MODULE 1 LESSON 1- pH of Common Solutions
Measurement of pH (potentiometric hydrogen ion concentration)
➔ measures the acidity level of a solution (chemical in an aqueous/water solution)
➔ any changes to the pH level can disrupt living systems
➔ rated on a 14-point scale
◆ Pure water- neutral pH close to 7.0 at 25 degrees Celsius
◆ 7: neutral
◆ >7: acidic (0-6)
◆ <7: basic/alkaline (8-14)
➔ pH is measured using pH paper or a pH meter.
➔ Not All Liquids Have a pH Value, pH only has meaning in an aqueous
solution (in water).
◆ vegetable oil, gasoline, pure alcohol
pH of Common Acids
pH
Common Acidic Substances
0
Hydrochloric Acid (HCl)
1.0
Battery Acid (H2SO4 sulfuric acid) and stomach acid
2.0
Lemon Juice
2.2
Vinegar
3.0
Apples, Soda
3.0 to 3.5
Sauerkraut
3.5 to 3.9
Pickles
4.0
Wine and Beer
4.5
Tomatoes
4.5 to 5.2
Bananas
around 5.0
Acid Rain
5.0
Black Coffee
5.5ish
Human skin, hair, and nails
6.0
urine, saliva
* Milk is considered neutral (since it’s slightly acidic) but becomes more acidic overtime.
Neutral pH chemicals
7.0
Pure water
* Distilled water tends to be slightly acidic due to dissolved CO2 & other gasses.
* Blood is close to neutral but is slightly basic.
pH of Common Bases
pH
Common Basic Substances
7.0 to 10
Shampoo
7.4
Human Blood
7.4
Human Tears
7.8
Egg
around 8
Seawater
8.3
Baking Soda (Sodium Bicarbonate)
around 9
Toothpaste
10.5
Milk of Magnesia
11.0
Ammonia
11.5 to 14
Hair Straightening Chemicals
12.4
Lime (Calcium Hydroxide)
13.0
Lye
14.0
Sodium Hydroxide (NaOH)
Other pH Values
Soil - pH ranges from 3 to 10
Plants - pH between 5.5 and 7.5
Buffer Solution - prepared to maintain a pH near 7
*dissolving NaCl (salt) in water doesn’t change its pH value.
How to measure pH?
❖ pH paper strips
➢ color of the test strips corresponds to a pH range
➢ color change depends on the type of indicator dye used to coat the
paper, the result needs to be compared against a chart of standard.
➢
❖ pH indicator
➢ Drops of pH indicator on a small sample of a substance, then observe
test change.
➢
❖ pH test kits
➢ available to test liquids
➢ designed for a particular application, like aquaria or swimming pools
➢ fairly accurate but may be affected by other chemicals in a sample.
➢
❖ pH meter
➢ most accurate method of measuring pH
➢ Measures [H+] by means of electrodes immersed in the test solution
➢ more expensive than test papers or kits
➢ require calibration
➢ generally used in schools and labs.
➢
pH Equation
➔ Proposed in 1909 by Danish biochemist Søren Peter Lauritz Sørensen
➔ pH = -log[H+]
◆ log is the base-10 logarithm
◆ [H+] stands for the hydrogen ion concentration in units of moles per liter
solution.
➔ "pH" abbreviation for power of Hydrogen, comes from the German word
"potenz," (power) combined with H, the element symbol for hydrogen
* chemicals that have very low or very high pH are often corrosive and can produce
chemical burns.
* better dilute them first in pure water to test their pH (value unchanged, risk reduced)
Water [H+]: 1x10-7 moles/L
● Acid [H+]: >1x10-7 moles/L
● Bases [H+]: <1x10-7 moles/L
MODULE 1 LESSON 2- BUFFER PREPARATION
Bicarbonate buffering system - keeps blood’s pH in check
Buffer Solutions
➔ An aqueous solution mixture of a weak acid and its conjugate base/weak base
and its conjugate acid
◆ These two components maintain a pH balance that resists change when
strong acids or bases are added to it.
➔ Conjugate base - formed by removing a proton (minus 1 H+) from the same
type of acid/by combining a weak base with its conjugate acid
◆ use of conjugates is what gives a buffer solution its resistance to pH
changes
◆ It creates an equilibrium between the acids and the bases
➔ Examples:
◆ Blood ( contains a bicarbonate buffering system)
◆ TRIS buffer
◆ Phosphate buffer
BUFFER
pKa
pH range
citric acid
3.13., 4.76, 6.40
2.1 to 7.4
acetic acid
4.8
3.8 to 5.8
KH2PO4 (Potassium dihydrogen
7.2
6.2 to 8.2
borate
9.24
8.25 to 10.25
CHES
9.3
8.3 to 10.3
phosphate) or (Monopotassium
Phosphate)
*strong acids (HCl) is added to lower the pH of acidic buffers
*strong bases (NaOH) is added to lower the pH of basic/alkaline buffers
➔ Functions of a buffered pH:
◆ required for most enzymes to function correctly
◆ used to ensure proper color concentration when using dyes
◆ used to calibrate equipment, especially pH meters (might be
miscalibrated if a buffer is not present)
*buffer solutions do not necessarily have a neutral pH, just a balanced one
➔ Buffer Capacity
◆ The amount of a strong acid or base that a buffer solution can take before
significant pH changes occur
◆ differs depending on the core components of the buffer solution and how
much of the strong acid or base is added to the solution
● Adding strong acid = amount of base in the buffer solution
● Adding strong base = amount of acid in buffer solution
MODULE 1 LESSON 3- BUFFER SYSTEMS
Buffer system
➔ important since most enzymes require a very specific pH ranges in order to work
properly
◆ Human blood
● average pH of human blood is 7.4.
● If the pH goes below 6.8 or above 7.8 death may result.
● To hold the pH of the blood close to 7.4, the body uses three buffer
systems: carbonates, phosphates, and proteins.
➔ 3 KINDS OF BUFFER SYSTEMS
◆ Carbonates
● Most important
● made up of Carbonic acid – Bicarbonate ion in a ratio of 10:1
● Normal concentration in blood : 0.025 M HCO3- and 0.0025 M
H2CO3
NaHCO3 + HCl → H2CO3+NaCl
(sodium bicarbonate) + (strong acid) → (weak acid) + (salt)
H2CO3 + NaOH→HCO3- + H2O
(weak acid) + (strong base)→(bicarbonate) + (water)
◆ Phosphates
● Second most important
● made up of sodium dihydrogen phosphate (Na2H2PO4−) – sodium
hydrogen phosphate (Na2HPO42-) ion in a ratio of 1.6 : 1 ratio
HCl + Na2HPO4→NaH2PO4 + NaCl
(strong acid) + (weak base) → (weak acid) + (salt)
NaOH + NaH2PO4→Na2HPO4 + H2O
(strong base) + (weak acid) → (weak base) + (water)
◆ Proteins
● Almost all proteins can function as buffers.
● made up of amino acids, which contain positively charged amino
groups and negatively charged carboxyl groups.
● charged regions of these molecules can bind hydrogen and
hydroxyl ions, and thus function as buffers
● accounts for two-thirds of the buffering power of the blood and
most of the buffering within cells.
○ HEMOGLOBIN
◆ principal protein inside red blood cells
◆ ⅓ of the mass of the cell
● Normal arterial blood pH is restricted to a very narrow range of 7.35 to 7.45
○ ACIDOSIS
◆ occurs when there is overproduction of acid in the
body fluids or the blood pH falls below 7.35.
◆ Can either by caused by:
● Metabolic Acidosis
○ excessive loss of bicarbonate from the
blood
● Respiratory Acidosis
○ buildup of carbon dioxide in the blood
that results from poor lung function or
depressed breathing
◆ Symptoms include headache and confusion, being
lethargic, and fatigue.
○ ALKALOSIS
◆ occurs when there is excessive blood alkalinity in the
body fluids or the blood pH rises above 7.45
* a pH above 7.8 is fatal
◆ Can either by caused by:
● Metabolic Alkalosis
○ loss of acid from the blood
● Respiratory Alkalosis
○ low level of carbon dioxide in the blood
that results from rapid or deep breathing
◆ Symptoms include cognitive impairment (which can
progress to unconsciousness), tingling or numbness
in the extremities, muscle twitching and spasm, and
nausea and vomiting.