PH

The pH of a solution is a measure of its acidity or alkalinity. The pH scale ranges from 0 to 14. A

change of one unit on the pH scale represents a change in the concentration of hydrogen ions by a factor

of 10, a change in two units represents a change in the concentration of hydrogen ions by a factor of 100.

Thus, small changes in pH represent large changes in the concentrations of hydrogen ions. Pure water

is neutral. It is neither acidic nor basic and has a pH of 7.0. Anything below 7.0 (ranging from 0.0 to

6.9) is acidic, and anything above 7.0 (from 7.1 to 14.0) is alkaline. The blood in your veins is slightly

alkaline (pH = 7.4). The environment in your stomach is highly acidic (pH = 1 to 2). Orange juice is

mildly acidic (pH = approximately 3.5), whereas baking soda is basic (pH = 9.0).

Acids are substances that provide hydrogen ions (H+) and lower pH, whereas bases provide

hydroxide ions (OH–) and raise pH. The stronger the acid, the more readily it donates H+. For example,

hydrochloric acid and lemon juice are very acidic and readily give up H+ when added to water.

Conversely, bases are those substances that readily donate OH–. The OH– ions combine with H+ to

produce water, which raises a substance’s pH. Sodium hydroxide and many household cleaners are very

alkaline and give up OH– rapidly when placed in water, thereby raising the pH.

40 Matthew R. Fisher

How is it that we can ingest or inhale acidic or

basic substances and not die? Buffers are the key.

Buffers readily absorb excess H+ or OH–,

keeping the pH of the body carefully maintained

in the aforementioned narrow range. Carbon

dioxide is part of a prominent buffer system in the

human body; it keeps the pH within the proper

range. This buffer system involves carbonic acid

(H2CO3) and bicarbonate (HCO3–) anion. If too

much H+ enters the body, bicarbonate will

combine with the H+ to create carbonic acid and

limit the decrease in pH. Likewise, if too much

OH– is introduced into the system, carbonic acid

will combine with it to create bicarbonate and

limit the increase in pH. While carbonic acid is an

important product in this reaction, its presence is

fleeting because the carbonic acid is released

from the body as carbon dioxide gas each time we

breathe. Without this buffer system, the pH in our

bodies would fluctuate too much and we would

fail to survive. What is pH? All chemicals can be grouped by their properties. One important property of many chemicals is pH in solutions with water, which indicates whether a substance, when dissolved in H2O forms an acidic solution, a basic solution, or a neutral solution. What is pH? The p stands for potential, or power. H is the chemical symbol for the element hydrogen and represents the hydrogen ion (H+). So pH is the power (or concentration) of hydrogen ions in a solution. Basically, if you have a lot of H+ you are an acid. If you have a little you are a base. The pH of a solution indicates how acidic or basic (alkaline) the solution is. What causes pH? Pure water breaks apart to a small extent, by dissociation, into its constituent ions: H+ is the hydrogen ion and OH- is the hydroxide ion. In pure water, the concentration of H+ is 1x10 -7 and the pH (- log [H+]) = 7. Sometimes chemicals dissociate into ions that produce a LOT of H+ or OH-. These chemicals we call acids and bases. Acids, like HCL, dissociate (break up) into H+ and Cl-. This increases the H+ concentration, making HCL “hydrochloric acid”. Bases, like NaOH, dissociate into Na+ and OH-. This increases the OH- concentration, making sodium hydroxide, into a strong base. Generally, bases are not indicated in the name. The hydroxide ion (OH-) should provide the clue. Acids and Bases Acids are a group of chemicals that release H+ when they dissolve in H2O increasing [H+]. When a food tastes sour, it usually contains acid. Bases are chemicals that remove H+ when they dissolve in H2O decreasing in [H+]. Bases typically have no characteristic taste but they tend to feel slippery or soapy. How acidic or basic a solution is has huge implications in biological systems. If a protein (our biological machines) are exposed to a pH that is out of a certain range, the extra H+ or OH- ions will disrupt the structure of the protein, causing it to malfunction. We call this "denaturing" the protein. Heat and extreme pH can both cause proteins to denature and we see this in cooking all of the time. When you put an egg on a hot skillet you denature the proteins and turn the egg protein into breakfast. When you place raw fish in a bowl of fresh squeezed lime juice (citric acid) you make ceviche. In almost all cases once a protein is "denatured" it is broken forever! Buffers To keep the shape (the structure) of important proteins intact, biological systems use buffers to maintain the pH around a certain level. A buffer is a chemical that resists pH change- meaning if a solution is buffered at pH 10, it will not be changed easily. A buffer solution can be anywhere on the pH scale, the important thing to remember is that it keeps pH the same despite added acid or base. We have buffers in our stomach keeping our stomach pH low, even if you eat foods that are alkaline (basic). Practice your application! If pH is a measure of H+ in a solution, what would happen if you add H2SO4? answer: there will be more of the H+ in solution, lowering the pH! (I bolded the H) If you add NaCl to pure water, what should happen to the pH? answer: ??? give it your best shot, then google it to check your answer or email me to confirm! Here are some values you must memorize to do well on the quizzes pH is measured on a scale from 0-14. pH levels from 0 - < 7.0 are acidic. pH levels from > 7.0 - 14 are basic. A pH of exactly 7.0 is neutral.

Why is pH important? Many chemical reactions in water (aqueous solutions, e.g., lakes, oceans, the atmosphere, your blood and other tissues) are controlled (rate, extent, etc.) by the pH of the solution. Many molecules are altered by exposure to positively charged hydrogen ions or negatively charged hydroxide ions (Importantly, proteins, enzymes, pH paper) Human fluid systems (blood, interstitial fluid, CSF, urine, etc.) require very narrow pH ranges to function properly. Human digestion requires very low pH in the stomach and a higher, slightly alkaline, pH in the intestines. Organisms are very sensitive to changes in acidity and alkalinity. For example, some bacteria can only grow in acidic solutions, while some marine organisms can only survive in slightly alkaline (basic) environments. Our bodies are very sensitive to pH change as well. If there is an disruption in pH for an extended period, we can easily die. The reason for such sensitivity is explained in part by the characteristics of proteins – especially enzymes. Any change in temperature or pH that goes beyond an appropriate level to maintain normal enzyme functions causes the enzyme to distort its shape (denature). Change of enzyme shape means loss of reactivity, since enzyme function is shape-specific. Enzymes work like locks and keys. If there is any change in shape to the key or the lock, the interaction will not work. pH can directly alter these shapes. Due to the fact our bodies rely on enzymes (type of protein) for just about all of our biochemistry, we also rely on the perfect pH to conduct that chemistry. To prevent massive swings in pH, living systems possess chemical compounds that act as pH buffers. They react with hydroxide (OH-) and hydrogen (H+) ions to prevent the pH from changing too much or too quickly. This helps protect the organism from damage and death. pH buffers are usually proteins, but many others exist. I think about buffers like the bumpers on your car. They can protect your car from small accidents, but not a massive crash. Generally, buffers can cushion small pH swings, but can be overwhelmed in there is a massive amount of H+ or OH-.

To ace your closed-note pH lab quiz, focus on these key concepts and memorization points directly from your notes:

1. Core Definitions & Values to Memorize

• pH Scale: Ranges from $0$ to $14$.
  • Acidic: $ext{pH} < 7.0$ (high concentration of $H^+$).
  • Neutral: $ext{pH} = 7.0$ (pure water, where $[H^+] = 1 imes 10^{-7}$).
  • Basic / Alkaline: $ext{pH} > 7.0$ (high concentration of $OH^-$).
• Logarithmic Scale Factor: A change of $1$ pH unit represents a $10 imes$ change in $[H^+]$ concentration. A change of $2$ units represents a $100 imes$ change.

1. Biological Examples & Specific pH Values

• Stomach environment: Highly acidic ($ext{pH} = 1$ to $2$).
• Orange juice: Mildly acidic ($ext{pH} imes 3.5$).
• Blood: Slightly alkaline ($ext{pH} = 7.4$).
• Baking soda: Basic ($ext{pH} = 9.0$).

1. Acids, Bases, and Dissociation Mechanisms

• Acids: Release $H^+$ ions when dissolved in water, lowering pH (e.g., HCl
  ightarrow H^+ + Cl^-).
• Bases: Release $OH^-$ ions or remove $H^+$, raising pH (e.g., NaOH
  ightarrow Na^+ + OH^-). $OH^-$ combines with $H^+$ to form $H_2O$.

1. Buffers and Biological Importance

• Function: Buffers resist pH change by absorbing excess $H^+$ or $OH^-$ ions to maintain narrow required pH ranges.
• Human Body Buffer: The bicarbonate ($HCO<em>3^-$) and carbonic acid ($H</em>2CO<em>3$) system maintains blood pH. Excess $H^+$ combines with $HCO</em>3^-$ to form $H<em>2CO</em>3$, which is eventually exhaled as $CO_2$ gas.
• Denaturation: Extreme pH or heat distorts the shape of proteins and enzymes (lock-and-key model), causing irreversible loss of function.

1. Application Practice Questions

• Adding $H<em>2SO</em>4$: Increases $H^+$ concentration, lowering the pH.
• Adding $NaCl$ to pure water: $NaCl$ dissociates into $Na^+$ and $Cl^-$ ions, providing neither $H^+$ nor $OH^-$ ions. Therefore, the pH remains neutral ($7.0$).