Acids

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Acids Bases

• PROPERTIES STUFF

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Properties of Acids Bases

• There are two types of specialized solutions
  (acidic and basic)
• Acids have common properties
• taste sour, are corrosive to metals, change
  litmus (a dye extracted from lichens) red, and
  become less acidic when mixed with bases
• Bases have common properties
• feel slippery, change litmus blue, and become
  less basic when mixed with acids

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Names Formulas of Acids

• An acid is a solute that ionizes when aqueous,
  producing H ions
• Therefore the chem formulas of acids are of the
  general form HX
• Where X is a monatomic or polyatomic ion

• When the compnd HCl (g) dissolves in water to
  form HCl (aq), it is named as an acid.

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Names Formulas of Acids
5
Strong and Weak Acids Bases

• Acids can be classified as weak or strong
  depending on the degree to which they ionize in
  water
• Strong acids completely ionize in aqueous
  solns
• Weak acids only partially ionize in aqueous
  solns

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Strong and Weak Acids Bases
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Strong and Weak Acids Bases
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List of Strong Acids/Bases

• Hydrochloric acid (HCl)
• Hydrobromic acid (HBr)
• Hydroiodic acid (HI)
• Nitric acid (HNO3)
• Sulfuric acid (H2SO4)
• Perchloric acid (HClO4)

acids

• Any hydroxide base that includes a metal from
  Group I and Group II

bases
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Names Formulas of Bases

• A base is a solute that when aqueous produces
  OH-1 in water
• We use the ionic compound naming rules to name a
  base
• The name of the cation followed by the name of
  the anion
• NaOH ? sodium hydroxide

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Strong and Weak Acids Bases

• There are also strong and weak bases
• Strong bases dissociate com-pletely into metal
  cations hydroxide ions in soln
• All soluble hydroxides are strong
• Weak bases react with water to form the OH-
  the conjugate acid of the base
• Ammonia (NH3) is a weak base

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Properties of Acids/Bases
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Hydrogen Ions and Acidity

• Weve learned that water is a collection of polar
  molecules in constant motion connected by
  hydrogen bonds
• Collision theory indicates that
• Occasionally, the collisions between water
  molecules are energetic enough to transfer a
  hydrogen from one water molecule to another
• A water molecule that loses a
  hydrogen becomes a OH- ion

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Hydrogen Ions and Acidity

• A water molecule that gains a hydrogen becomes a
  positively charged ion, AKA hydronium, H3O
• This reaction is called the self-ionization of
  water
• Establishes an equilibrium

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Hydrogen Ions and Acidity

• The Hydrogen ions in aqueous solution have
  several aliases.
• Protons, Hydrogen ions, Hydronium ions,
  Solvated protons
• And can symbolized by
• H3O and/or H
• In pure water, this self-ionization occurs to a
  very small extent
• H3O1.0x10-7
• OH-1.0x10-7

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Hydrogen Ions and Acidity

• Notice the concs of the two components are equal
• This described as a neutral soln
• For aqueous solns, the product of H OH-
  equals 1.0x10-14
• H3OOH-1 1.0x10-14
• The product of the conc of the ions will always
  equal 1.0x10-14
• called the ion-product constant for water (Kw)

Kw
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Hydrogen Ions and Acidity

• Therefore the ions are interdependent
• when H3O increases then OH- decreases
• If additional ions of either component are added
  the equilibrium shifts to compensate

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Hydrogen Ions and Acidity

• Of course not all solns are neutral
• When a substance dissolves in water, which
  contributes H, the H increases, so it
  produces an acidic soln (H gt OH-)
• When a substance dissolves in water and
  contributes into OH-, the OH- increases, so
  it produces a basic soln (OH- gt H)

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Hydrogen Ions and Acidity
19
Hydrogen Ions and Acidity
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The pH Concept

• Using conc to express the hydrogen ion content is
  difficult
• A more widely used system is the pH
  (Potential Hydrogen) scale
• A logarithmic scale (log base 10)
• Ranges from 0 to 14 (can be lt0 and gt 14
• HOH- corresponds
  to 7 on the pH scale (neutral)
• pH of 0 is considered highly acidic
• pH of 14 is considered highly basic

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The pH Scale
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The pH Concept
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The pH Concept

• Calculating the pH of a solution is
  straightforward
• pH -logH3O
• You can also calculate pOH, but it isnt
  used as often
• pOH -logOH-1
• pOH would be a scale to decide how basic a
  substance is
• If you know the pH you can calculate pOH
  automatically
• pH pOH 14

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The pH Concept

• Remembering the equilibrium of H3O and OH-
  and that the sum of the pH and pOH always
  equals 14
• Using some simple math, we can bounce back and
  forth between pH, pOH, H, OH-

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pOH14 - pH
pH
pOH
pH14 - pOH
H10(-pH)
OH-10-pOH
pH-logH
pOH-logOH-
OH- 1x10-14/H
OH-
H
H 1x10-14/OH-
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A pH Example

• Given the following info
• fill in the missing pieces.

H pH OH- pOH A or B
12
.0025
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pH Practice

• A student is making a solution of calcium
  hydroxide. She mixes 7.55 grams of calcium
  hydroxide into 500 ml of water.
• What is the concentration of the solution?
• What is the OH-?
• What is the pOH of the solution?
• What is the pH of the solution?
• What is the H3O?
• If she then dilutes the solution to ¼ of its
  original concentration what is the OH-?
• What is the new pOH pH of the solution?

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The pH Concept

• People need to be able to measure the pH of the
  solns they use
• maintaining the correct acid-base balance in a
  pool
• Creating soil conditions ideal for plant growth
• Making medical diagnoses
• Indicators or pH meters are often used in
  measuring pHs
• An indicator is an acid or base that changes
  color in a known pH range

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Acid-Base Indicators
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Acid-Base Indicators

• Knowing the range over which the color change
  occurs gives a rough estimate of pH
• Although indicators are useful tools, they are
  limited
• Some are dependent on temp
• If the soln being tested isnt colorless, the
  indicator may not show up well
• Dissolved salts in a soln may affect the
  dissociation of the indicator

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Acid-Base Indicators
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Acid-Base Indicators

• A pH meter usually gives a more
  accurate, more precise measurement of pH
• The color and cloudiness of the unknown solution
  arent an issue
• Meters are used in hospitals, sewage plants,
  industry, etc.
• When pHs of .01
  matter

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Using Indicators
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Theories Arrhenius Acids Bases

• Chemists had recognized the properties of acids
  bases
• but they were not able to explain the chemical
  theory of this behavior
• Svante Arrhenius proposed a new way of thinking
  about acids bases
• Acids produce H in soln
• Bases produce OH- in soln

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Theories Arrhenius Acids Bases

• The table on the next slide lists some common
  acids
• An acid that contains one ionizable hydrogen is
  called a monoprotic acid
• An acid that contains two ionizable hydrogens is
  called a diprotic acids
• Three ionizable H are called triprotic acids

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Theories Arrhenius Acids Bases
SOME COMMON ACIDS SOME COMMON ACIDS
Name Formula
Hydrochloric acid HCl
Nitric Acid HNO3
Sulfuric Acid H2SO4
Phosphoric Acid H3PO4
acetic Acid CH3COOH
Carbonic Acid H2CO3
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Theories Arrhenius Acids Bases

• Not every hydrogen is created equal
• Only those Hs attached to highly electronegative
  atoms are acidic

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Theories Arrhenius Acids Bases

• Arrhenius bases are soluble hydroxides
• The most common in sodium hydroxide
• Sodium reacts with water to produce sodium
  hydroxide
• Extremely caustic, commonly known as lye, a major
  component of products used to clean clogged drains

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Arrhenius Acids/Bases
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Theories Bronsted-Lowry A B

• The Arrhenius definition of acids and bases is
  not a very comprehensive one
• Its too narrow a definition, doesnt include
  all substances with acidic or basic properties
• For example solutions of NH3 and NaCO3 are
  basic, but neither contain the hydroxide ion

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Theories Bronsted-Lowry A B

• 2 chemists in 1923, independently proposed an
  alternative theory (Bronsted Lowry Theory)
• An acid is a hydrogen ion donor
• A base is a hydrogen ion acceptor
• This new theory allows for ammonias basic
  character, and other discrepencies
• When ammonia is dissolved in water it accepts a
  hydrogen ion from the water

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Bronsted-Lowry Acids/Bases
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Theories Bronsted-Lowry A B

• The acceptor (NH3) is labeled a
  Bronsted-Lowry base
• The donor (H2O) is labeled a Bronsted-Lowry
  acid
• H are transferred from H2O to NH3
• Causes the OH- conc to be greater than it is in
  pure H2O
• therefore, ammonia solns are basic

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Theories Bronsted-Lowry A B

• The ammonia interaction with water is in
  equilibrium
• The NH4 will donate its acquired H to the OH-
  to give NH3 H2O
• In the reverse direction NH4 acts as a B-L acid
  the OH- acts as a B-L base
• The reverse directions compon-ents become
  conjugates of the parent acids and bases
• A conjugate acid is the particle that results
  from a base accepting a H

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Theories Bronsted-Lowry A B

• A conjugate base is the particle that results
  from an acid donating its H
• 2 components related by the loss or gain of a
  single H are called conjugate acid/base pairs
• The NH3 molecule and the NH4 ion are a conjugate
  acid/base pair
• The H2O molecule and OH- ion are also a conjugate
  acid/base pair

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Theories Bronsted-Lowry A B
Acid - Base Pairs
NH3 H2O ? NH4 OH-
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Theories Bronsted-Lowry A B
Acid - Base Pairs
HCl H2O ? Cl- H3O
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Theories Bronsted-Lowry A B

• Did you notice that when water was interacting
  with NH3 it was acting as an acid, and that when
  it interacted with HCl it acted as a base?
• A substance that can act as both an acid and a
  base is amphoteric
• if an acid is present H2O acts as a base/if a
  base is present H2O acts as an acid

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Conjugate Acid-Base Pairs
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Theories Lewis acid/base

• In the 3rd definition of an acid/base Gilbert
  Lewis defined an acid as an e- acceptor, while
  a base is an e- donor
• A Lewis acid has too few e-s
• A Lewis base has too many e-s

Lewis acid
Lewis base
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Theories Acid-Base Definitions
DEFINITIONS OF ACIDS BASES DEFINITIONS OF ACIDS BASES DEFINITIONS OF ACIDS BASES
TYPE ACID BASE
Arrhenius H producer OH- producer
Bronsted-Lowry H donor H acceptor
Lewis electron-pair acceptor electron-pair donor
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Definitions Practice

1. Identify (by formula) the conjugate bases of the
   following acids

NH41
H2O
HSO4-1
HCl
HPO4-2

1. Identify (by formula) the conjugate acids of the
   following bases

H2O
S-2
CO3-2
NO2-1
F-1

1. Identify the acids in the following rxn. Next
   Identify the bases.

HNO2 H2O ? H3O NO2-

1. Identify the Lewis acid and base in the following
   rxn

NO NO3- ? N2O4
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Acid/Base Equilibria

• The equil expression for the diss-ociation of a
  weak acid would be
• Since weak acids only partially ionize, the
  reverse direction is favored

• For dilute solns, the conc of water is a constant

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Acid/Base Equilibria

• So we can pull it out of the equilibrium
  expression and it becomes an acid dissociation
  constant expression (Ka)

• The Ka is the ratio of the conc of the
  dissociated or ionized form of an acid to the
  conc of the nonionized form

Only 1 H at a time
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Acid/Base Equilibria

• It reflects how much acid ion is produced at
  equilibrium
• The stronger the acid the larger the Ka, because
  the conc of the H3O is large and in the numerator

ACID IONIZATION Ka
Oxalic (diprotic) HOOCCOOH ? H HOOCCOO- HOOCCOO- ? H OOCCOO2- 5.6x10-2 5.1x10-5
Phosphoric (triprotic) H3PO4 ? H H2PO4- H2PO4- ? H HPO42- HPO42- ? H PO43- 7.5x10-3 6.2x10-8 4.8x10-13
Acetic CH3COOH ? H CH3COO- 1.8x10-5
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Acid/Base Equilibria

• The equilibrium expression of the eqn of a basic
  solution can be written as well

Base dissociation constant (Kb)
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Acid/Base Equilibria

• The Kb is the ratio of the conc of the conj acid
  times the conc of the OH- to the conc of the conj
  base
• The smaller the value of Kb, the weaker
  the base

Note concentrated dilute indicate how much
(in moles) of an acid or base is dissolved in
soln. The terms strong or weak refers to extent
of ionization.

• Kb and Ka are related through Kw

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Acid/Base Equilibria Example
A 0.1000M solution of acetic acid is only
partially ionized. From measurements of pH of
the solution, H is determined to be
1.34x10-3M. What is the acid dissociation
constant (Ka) of acetic acid?
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Ka/Kb Practice

1. Acetic acid (vinegar) is a weak acid. If
   the HC2H3O2 0.200 M and H 0.0019 M,
   calculate Ka and the Kb for the dissociation of
   acetic acid.
2. Timethylamine (TMA) is a weak base. If the
   initial concentration of TMA is 0.390M and the
   equilibrium conc of OH- is 4.4 x 10-3 M,
   calculate the Kb and the Ka for the dissociation
   of TMA.
3. Chlorous acid, HClO2, has a Ka of 1.0 x
   10-2. Calculate the pH of a 0.10 M solution of
   chlorous acid.

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Titration
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AcidBase Titration

• The conc of an acid especially a weak one or a
  weak base in water is difficult to measure
  directly.
• But we can calculate the conc from the results of
  titration.
• A titration is a carefully controlled
  neutralization rxn.
• Titration is a lab technique for measuring the
  conc of an unknown acid or base

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AcidBase Titration

• The average titration involves three components
• An acid or base of unknown concentration
• You need a
  standard soln
• A standard is
  a solution of
  known conc
• Should be
  opposite of the
  unknown
• An indicator

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AcidBase Titration

• The purpose of the indicator is used to indicate
  when the acid and base ions are equal
• Lets the person performing the titration know
  when neutralization has occurred
• Phenolphthalein is a very common indicator chosen
• It undergoes a color change at a pH of about 7.6
• clear in acid
• Light pink in neutral
• Dark pink in base

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AcidBase Titration

• In titration, the stand-ard is slowly added to
  the unknown soln
• As the 2 solns mix, the acid in one neutralizes
  the base in the other
• Eventually, enough standard soln is added to
  neutralize all of the acid or base in the unknown
  soln.
• Indicator changes color

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Procedure for Titration

• The point at which this occurs is called the
  equivalence point.
• We can calculate the conc from the results of
  titration
• _at_ the equivalence point the H30 OH-
• The point at which the indicator changes color is
  called the end point of the titration.
• If the indicator is chosen correctly the end
  point should be very close to the equivalence
  point.

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Procedure for Titration

• Therefore, at approximately the end point
  of a titration the total mols of H donated by
  the acid is equal to the total mols of H
  accepted by the base.
• We call these molar equivalences (stoichiometric
  equivalences)
• Calculate moles of standard soln
• Calculate moles of unknown (stoichiometry)
• Claculate conc of unknown

Remember (conc)(vol in L) moles
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Sample Problem

• Solutions of NaOH are used to
  unclog drains.
• A 43.0 ml volume of NaOH was titrated with 32.0
  ml of .100M HCl.
• What is the molar concentration of the
  NaOH solution?

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Sample Problem

• 200 ml of a 0.1 M solution of sodium
  hydroxide was needed to neutralize 125 ml of a
  sulfuric acid solution. What conc was the
  sulfuric acid soln?

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Titration Practice

1. Calculate the conc of a nitric acid solution if a
   20 ml sample of the acid required an average vol
   of 55 ml of a 0.047 M soln of Ba(OH)2 to reach
   the endpoint of the titration.
2. 15 ml of 12 M Hydrochloric acid is diluted to 250
   ml. A 10.0 ml aliquot is measured and titrated
   with an unknown conc of Potassium hydroxide (vol
   of 20ml). What is the conc of the hydrochloric
   acid? What is the conc of the Potassium
   hydroxide solution?

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Titration Curves

• We can also track the titration using a pH meter
  or indicator paper to record the pHs as we add
  excess acid or base.
• If we graph the results, we get a
  characteristic curve called a titration curve
• Graphing pH vs. volume of titrant
• Using a titration curve we can determine the
  equivalence point by analyzing the vertical of
  the titration curve.

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Strong Acid/Base Titration
pKapH
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Strong Acid/Base Titration
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Strong Acid/Base Titration
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Weak Acid w/Strong Base
Buffer Region
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Strong Acid/Weak Base Titration
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Polyprotic acid titration
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Polyprotic Acid Titration
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Buffers

• If you add 10ml of 0.1M NaOH to 1 L of pure H2O
  you would increase the pH from 7.0 to 11.0
• If you dissolve a mixture of sodium acetate and
  acetic acid into water and then add the same
  amount of base, the pH will only increase by 0.01
  
• The mixture acts as a buffer, which is a soln in
  which pH remains relatively constant with the
  addition of small amounts of acid or base.

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Buffers

• A buffer soln consists of a WA or a WB and one
  of its salts
• The acetic acid (CH3COOH) and its anion (CH3COO-)
  act as reservoirs of neutralizers
• They react with any OH- or H added to the soln
• A buffer doesnt have unlimited neutralizing
  power
• The amnt of A or B that can be added to a buffer
  soln before a sig. ?pH occurs is the buffer
  capacity

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Buffers
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Buffers
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Buffers

• Buffers are used to maintain the pH within an
  acceptable range.
• Human blood pH is 7.4, and a ?pH of 0.2 units in
  either direction is considered dangerous
• Acidosislt7.4ltalkalosis
• The body must get rid of about 15 mols of
  potential acid per day
• It uses a mixture of buffers like bicarbonate,
  the lungs reduce CO2 (considered acidic), and the
  kidneys filter out acids and bases

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Buffers
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