1 / 15

CHAPTER 19 ACIDS & BASES

CHAPTER 19 ACIDS & BASES. The pH scale. The definition of pH is defined as the negative logarithm of the hydrogen ion-concentration. Expressed mathematically, pH is shown as: pH=- log[H + ]. The pH scale ranges from 0 to 14.

taite
Download Presentation

CHAPTER 19 ACIDS & 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. CHAPTER 19 ACIDS&BASES

  2. The pH scale • The definition of pH is defined as the negative logarithm of the hydrogen ion-concentration. • Expressed mathematically, pH is shown as: pH=-log[H+]. • The pH scale ranges from 0 to 14. • A pH of 7 is considered neutral. Because the number of hydrogen ions equal the number of hydroxide ions. • A solution with a pH of 0 is considered to be highly acidic, and a solution with a pH of 14 is highly basic or alkaline.

  3. pH scale continued…. • Along with ranging from 0 to 14, the pH scale also corresponds with it’s numbers by having different colors for acids and bases. • Acids usually range from red to light orange while bases range from yellowish-green to blue.

  4. The Properties of Acids • The properties of acids include: having a sour taste, they change color of an acid-base indicator, they react with compounds containing hydroxide ions to form water and salt, some react with metals to produce hydrogen (H2) gas which can be useful test in the lab, and they can be strong or weak electrolytes in aqueous solution.

  5. Strong Acids • A strong acid is an acid that ionizes completely in aqueous solutions. • Such as: HCl(g)+H20(l)H3O+(aq)+ Cl-(aq) • Strong acids are also a weak conjugate base. • Because if an acid is good at donating a H+ ion, it has a difficult time getting it back. • Some common strong acids include: Hydrochloric acid, Nitric acid, and Sulfuric acid.

  6. Strong Bases • Bases are considered strong bases when they disassociate completely into metal ions and hydroxide ions in aqueous solutions. • Some common strong bases are Potassium hydroxide, and sodium hydroxide.

  7. Hydroniumvs Hydroxide ion • The Hydronium ions is formed when a water molecule gains a hydrogen ion. Therefore, the resulting charge of the new molecule is positive. [H3O+] • Inversely, the Hydroxide ion is formed when water loses a hydrogen and becomes a negatively charged ion. [OH-]

  8. Arrenhius acids and bases • Arrenhius defined an acid as hydrogen-containing compounds that ionize to yield hydrogen ions [H+] in aqueous solution. He also said that bases are compounds that ionize to yield hydroxide ions [OH-] in aqueous solution. • Identify one of the following as an Arrenhius base: LiOH, NH3, H2PO4, or CH3COOH

  9. Acid vs Base on pH scale • On the pH scale the pH of acids are lower than the pH of bases • Acids range from 0 to 6. While bases range from 8 to 14. • Acids also turn an acid-base indicator from dark red to orange depending on the strength of the acid. • Bases turn the color of an acid-base indicator from green to blue, depending on how strong the base is.

  10. Lewis Acid vs. Lewis Base • A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. • A Lewis base is a substance that can donate a pair of electrons to form a covalent bond. • An example of a Lewis base is in the reaction H++O—HH—O—H. • The hydroxide is a Lewis base in this reaction.

  11. Bronsted-Lowry acid vs base • The Bronsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogen-ion acceptor. • Consequently, when a Bronsted-Lowry acid loses a hydrogen ion, it becomes a conjugate base. • Also, when a Bronsted-Lowry base accepts a hydrogen ion, it forms the conjugate acid of the base. • Find the conjugate acid-base pair in the following reaction: NH3+H2O  NH4+ + OH-

  12. Finding [OH-] when given [H+] • If the [H+] in a solution is 1.0 X 10-5M , is the solution acidic, basic, or neutral? What is the [OH-] of this solution? • Given: [H+] 1.0x10-5M • Find: [OH-] and acidic basic or neutral • Use: Kw=[H+] X [OH-]=1.0 X 10-14 • Plug-in: [OH-]= 1.0 X 10-14/ 1.0 X 10-5 • Solve: [OH-]= 1.0 X 10-9 the solution is acidic

  13. Now do one yourself…. • The [H+] in a solution is 2.4 X 10-7M, what is the [OH-] of the solution?

  14. Finding pH given [OH-] • What is the pH of a solution if [OH-]= 4.0 X 10-11M • Given: [OH-]= 4.0 X 10-11 • Find: pH • Use: pH= -log[H+] and Kw=[H+] X [OH-]=1.0 X 10-14 • Plug-In: [H+]= 1.0 X 10-14/4.0 X 10-11 • Solve: [H+]= 2.5 X 10-4M pH=-log[2.5 X 10-4M] • pH= 3.60

  15. Now do one yourself… • Calculate pH the if [OH-]=4.3 X 10-5M

More Related