recall from general chemistry that atomic size is a periodic property. • For main group elements, atomic radii increase going down a group and increase going across a period.

  23. Structure and Acidity • For the halogens, iodine has the largest atomic radii, fluorine has the smallest I > Br > Cl > F. Anions are always larger than the atoms from which they are derived. For anions, nuclear charge is unchanged but the added electron(s) introduce new repulsions and the electron clouds swell. Among the halide ions, I- has the largest atomic radius, and F- has the smallest atomic radius. Thus, HI is the strongest acid because the negative charge on iodide ion is delocalized over a larger volume than the negative charge on chloride, etc. Note that the result of both resonance and the inductive effect is due to the delocalization of charge.

  24. Lewis Acids and Bases • Lewis acid: Any molecule or ion that can form a new covalent bond by accepting a pair of electrons. • Lewis Base: Any molecule or ion that can form a new covalent bond by donating a pair of electrons.

  25. Lewis Acids and Bases When HCl dissolves in water, for example, the strongest available Lewis base is an H2O molecule, and the following proton-transfer reaction takes place. When HCl dissolves in methanol, the strongest available Lewis base is a CH3OH molecule, and the following proton-transfer reaction takes place.

  26. Lewis Acids and Bases

  27. Lewis Acids and Bases • Another type of Lewis acid we will encounter in later chapters is an organic cation in which a carbon is bonded to only three atoms and bears a positive formal charge. • such carbon cations are called carbocations

  28. Acids and Bases End Chapter 2

More Related