1604_L5_Chemical reaction_AY25-26_teacher version

General form: A-B + C → AC + B. Clinical relevance: metal ion displacement can affect protein function and is relevant in heavy metal poisoning.

The components of a chemical equation include:

1. Reactants - substances that undergo a chemical change.
2. Products - substances formed as a result of the chemical change.
3. Coefficients - numbers that indicate the relative amounts of reactants and products.
4. Symbols for state of matter - such as (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solutions.

According to the Brønsted-Lowry theory:

• Acids are H+ donors.
• Bases are H+ acceptors.

Acids can be strong (fully dissociate) or weak (partially dissociate) in water.

pH = -log[H3O+]

pH < 7: acidic; pH = 7: neutral; pH > 7: basic.

A buffer resists changes in pH when acids or bases are added. The bicarbonate buffer system:

CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3−

• When H+ is added, HCO3− binds H+ to form H2CO3, minimizing pH change.
• When OH− is added, H2CO3 donates H+ to neutralize OH−.
• The respiratory system controls CO2 (shifts equilibrium), and the kidneys regulate HCO3− to maintain blood pH.

The major types of chemical reactions in the body include:

1. Synthesis (combination) reactions: build complex molecules from simpler ones.
2. Decomposition reactions: break complex molecules into simpler ones.
3. Single replacement reactions: one element replaces another in a compound.
4. Double replacement reactions: ions are exchanged between compounds.
5. Combustion (oxidation) reactions: oxygen-involving reactions that release energy.

Acid-base reactions:

• Maintain pH homeostasis critical for enzyme activity and metabolism.
• Participate directly in metabolic pathways and transport (e.g., proton transfers).
• Are stabilized by buffer systems (bicarbonate, phosphate, protein) to prevent harmful pH shifts.

The pH scale matters because:

• Enzymes have optimal pH ranges; deviations impair activity.
• Many metabolic reactions are pH-dependent.
• Organ systems (blood, cells) require stable pH for proper function.

Reactants: The starting substances in a chemical reaction. Products: The new substances formed as a result of the reaction.

General form: A + B → AB. Clinical relevance: involved in biosynthesis and conditions like hypoxia management and anemia therapy.

Example: ATP → ADP + Pi + Energy. Clinical significance: relates to fatigue and metabolic disorders due to impaired ATP production.

General form: AB + CD → AC + BD. Example: HCl + NaHCO3 → NaCl + H2CO3 (antacid neutralization), important for acid-base balance.

Catabolism: the breakdown of molecules that releases energy. Function: provides ATP for cellular activities (e.g., glucose → ATP + CO2 + H2O).

Anabolism: synthesis of molecules requiring energy. Example: amino acids → proteins for tissue repair and enzyme production.

Oxidation-reduction (redox) involves electron transfer. Role: central to ATP production in mitochondria (e.g., NAD+ accepts electrons to become NADH).

Anabolic reactions:

• Require ATP
• Build larger molecules from smaller ones (e.g., glycogen synthesis)

Catabolic reactions:

• Release energy (ATP and heat)
• Break down larger molecules into smaller ones (e.g., glycogen → glucose)

Redox reactions involve electron transfer between substances. Significance: essential for energy production, detoxification, and many metabolic pathways.

Catabolism: the process that results in net release of free energy (exergonic).

Anabolism describes synthesis of complex proteins from amino acid subunits.

Catabolism is the term for breakdown of lipids into fatty acids and glycerol.

The breakdown of glycogen into glucose in the liver is catabolism (glycogenolysis).

The process of creating new DNA strands during replication is anabolism (biosynthesis).

Brønsted-Lowry definition:

• Acid: H+ donor
• Base: H+ acceptor

HCl donates H+ to H2O, forming H3O+ (hydronium) and Cl−; demonstrates acid donating a proton to a base.

Developed independently by J. N. Brønsted (Denmark) and T. M. Lowry (Great Britain) in 1923; they expanded acid-base definitions beyond OH−-containing bases.

A strong acid completely ionizes (100%) in water, producing high ion concentration; a weak acid only partially dissociates, yielding lower ion concentration.

The strength of an acid is inversely related to the strength of its conjugate base. Strong acids have weak conjugate bases, while weak acids have stronger conjugate bases.

1. Hydroiodic acid (HI) - Conjugate base: I⁻ (Iodide ion)
2. Hydrobromic acid (HBr) - Conjugate base: Br⁻ (Bromide ion)
3. Sulfuric acid (H₂SO₄) - Conjugate base: HSO₄⁻ (Hydrogen sulfate ion)

Weak acids partially dissociate in solution, resulting in a higher concentration of their conjugate bases. In contrast, strong acids completely dissociate, leading to a lower concentration of their conjugate bases.

The conjugate base of acetic acid (HC₂H₃O₂) is acetate ion (C₂H₃O₂⁻).

The conjugate base of the hydronium ion (H₃O⁺) is water (H₂O).

The pH scale indicates the acidity of a solution, with values ranging from 0 to 14. Solutions are considered acidic when pH is less than 7, neutral at a pH of 7, and basic when pH is greater than 7.

• Acidic: pH < 7
• Neutral: pH = 7
• Basic: pH > 7

• Acidic solution: pH < 7.0, [H3O+] > 1.0 × 10^-7 M
• Neutral solution: pH = 7.0, [H3O+] = 1.0 × 10^-7 M
• Basic solution: pH > 7.0, [H3O+] < 1.0 × 10^-7 M

The pH scale is a logarithmic scale that corresponds to the concentration of hydronium ions [H3O+] in aqueous solutions. It is calculated as the negative logarithm (base 10) of the [H3O+]. For example, if [H3O+] is 1.0 × 10^-2 M, then pH = -log[1.0 × 10^-2] = 2.00.

To find the pH of a solution with a given [H3O+] concentration, use the formula: pH = -log[H3O+]. For example, for [H3O+] = 4.0 × 10^-5, the pH would be calculated as follows:

1. Substitute the value into the formula: pH = -log[4.0 × 10^-5].
2. Calculate the logarithm to find the pH value.

In a neutralization reaction, an acid reacts with a base to produce salt and water. The salt formed consists of the anion from the acid and the cation from the base.

Antacids are used to neutralize excess stomach acid.

Some antacids are made of aluminum hydroxide and magnesium hydroxide mixtures.

Antacids are used to neutralize excess stomach acid.

Amphojel contains Aluminum hydroxide (Al(OH)3) as its base.

The antacids that contain magnesium hydroxide (Mg(OH)2) include:

1. Milk of magnesia
2. Mylanta, Maalox, Di-Gel, Gelusil, Riopan
3. Bisodol, Rolaids

The reaction is:

Mg(OH)2(aq) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)

The products of the reaction are:

1. Sodium chloride (NaCl)
2. Carbon dioxide (CO2)
3. Water (H2O)

The reaction is:

NaHCO3(s) + HCl(aq) → NaCl(aq) + CO2(g) + H2O(l)

A buffer solution maintains the pH by resisting changes when acids or bases are added, absorbing H3O+ or OH- to keep the pH stable.

Maintaining a pH close to 7.4 in blood is crucial because a change in blood pH affects the uptake of oxygen and various cellular processes.

Buffers are important for the proper functioning of cells and blood by stabilizing pH levels, which is essential for various biochemical reactions.

When an acid or base is added to water, the pH changes drastically, while a buffer solution only experiences a slight change in pH.

The key reaction is: CO2 + H2O = H2CO3 = H+ + HCO3

• The Respiratory System: Controls the release of carbon dioxide (CO2) by adjusting the breathing rate.
• The Kidneys: Regulate bicarbonate ion (HCO3¯) levels by either reabsorbing them into the blood or excreting them in urine based on the body's needs.

Carbonic acid (H2CO3) is formed from the reaction of carbon dioxide (CO2) and water (H2O) and can dissociate into hydrogen ions (H+) and bicarbonate ions (HCO3-), helping to maintain pH balance in the body.

The bicarbonate buffer system neutralizes excess H⁺ by converting it into carbon dioxide (CO₂), which is then exhaled by the lungs. This process helps to stabilize the pH of the blood.

1. Buffering: H⁺ + HCO₃⁻ → H₂CO₃ (Free H⁺ is neutralized by bicarbonate, forming weak carbonic acid.)

2. Conversion: H₂CO₃ → CO₂ + H₂O (The carbonic acid breaks down into carbon dioxide and water.)

3. Exhalation: CO₂ is exhaled (The lungs remove the CO₂, eliminating the acid load from the body.)

Overall Result: H⁺ + HCO₃⁻ → CO₂ + H₂O (The hydrogen ion is converted and removed, stabilizing pH.)

Acidosis occurs when blood pH is below 7.35, indicating it is too acidic. Alkalosis occurs when blood pH is above 7.45, indicating it is too alkaline.

Hypoventilation leads to an increase in CO2, causing a shift to the right in the equilibrium reaction, resulting in respiratory acidosis. A clinical example includes conditions like COPD and asthma.

Hyperventilation causes a decrease in CO2, leading to a shift to the left in the equilibrium reaction, resulting in respiratory alkalosis. This can occur in situations such as anxiety or fever.

Ketoacidosis results in an increase in H+, causing a shift to the left in the equilibrium reaction, leading to metabolic acidosis. This is commonly seen in conditions like diabetes and sepsis.

Vomiting HCl leads to a decrease in H+, causing a shift to the right in the equilibrium reaction, resulting in metabolic alkalosis. A clinical example is pyloric stenosis.