Lecture 1: Water and Aqueous System PDF

Lecture 1 Water and Aqueous System Dr. Sophie SHI Ling [email protected] School of Science and Technology 1 Content I. Properties of water II. Interactions and bonding III. Ionization of water, acids and bases IV. Buffer sy...

Lecture 1 Water and Aqueous System Dr. Sophie SHI Ling [email protected] School of Science and Technology 1 Content I. Properties of water II. Interactions and bonding III. Ionization of water, acids and bases IV. Buffer system and mechanism of blood pH 2 I. Properties of water II. Interactions and bonding 3  https://www.youtube.com/watch?v=9iMGFqMmUFs 4  Chemical elements of life Ø Most abundant elements in cell: carbon (C), hydrogen (H), nitrogen (N), oxygen (O), phosphorus (P), sulphur (S) Account for over 99% of the atoms in human body Ø H, C, N and O have relatively low atomic numbers electron shell nucleus atomic number H C 1 N 6 7 mass number 1 12 14 O P S 8 15 16 5 16 31 32 Molecular bonds and forces Ø Covalent bond Forms when atoms share one or more pairs of electrons Example: H2 (hydrogen gas) Ø Ionic bond Consists of ions and forms through the electrical force between oppositely charged ions Example: NaCl Ø Hydrogen bond An attraction between two atoms that already participate in other chemical bonds One of the atoms is hydrogen, while the other may be any electronegative atom, such as oxygen, chlorine, or fluorine Example: HF (Hydrogen fluoride), H2O Hydrogen fluoride 6  Molecular bonds and forces Ø Hydrophobic interaction Polar vs Non-polar ‒ Polar molecule: have an uneven distribution of charge across their geometry resulting in one side being positive and the other side negative ‒ Non-polar molecule: molecules that do not have any electrical charges or partial charges Hydrophobic (“water hating”): the interaction between nonpolar molecules or groups in water Example: lipid-water interaction Ø Van Der Waals force Distance-dependent forces between atoms and molecules not associated with covalent or ionic chemical bonds Weak interaction 7 Basic properties of water Ø Structural properties Chemical formula: H2O Consists of one atom of oxygen (O) and two atoms of hydrogen (H) A polar molecule ‒ the oxygen has a large nucleus that attracts electrons (strong bond) - ‒ causing the oxygen atom to be negatively charged ‒ whereas the hydrogens are positively charged + (weak bond) Allows water molecules to interact with each other by hydrogen bonds ‒ attraction between positive H in one H2O molecule to negative O in another H2O ‒ weak bond 8 Important properties of water Ø Cohesion and adhesion Cohesion: water molecules stick to each other Example: drop a tiny amount of water onto a very smooth surface, the water molecules will stick together and form a droplet, rather than spread out over the surface Adhesion: water molecules stick to other molecules or substances ‒ Example: the water droplets are stuck to the end of the pine needles 9  Important properties of water Ø High boiling point Water’s boiling point (100°C) is higher than the boiling points of similar substances due to the hydrogen bonds Ø High specific heat More energy to increase the temperature of water Water plays a very important role in temperature regulation This property helps to maintain homeostasis Ø Solvent for polar molecules The polarity and ability of water to form hydrogen bonds makes it an excellent solvent Most of the chemical reactions important to life take place in a watery environment inside of cells Water's capacity to dissolve a wide variety of molecules is key in allowing these chemical reactions to take place 10 III. Ionization of water, acids and bases 11  Ionization of water ØDue to the polar structure, water can act as either an acid or a base Acid: can donate a proton (H+) to a base Base: can use a lone pair of electrons to accept a proton (H+) ØAutoionization of water 12  Ionization of water Ø Equilibrium constant (Kw) for the autoionization of water: Kw = [H+][OH-] or [H3O+][OH−] [H+]: Concentration of H+ [H3O+]: Concentration of H3O+ [OH−]: Concentration of OH- Ø At 25 ℃, the experimentally determined value of Kw in pure water is 1.0×10−14 Kw = [H+][OH−] = 1.0×10−14 Ø In a sample of pure water, the concentrations of H+ and OH− are equal to one another [H+] = [OH−] = 1.0×10−7 Ø No matter whether the aqueous solution is an acid, a base, or neutral, the Kw remains constant For acidic solutions, [H+] > [OH−] For basic solutions, [OH−] > [H+] For neutral solutions, [H+] = [OH−] = 1.0×10−7M 13  The Concentration of H+ (protons) in solution (pH) Ø pH The measure of the concentration (in molar or moles/liter) of hydrogen ions (H+) in solution Defined as the negative log of [H+], which will give a positive value for pH pH= –log [H+] [H+]: Concentration of H+ The pH of pure water is 7 at 25 °C Classification Relative ion pH at 25℃ conc. acidic [H+]>[OH-] pH < 7 neutral [H+]=[OH-] pH = 7 basic [H+] 7 14 Strong acids and strong bases Ø Strong acids Acids that are completely or nearly 100% ionized in their solutions The ionization equation of a strong acid (HA): HA H+ + A- Ø Strong bases Bases that are completely or nearly 100% ionized in their solutions The ionization equation of a strong base (BOH): BOH OH − + B+ Strong acids Strong bases hydrochloric acid (HCl) sodium hydroxide (NaOH) nitric acid (HNO3) potassium hydroxide (KOH) sulfuric acid (H2SO4) calcium hydroxide (Ca(OH)2) perchloric acid (HClO4) lithium hydroxide (LiOH) hydrobromic acid (HBr) strontium hydroxide (Sr(OH)2) hydroiodic acid (Hl) barium hydroxide (Ba(OH)2) 15 Weak acids and weak bases Ø Weak acids Acids that partially dissociates when it is dissolved in a solvent Because this dissociation does not go 100% to completion, the ionization is considered as a reversible reaction: HA H+ + A- Ø Weak bases Bases that partially dissociates when it is dissolved in a solvent The ionization equation of a weak base (BOH): BOH OH − + B+ Ø Ionization Constant Acid ionization constant: Base ionization constant: pKa and pKb: 16 Weak acids and weak bases Ø Common weak acids and weak bases Weak acids Weak bases Acid name Formula Base Name Formula Formic HCOOH Ammonia NH3 Acetic CH3COOH Trimethyl N(CH3)3 ammonia Trichloroacetic CCl3COOH Pyridine C5H5N Hydrofluoric HF Ammonium NH4OH hydroxide Hydrocyanic HCN water H2O Hydrogen sulfide H2S HS- ion HS- Water H2O Conjugate bases e.g. HCOO- of weak acids Conjugate acids NH4+ of weak bases 17  Conjugate acids and conjugate bases Ø Conjugate acids The conjugate acid of a base is what results after the base has accepted a proton This species is an acid because it can give up a proton to re-form the original base Ø Conjugate bases The conjugate base of an acid is what remains after the acid has donated a proton This species is a base because it can accept a proton to re-form the original acid 18 IV. Buffer system and mechanism of blood pH 19 What is buffer? Ø Definition An aqueous solution that resists changes in pH upon the addition of an acid or a base Adding water to a buffer or removing water from the buffer does not change the pH of a buffer significantly Constituent a pair of a weak acid and its conjugate base, or a pair of a weak base and its conjugate acid Ø Preparation of buffer mixing a large volume of a weak acid with its conjugate base (eg. acetic acid – acetate ion, CH3COOH – CH3COO-) mixing a large volume of weak base with its conjugate acid (eg. ammonia – ammonium ion, NH3 – NH4+) Ø Example of buffer——Phosphate buffer Consists of a weak base (HPO42-) and its conjugate acid (H2PO4-) The pH of a phosphate buffer is usually maintained at a physiological pH of 7.4 20  How does a buffer work? When a strong acid is added, the base present in the buffer neutralizes the hydrogen ions (H+) When a strong base is added, the acid present in the buffer neutralizes the hydroxide ions (OH-) Example 1: Weak acid and its conjugate base acetic acid & its conjugate base acetate ion (CH3COOH & CH3COO-) H+ H+ OH- OH- CH3COO- + H3O+ CH3COOH + H2O CH3COOH + OH- CH3COO- + H2O already exist in buffer already exist in buffer 21  How does a buffer work? Example 2: Weak base and its conjugate acid Ammonia & its conjugate base ammonia ion (NH3 & NH4+) H+ H+ OH- OH- NH3 + H3O+ NH4+ + H2O NH4+ + OH- NH3 + H2O already exist in buffer already exist in buffer 22 Henderson-Hasselbalch approximation For a weak acid HA and its conjugate base A− : HA H+ + A- If [A-] = [HA], log ([A-]/[HA])=log1=0, pH=pKa For a weak base B and its conjugate acid HB+ : pH+pOH=14 B + H2O OH- + HB+ If [HB+] = [B], log ([HB+]/[B])=log1=0, pOH=pKb 23 pH+pOH=14 derived from the expression of water ionization: Kw = [H+][OH–]=1x10-14 To get the pH and pOH correlation, we take a negative log of both sides of the equation: -log [H+] + (- log [OH–]) = – log 10-14 Notice that the terms on the left represent the pH and pOH of the solution (-log [H+] = pH, – log [OH–] = pOH). Therefore, we can write that: pH + pOH = 14 24 Kw = Ka x Kb Ø We know that Kw = [H+][OH−] = 1.0×10−14 Ø For a buffer system Dissociation of a weak acid: HA H+ + A- Kw = Ka x Kb is also true for weak base and its conjugate acid Dissociation of its conjugate base: A- + H2O OH- + HA HA [A-] Ka x Kb = x = [H+] [OH−] = Kw 25 Blood as a buffer Human blood contains a buffer of carbonic acid (H2CO3) and bicarbonate anion (HCO3—) maintain blood pH between 7.35 and 7.45 a value higher than 7.8 or lower than 6.8 can lead to death Blood buffer is: HCO3- + H+ H2CO3 With the following simultaneous equilibrium H2CO3 CO2+ H2O When acidic substance enters blood: HCO3- + H+ H2CO3 When basic substance enters blood: H2CO3 + OH- HCO3- + H2O 26 Disorders of acid-base balance Ø Acidosis Acidosis occurs when the blood pH falls below 7.35 It can be classified into two main types: q Respiratory acidosis Caused by an accumulation of carbon dioxide (CO₂) due to respiratory issues Can result from anything that interferes with respiration, such as pneumonia, asthma and chronic obstructive pulmonary disease (COPD) Symptoms may shortness of breath, confusion, drowsiness, and cyanosis (a bluish-purple color of the skin) q Metabolic acidosis Results from an excess of acids or a loss of bicarbonate (HCO₃⁻) Common causes include diabetic ketoacidosis, renal failure, and lactic acidosis Symptoms may include rapid breathing or long, deep breathing, fatigue, confusion or dizziness 27 Disorders of acid-base balance Ø Alkalosis Alkalosis occurs when the blood pH rises above 7.45 It can be classified into two main types: q Respiratory alkalosis Caused by excessive loss of CO₂ Often due to hyperventilation from anxiety, pain, or respiratory diseases Hyperventilation is breathing that is deeper or more frequent than normal Symptoms may include dizziness, tingling in the extremities, and muscle cramps q Metabolic alkalosis Results from an excess of bicarbonate (HCO₃⁻) or a loss of acids Causes can include prolonged vomiting, diuretic use, or excessive antacid consumption Symptoms can include muscle twitching and hand tremors 28 Disorders of acid-base balance Ø Diagnosis Diagnosis usually involves measuring arterial blood gases (ABGs) to assess pH, CO₂, and bicarbonate levels Ø Treatment Administering bicarbonate for acidosis or acidifying agents for alkalosis Adjusting ventilation in respiratory disorders Intravenous fluids and electrolytes to correct metabolic imbalances 29 References Devlin, T. M. (2011). Textbook of biochemistry : with Clinical Correlations (7th ed). John Wiley & Sons. Litwack, G. (2022). Human Biochemistry (2nd ed.) Academic Press. Poian, A. D. T., & Castanho, M. A. R. B. (2021). Integrative Human Biochemistry: A Textbook for Medical Biochemistry (2nd ed. 2021 ed.) Springer Tansey, J. T. (2020). Biochemistry: An Integratve Approach with Expanded Topics, WileyPLUS NextGen Card with Loose-leaf Set Multi-Semester: An Integrative Approach with Expanded Topics (1st ed.). Wiley (WileyPLUS Products). 30