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Acid Base Rxns. Reaction with an “active metal” and an acid. Mg + 2HCl H 2 + MgCl 2. Mg + HCl. Zn + 2HCl H 2 + ZnCl 2. Zn + HCl. Al + HCl. 2Al + 6HCl 3H 2 + 2AlCl 3. Write the balanced equation for the reaction that occurs between. What type of Reaction?.

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write the balanced equation for the reaction that occurs between

Mg + 2HCl H2 + MgCl2

Mg + HCl

Zn + 2HCl H2 + ZnCl2

Zn + HCl

Al + HCl

2Al + 6HCl 3H2 + 2AlCl3

Write the balanced equation for the reaction that occurs between
what type of reaction
What type of Reaction?

Single Replacement Reactions

neutralization reactions
Neutralization Reactions

Despite heavy fire, McAllister’s aim

was true, and his carefully measured

hydrochloric acid found its mark in

the enemy’s reservoir of sodium hydroxide

McAllister grinned, one of the

enemy’s strongest bases had finally been neutralized

neutralization an acid base rxn

Hydrochloric acid + sodium hydroxide ?

ACID + BASE SALT + WATER

HCl + NaOH NaCl + HOH

Neutralization: An Acid Base Rxn

Hydroxide ion

Complete neutralization occurs when

The acid H+ and the base OH- are 1:1

complete neutralization

Net ionic equation:

H+ + OH- HOH

Complete neutralization . . .
  • Occurs when H+ and OH- are added together in a 1:1 molar ratio
  • MaVa = MbVb
  • BE CAREFUL
    • If the acid or base has a subscript you have to account for it in the formula.
    • saMaVa = sbMbVb
    • Ex. If you have 1L of 1M HCL, what M is necessary to neutralize 1L of Mg(OH)2?
titrations
Titrations
  • Lab procedure used to determine a concentration (M) of a solution.
  • Titrations use an indicator to signal when the “endpoint” is reached
    • Endpoint is when the solution is neutralized
    • Indicator changes color
titration neutralization practice
Titration / Neutralization practice
  • 55mL of an 8M HCl solution is completely neutralized by 26mL of Ca(OH)2
    • Write the balanced neutralization reaction.
    • What is the concentration of base that was used?
    • What is the net ionic equation?
    • What are the spectator ions?
salt hydrolysis reaction of a salt with water
Salt Hydrolysis Reaction of a salt with water
  • Salts are ionic compounds
    • contain a + cation and a – anion
    • Derived from acids and bases
  • So when dissolved in water some salts produce:
    • A neutral solution
    • A basic solution
    • An acidic solution
salts that hydrolize water produce alkaline solns

Cation of NaOH

Anion of CH3COOH

Salts that hydrolize water & produce alkaline solns
  • Consists of a cation of a strong base and a anion of a weak acid

NaCH3COO

slide13

NaCH3COO  Na+ + CH3COO-

  • When this salt dissolves in water dissociation occurs.
  • Water also ionizes

H2O  H+ + OH-

  • There is a force of attraction between the
  • H+ of the water and the acetate anion
  • & a force of attraction between the Na+ and
  • OH-
slide14
YES I KNOW . . .
    • Acetic acid CH3COOH is formed
      • BUT it has a low Ka value and doesn’t ionize completely
      • Most of the H+ stays bonded
    • NaOH is formed BUT it stays completely ionized as Na+ and OH-
    • If there are more OH- than H+ in solution the solution is considered BASIC!!!!!!!!!
some salts hydrolyze water produce acidic solns

Weak base

bonds

Strong acid

stays ionized

Some salts hydrolyze water & produce acidic solns
  • The salt of a strong acid and a weak base

NH4Cl dissolved in water produces:

NH4+ Cl- H+ and OH-

slide16
Salts of a weak base and a weak acid may hydrolyze water, but pH needs to be tested.
  • Salts of strong bases and strong acids do not hydrolyzewater and form neutral solutions
    • NaCl
buffers
Buffers
  • A buffer is a mixture of chemicals that make a solution resist a change of pH
    • pH remains relatively constant when adding an acid or base
  • A buffer is either a solution of a weak acid and one of its salts or a weak base and one of its salts
    • The salt cation is only a spectator ion and is not involved in the reaction
buffer capacity
Buffer Capacity
  • Buffers can’t stop the change in pH they can only resist it
    • There comes a point when the buffer is “used up”
    • It can no longer take H+ ions out of solution and the pH begins to change
slide19
This is important when the system is fragile
    • Bloodstream pH 7.3 > 7.5
    • If < 6.9 or > 7.7 the person dies
  • Buffer aids in maintaining homeostasis of the individual
  • The buffer species in the blood is

carbonic acid / hydrogen carbonate ion

H2CO3 / HCO3-

slide20
If the blood is too alkaline, the reaction decreases the amount of OH- in the bloodstream

H2CO3 + OH- HOH + HCO3-

  • If excess H+ is in the blood, the reaction
  • decreases it
  • H+ + HCO3- H2CO3
practice
Practice
  • Write an equation to show the addition of an acid and a base to the following buffer systems

NH4+ / NH3

CH3COOH / CH3COO-

H2PO4- / HPO4-