Chapter 2 Chemistry Comes Alive PDF

Chapter 2 Chemistry Comes Alive Antoine Lavoisier was central to the eighteenth-century chemical revolution....

Chapter 2 Chemistry Comes Alive Antoine Lavoisier was central to the eighteenth-century chemical revolution. Dr. Anna Howard, MD Learning Outcomes  Matter Matter is the “stuff” of the universe. Matter is anything that occupies space and has mass. Mass is not the same as weight. The mass of an object is equal to the actual amount of matter in the object. It can be seen, smelled, and felt. Weight varies with gravity. States of Matter: Matter exists in solid, liquid, and gaseous states. Examples of each state are found in the human body. Solids, like bones and teeth, have a definite shape and volume. Liquids such as blood plasma have a definite volume, but they conform to the shape of their container. The air we breathe is a Gas.  Energy Energy is the capacity to do work, or to put matter into motion. It has no mass, does not take up space, and we can measure it only by its effects on matter. ✓ Potential energy is stored energy Forms of Energy: ✓ Kinetic energy - energy in action Chemical energy is the form stored in the bonds of chemical substances. Electrical energy results from the movement Mechanical energy is energy directly involved in moving matter Radiant energy, or electromagnetic radiation, is energy that travels in waves. 118 elements In addition - Recommended watching Mendeleev’s Periodic Table | Classification of Elements https://www.youtube.com/watch?v=0vIScP0aXN8 In addition - recommended reading Mendeleyev's Dream: The Quest for the Elements by Paul Strathern https://www.thriftbooks.com/ Four elements make up about 96% of body weight 9 elements Structure of Atoms Atom* - smallest part of an element, has a subatomic particles: central nucleus with protons (p+ positively charge) neutrons (neutral, no charge) And electrons (e- the negatively charged particles) around the center of the nucleus. All atoms are electrically neutral because the number of protons=number of electrons *Greek word “indivisible” Two models of the structure an atom  ATOMIC SYMBOL Mass of protons + Mass of neutrons = Atomic Mass The atomic number of any atom is equal to the number of protons in its nucleus and is written as a subscript to the left of its atomic symbol. *atomic mass unit = 1 amu https://youtu.be/DkbUDqDlS34 Atomic structure of the three smallest atoms H has an atomic number =1 He has an atomic number =2 Li has an atomic number =3 1. How many protons does a carbon atom have? 2. How many electrons does an oxygen atom have? 3. How many neutrons does a nitrogen atom have? The Role of Electrons in Chemical Bonding Electron shells - regions of space that surround the atomic nucleus The atoms have electrons in seven shells (the periodic table has seven rows) Each electron shell represents a different energy level (the farthest from the nucleus have the greatest potential energy) Each electron shell can hold a specific number of electrons: shell 1 accommodates only 2 electrons, shell 2 holds a maximum of 8. The number of electrons in each shell of an atom is called its electron configuration. Electrons determine the chemical behavior of atoms. When the outermost energy level of an atom contains 8 electrons - the atom is stable, or chemically inert, unreactive. Chemically inert (unreactive) and reactive elements Atomic Structure: Protons, Electrons & Neutrons https://youtu.be/EMDrb2LqL7E  How is Matter Combined? Compound vs Molecule? Atoms bound together form molecules If two or more atoms of the same element combine, the resulting substance - a molecule of that element. Ex.: a molecule of oxygen gas (O2), hydrogen gas (H2). When two or more different kinds of atoms bind, they form molecules of a compound. Compound is a substance made up of two or more different chemical elements combined in a fixed ratio. Ex.: Two hydrogen atoms combine with one oxygen atom to form the compound or molecule of water (H2O). Tri-County Technical College. Website created by Phillip Gilmour  Mixtures Figure 2.4  The three basic types of mixtures. Mixtures are substances * composed of two or more components physically intermixed. All three types of mixtures are found in both living and nonliving systems. Living material is the most complex mixture of all. Solutions are Ex.: Cytosol - the Ex.: mixture of sand homogeneous mixtures, semifluid material in and water. may be gases, liquids, or living cells solids. *you will not see the path of light.  Solutions and its concentration measurement Solutions are homogeneous mixtures of components that may be gases, liquids, or solids. Solvent - the substance present in the greatest amount. Water is the body’s chief solvent, in which gases, liquids, or solids are dissolved in. Solutes - substances present in smaller amounts (dissolved in the solvent). Concentration of Solutions may be indicated in various ways: ✓ percent (parts per 100 parts) of the solute in the total solution ✓ milligrams per deciliter (mg/dl) is a concentration measurement commonly used to measure the blood concentration of glucose, cholesterol, and so on. (A deciliter is 100 milliliters or 0.1 liter.) ✓ molarity (mo-lar′ĭ-te), or moles per liter (M.) A mole of any element or compound is equal to its atomic weight (in periodic table) or molecular weight (sum of the atomic weights) in grams.  Distinguishing Mixtures from Compounds 1. No chemical bonding occurs between the components of a mixture. Remember they are only physically intermixed. 2. It’s components can be separated by physical means— straining, filtering, evaporation, but in compounds, by contrast, can be separated only by chemical means (breaking bonds) 3. Some mixtures are homogeneous - that the mixture has exactly the same composition, and the others are heterogeneous. Check Your Understanding Which of the following is incorrectly matched? blood—suspension water (H2O)—molecule hydrogen gas (H2)—compound sugar dissolved in water—solution  What is it a chemical bonds? When atoms combine with other atoms, they are held together by chemical bonds - an energy relationship between the electrons of the reacting atoms. Three major types of Chemical Bonds: ionic, covalent, and hydrogen bonds. 1. Ionic Bond – between atoms is formed by the transfer of one or more electrons from one atom to the other. ✓ The atom that gains one or more electrons - the cation anion electron acceptor, has a negative charge – an anion. ✓ The atom that loses electrons - the electron donor has a positive charge, a cation. Most ionic compounds are formed the salt. Example: table salt - Sodium chloride (Na+CL-)  Formation of Ionic Bonds between Figure 2.6 Na (11) and Cl (17) Octet rule, or rule of eights - atoms tend to interact in such a way that they end up having 8 electrons in their valence shell.  Type of Chemical Bonds 2. Covalent (together) Bonds– electron sharing produces molecules in which the shared electrons occupy a single orbital common to both atoms and constitute covalent bonds Covalent bonds are more common in organic chemistry than ionic bonds They are also known as molecular bonds. https://youtu.be/OTgpN62ou24 Formation of double or triple covalent bonds  Polar and Nonpolar Molecules ✓ Nonpolar Molecules formed are electrically balanced because they do not have separate ‘+’ and ‘-’ poles of charge. Example: carbon dioxide (CO²) ✓ Polar Molecules – nonsymmetrical molecules that have atoms with different electron-attracting abilities. Figure 2.8 Carbon dioxide and water molecules have Example: water (H²O) different shapes, as illustrated by molecular models.  Chemical Bonding – Ionic vs. Covalent Bonds https://youtu.be/OTgpN62ou24  Hydrogen Bonds Figure 2.10 Hydrogen bonds. 3. Hydrogen Bonds form when a hydrogen atom already covalently linked to one electronegative atom (nitrogen or oxygen) is attracted by another electron-hungry atom. ✓ Hydrogen bonding is common between dipoles such as water molecules. Examples: proteins and DNA have numerous hydrogen bonds that help maintain and stabilize their structures. Check Your Understanding Which of the following is an example of a hydrogen bond? 1. bond between oxygen and hydrogen atoms 2. bond between carbon and hydrogen atoms 3. bonds between two separate water molecules 4. bond between two hydrogen atoms  Chemical Equation We can write chemical reactions in symbolic form as chemical equations. Ex.: H+H (reactants) = H2 (Product - hydrogen gas) It contains the following information: 1. The reactants: The number and kinds of the interacting substances. 2. The products: The chemical composition of the result of the reaction. 3. The relative proportions: Balanced equations indicate the relative proportion of each reactant and product.  Types of Chemical Reactions Figure 2.11 Anabolic activities in body Catabolic processes in body cells cells  Reversibility of Chemical Reactions All chemical reactions are potentially reversible: A+B B+A Such a chemical reaction is said to be in a state of chemical equilibrium (when 10 leave, 10 more may go in) Many biological reactions show so little tendency to go in the reverse direction that they are irreversible for all practical purposes (during cellular respiration)  Factors Influencing the Rate of Chemical Reactions Temperature. Higher temperatures increase the kinetic energy of particles and the force of their collisions, increasing the rate of chemical reactions Concentration. High concentrations of reacting particles increase the chances of successful collisions, and reactions progress faster Particle size. The smaller the reacting particles, the faster the chemical reaction Catalysts are substances that increase the rate of chemical reactions without themselves becoming chemically changed or part of the product. Biological catalysts are called enzymes. Check Your Understanding Which of the following is an example of a decomposition reaction? 1. transfer of a phosphate group from ATP to a glucose molecule 2. joining small molecules called amino acids into large protein molecules 3. transfer of an electron between sodium and chloride ions 4. breakdown of glycogen to multiple glucose molecules  PART 2: BIOCHEMISTRY Biochemistry is the study of the chemical composition and reactions of living matter.  Organic & Inorganic Compounds Organic Compounds – contain carbon 1. Proteins 2. Carbohydrates 3. Lipids (Fats) 4. Nucleic acids Inorganic Compounds – usually don’t contain carbon 1. Water 2. Salts 3. Acids 4. Bases  Inorganic Compounds Water – the most important inorganic compound in living material. The most important properties of water: 1. High Heat capacity - takes much energy input (in the form of heat) to change its temperature. It prevents sudden changes in temperature caused by external factors 2. High Heat of Vaporization - takes much energy input to change it from liquid to gas. This property is extremely beneficial when we sweat. 3. Water is the universal solvent - the body’s major transport medium because it is such an excellent solvent. 4. Reactivity - important reactant in many chemical reactions. 5. Cushioning – by forming a resilient cushion around certain body organs, protect them from physical trauma.  Inorganic Compounds 2. A salt is an ionic compound containing cations other than H+ and anions other than the hydroxyl ion (OH−). Salts play vital roles When salts are dissolved in water, they dissociate into in body function! their component ions. All ions are electrolytes (conduct an electrical current in solution). Salts commonly found in the body: CaCO3 (calcium carbonate) → Ca2+ + CO32- NaCl- (sodium chloride) → Na+ + Cl- KCl- (potassium chloride) → K+ + Cl- Calcium phosphates are the most plentiful salts that make bones and teeth hard.  Inorganic Compounds Example: hydrochloric acid (HCl), an acid produced 3. Acids & Bases are electrolytes by stomach cells that aids digestion. It dissociates into a proton and a chloride ion: Acid - substance that releases hydrogen ions (H+ ), is also defined as proton donor. Base - substance that takes up hydrogen ions (H+) – proton acceptor. Example: Ionization of sodium hydroxide (NaOH): It reduces the acidity NaOH Na+ + OH- + H+ H 2O + Na+ of the proton solution. sodium ion hydroxyl ion water sodium ion  pH: Acid – Base Concentration H+ pH ✓ The pH-scale is based on the number of hydrogen ions (H+) in solution, expressed in terms of moles per liter, or molarity ✓ The pH scale runs from 0 to 14 and is logarithmic – each 1 pH unit represents a tenfold change. ✓ At a pH of 7(at which [H+] is 10−7M), the solution is neutral— neither acidic nor basic. The concentrations of H+ = OH- ✓ The greater the concentration of hydroxyl ions (OH-), the more basic - alkaline the solution. ✓ The lower pH, the more acidic the solution. Solutions with a pH below 7 are acidic. ✓pH homeostasis is essential! ! Blood pH = 7.35 – 7.45! Any change is fatal! NB! Watch a D2L video OH- Sorensen's pH scale The idea for a pH scale was devised by a Danish biochemist and part-time beer brewer named Soren Sorensen in 1909. Soren Sorensen developed the simple pH scale, which measures whether a substance is acidic or basic. Søren Sørensen | Science History Institute  Neutralization What happens when acids and bases are mixed? They react with each other in displacement reactions to form water and a salt. For example, when hydrochloric acid (HCl) and sodium hydroxide (NaOH) interact, sodium chloride (a salt) and water are formed: HCl + NaOH NaCl + H2O acid base salt water Although the salt produced is written in molecular form (NaCl), remember that it actually exists as dissociated sodium and chloride ions when dissolved in water.  Buffers Buffers – chemical combinations of a weak acid and a weak base that minimize changes in pH by releasing (high pH) or binding hydrogen ions (low pH). One such buffer is the bicarbonate buffer system. This buffer system consists of carbonic acid, a week acid, and bicarbonate, its week base. *Notice! The reaction arrows go both This is the general reaction: directions. CO2+H2O H2CO3 H⁺ + HCO3- (carbon dioxide) (carbonic acid) (bicarbonate ion) This characteristic of weak acids allows them to play important roles in the chemical buffer systems of the body.  pH Homeostasis of the Blood If the blood pH varies from the limit of 7.35 – 7.45 by more than a few tenths of a unit, it may be fatal. Ex.: During CPR, an arterial pH of 7.0 predicts a poor outcome, and patients presenting with an arterial pH of less than 6.85 rarely survive. Bicarbonate buffer system - a major chemical blood buffers, in which the weak acid is carbonic acid. It dissociates reversibly, releasing its corresponding weak base, bicarbonate ions , and protons : Response to drop in pH H2CO3 HCO3¯ + H⁺ proton H+ donor Response to rise in pH H+ acceptor (weak acid = (weak base = carbonic acid) bicarbonate) The chemical equilibrium between carbonic acid and bicarbonate ion resists changes in blood pH  Organic compounds Many biological molecules (carbohydrates and proteins for example) are macromolecules. Biological molecules are formed from their monomers, or subunits, by dehydration synthesis and broken down to the monomers by hydrolysis reactions. Figure 2.14  Organic Compounds - Carbohydrates Carbohydrates, a group of molecules that includes sugar and starches, represent 1-2% of cell mass, contains carbon, hydrogen, and oxygen (water). Classification: 1. Monosaccharides (“one sugar”) are the building blocks, of the other carbohydrates. Most common monosaccharaides: glucose and fructose Bread 2. Disaccharide (“two sugars”), polysaccharide (glycogen) - polymers, linked together by dehydration synthesis, ideal sugars storage products - Disaccharide  Carbohydrate Molecules Figure 2.15 part of DNA  Carbohydrate Molecules Figure 2.15  Carbohydrate Functions: 1. Easily used source of cellular energy (simple sugar) 2. Ideal sugars storage products (glycogen) 3. Structure (proteoglycans), in our genes Review in MyA&P/A&P Place Animations Disaccharides and Polysaccharides  Check Your Understanding Glycogen is an example of a _________. 1. Polysaccharide 2. Monosaccharide 3. Disaccharide  Organic Compounds - Lipids II. Lipids – molecules that are NOT soluble in water but dissolve easily in other lipids and in organic solvents such us alcohol. Review in MyA&P/A&P Place Animations Fats  ✓ energy storage ✓ cushioning ✓ insulation Review in MyA&P/A&P Place Animations Fats  Lipids (cont.) A. Triglyceride – neutral fat, has two types of building blocks: ▪ Fatty acids ▪ Glycerol Functions: 1. Major form of stored energy in the body 2. Fat deposits (in subcutaneous tissue and around organs) protect and Review in MyA&P/A&P Place Animations Fats insulate body organs Triglycerides consist of Glycerol and Three Figure 2.16 Fatty Acids  Saturated and Unsaturated Fatty Acids Figure 2.17  Lipids (cont.) Figure 2.18 B. Phospholipids are modified triglycerides, have polar and nonpolar regions Function: ✓Form cellular membranes Review in MyA&P/A&P Place Art Labeling Lipids  Lipids (cont.) C. Steroids – made of four interlocking hydrocarbon rings. Each ring shares one or two sides with another ring. Main types: Cholesterol is the basis for all body steroids (ingest in animal products, and our liver produces some) Vitamin D is produced in the skin Sex hormones – female and male hormones: testosterone, estrogen, progesterone Adrenocortical hormones – metabolic hormones  Steroids (cont.) Functions: Figure 2.19 Steroid structure. 1. Cholesterol is found in cell membranes and is the raw material for synthesis of vit. D, steroid hormones, and bile salts 2. Steroid hormones are vital to homeostasis 3. Reproduction - sex hormones 4. Сorticosteroids.  Organic Compound - Protein III. Protein composes 10-30% of cell mass and is the basic structural material of the body. Many play vital roles in cell function. The building blocks of proteins are 20 common types of amino acids (see Appendix C), a 20-letter “alphabet” used in specific combinations to form “words” (proteins). Amino acids have two important functional groups: 1. Amino group - a base group (—NH2) 2. Carboxyl acid group (—COOH) R group that makes each amino acid chemically unique.  Organic Compounds - Protein Amino acids are linked together by peptide bonds Peptide bonds are formed by dehydration synthesis and broken by hydrolysis reaction Figure 2.21  Protein Functions Figure 2.20 Examples of protein functions. Structural Levels of Proteins Note! Tertiary Structure is essential to function, and involves the amino acids’ R-groups Polypeptide chain Two or more polypeptide chains aggregate Hemoglobin Α- helix  Fibrous and Globular Proteins Proteins are classified according to their overall appearance and shape Fibrous proteins = structural proteins insoluble in water very stable, provide mechanical support and tensile strength to the body’s tissues Examples: collagen is a composite of the helical, keratin, elastin, and certain protein of muscles Globular proteins = functional proteins compact, spherical proteins water soluble, unstable, chemically active molecules Examples: antibodies – help to provide immunity; hormones, enzymes.  Enzymes are globular proteins that act as biological catalysts The functional enzyme consists of two parts: an apoenzyme (the protein portion) and a cofactor  Enzyme Action Enzymes lower the activation energy required for a reaction. Figure 2.23  Check Your Understanding Which example of protein structure involves interactions between the R- groups of distantly positioned amino acids along the polypeptide chain? 1. tertiary structure 2. alpha helix 3. primary structure  Organic Compounds - Nucleic Acids (DNA and RNA) IV. Nucleic Acids - polymer of nucleotides composed of carbon, oxygen, hydrogen, nitrogen, and phosphorus - Two major classes of molecules: - Deoxyribonucleic acid (DNA) - Ribonucleic acid (RNA) - Each nucleotide consists of three chemical components: 1. Nitrogen-containing base 2. Pentose sugar 3. Phosphate group  Structure of DNA and RNA - Five major varieties of nitrogen-containing Double helix bases can contribute to nucleotide structure: 1. Adenine - A 2. Guanine – G Large, two-ring bases - purines 3. Cytosine – C 4. Thymine – T Smaller, single- ring bases - 5. Uracil - U pyrimidines - Two major classes of molecules: Deoxyribonucleic acid (DNA) – genetic material Ribonucleic acid (RNA) –protein synthesis The bases in DNA are A, G, C, and T, and its pentose sugar is deoxyribose (as reflected in its name) RNA bases include A, G, C, and U (U replaces the T found in DNA), and its sugar is ribose  Structure Of Nucleic Acids https://www.youtube.com/watc h?v=apaP9a079po  Nucleic Acids (cont.) Adenosine triphosphate (ATP) is the primary energy-transferring molecule in cells. Structurally, it is an adenine- containing RNA nucleotide (two additional phosphate groups have been added) When the terminal phosphate groups is cleaved off, energy is released to do useful work and ADP is formed ATP → ADP + P + ENERGY Figure 2.27 Three examples of cellular work driven by energy from ATP.  What is ATP? https://www.youtube.com/ watch?v=23ZzI6WZS28 Consistency. Discipline. Hard work. When it comes right down to it – LIFE IS CHEMISTRY THANKS FOR LISTENING!