Course Competency 3. - Apply basic chemical concepts to the study of human physiology.
25. Define buffer and briefly explain the importance of buffer systems in maintaining homeostasis.
A buffer is a chemical substance that helps maintain a relatively constant pH in a solution, even in the face of addition of acids or bases. Buffering is important in living systems as a means of maintaining a fairly constant internal environment, also known as homeostasis. The maintenance of blood pH is regulated via the bicarbonate buffer. This system consists of carbonic acid and bicarbonate ions. When the blood pH drops into the acidic range, this buffer acts to form carbon dioxide gas. The lungs expel this gas out of the body during the process of respiration. During alkaline conditions, this buffer brings pH back to neutral by causing excretion of the bicarbonate ions through the urine. The respiratory pigment present in blood, hemoglobin, also has buffering action within tissues. It has an ability to bind with either protons or oxygen at a given point of time. Binding of one releases the other. In hemoglobin, the binding of protons occurs in the globin portion whereas oxygen binding occurs at the iron of the heme portion. At the time of exercise, protons are generated in excess. Hemoglobin helps in the buffering action by binding these protons, and simultaneously releasing molecular oxygen.
2. Define chemical element and list the major elements, lesser elements, and trace elements.
A chemical element is the simplest form of matter to have unique chemical properties.
21. Define mixture and distinguish among three types of liquid mixtures, giving examples of each: solution, colloid, suspension.
A mixture consists of substances that are physically blended but not chemically combined. Each substance retains its own chemical properties. A solution consists of particles of matter called the solute mixed with a more abundant substance (usually water) called the solvent. The solute can be a gas, solid, or liquid—as in a solution of oxygen, sodium chloride, or alcohol in water, respectively. Solutions are defined by the following properties: *The solute particles are under 1 nanometer (nm) in size. *The solute and solvent therefore cannot be visually distinguished from each other, even with a microscope. *Such small particles don't scatter light noticeably, so solutions are usually transparent. *The solute particles can pass through most selectively permeable membranes, such as dialysis tubing and cell membranes. *The solute doesn't separate from the solvent when the solution is allowed to stand. The most common colloids in the body are mixtures of protein and water, such as the albumin in blood plasma. Many colloids can change from liquid to gel states—gelatin desserts, agar culture media, and the fluids within and between our cells, for example. Colloids are defined by the following physical properties: *The colloidal particles range from 1 to 100 nm in size. *Particles this large scatter light, so colloids are usually cloudy *The particles are too large to pass through most selectively permeable membranes. *The particles are still small enough, however, to remain permanently mixed with the solvent when the mixture stands. The blood cells in our blood plasma exemplify a suspension. Suspensions are defined by the following properties: *The suspended particles exceed 100 nm in size. *Such large particles render suspensions cloudy or opaque. *The particles are too large to penetrate selectively permeable membranes. *The particles are too heavy to remain permanently suspended, so suspensions separate on standing. Blood cells, for example, form a suspension in the blood plasma and settle to the bottom of a tube when blood is allowed to stand without mixing
24. Define pH and understand the importance of hydrogen ion concentration to the pH concept. Understand how the pH scale can be used to describe a solution's acidity or alkalinity. Distinguish between the following: b. hydrogen ion and hydroxide ion
An Acid is a substance that gives H+ ions when dissolved in water. A Base is a substance that gives OH- ions when dissolved in water.
23. Define and give specific examples of acids, bases, and salts.
An Acid is a substance that gives H+ ions when dissolved in water. Acids are compounds that contain Hydrogen (Hydrochloric, HCl; Sulphuric, H2SO4; Nitric, HNO3). However, not all compounds that contain Hydrogen are acids (Water, H2O; Methane, CH4). Acids are usually compounds of non metals with Hydrogen and sometimes Oxygen. The three acids above react with water in the following ways: HCl ---- (H2O) ----> H+ + Cl- H2SO4 ---- (H2O) ----> 2H+ + SO42- HNO3 ---- (H2O) ----> H+ + NO3- Sulphuric, Hydrochloric and Nitric acids are inorganic. There are also organic acids. Acetic acid (found in vinegar) has the formula CH3CO2H. Not all the Hydrogen atoms give H+ ions in water. In acetic acid, only the Hydrogen attached to the Oxygen yield a H+ ion. CH3CO2H ---- (H2O) ----> H+ + CH3CO2- Many acids only show acidic properties when water is present. Acids are corrosive and can burn flesh and dissolve metal. A Base is a substance that gives OH- ions when dissolved in water. Bases are usually metal hydroxides (MOH). Examples include Sodium Hydroxide, NaOH, Calcium Hydroxide, Ca(OH)2. The solution of a base in water is called an alkali. The two bases above react with water in the following ways: NaOH ---- (H2O) ----> Na+ + OH- Ca(OH)2 ---- (H2O) ----> Ca2+ + 2OH- Alkalis have a soapy feel and can corrode. A Salt results when an acid reacts with a base. Both are neutralised. The H+ and OH- ions combine to form water. The non metalic ions of the acid and the metal ions of the base form the salt. Acid + Base ----> Salt + Water Examples: HCl + NaOH ----> NaCl + H2O H2SO4 + Ca(OH)2 ----> CaSO4 + H2O NaCl is Sodium Chloride (common salt); CaSO4 is Calcium Sulphate. The salt ions normally stay in solution. The salt crystalizes out when the water is removed. Some salts are insoluble. They will precipitate out when the acid and base are added together. Barium Chloride behaves like this: 2HCl + Ba(OH)2 ----> 2H2O + BaCl2 Barium Chloride (BaCl2) appears as a white precipitate. Salts can be formed in other ways. When metals dissolve in acids, a salt is formed along with Hydrogen: Acid + Metal ----> Salt + Hydrogen In the example below Magnesium dissolves in Sulphuric Acid to give Magnesium Sulphate and Hydrogen which appears in the form of bubbles: H2SO4 + Mg ----> MgSO4 + H2 Many metal carbonates are unstable. When they dissolve in acids, a salt is formed along with water and Carbon Dioxide: Acid + Metal Carbonate ----> Salt + Water + Carbon Dioxide In the example below Calcium Carbonate dissolves in Hydrochloric Acid: 2HCl + CaCO3 ----> CaCl2 + H2O + CO2 Calcium Carbonate has many forms (chalk, marble and limestone) which are used for building. These substances are corroded by even weak acids as in acid rain.
12. Describe and give examples of the three major types of chemical bonds: a. ionic bonds
An ionic bond is the attraction of a cation to an anion. Sodium (Na+) and chloride (Cl−) ions, for example, are attracted to each other and form the compound sodium chloride (NaCl), common table salt. Ionic compounds can be composed of more than two ions, such as calcium chloride, CaCl2. Ionic bonds are weak and easily dissociate (break up) in the presence of something more attractive, such as water. The ionic bonds of NaCl break down easily as salt dissolves in water, because both Na+ and Cl− are more attracted to water molecules than they are to each other.
4. Define and discuss the structure of an atom by describing the major subatomic particles (proton, neutron, and electron) and their typical arrangement within the atom.
At the center of an atom is the nucleus, composed of protons and neutrons. Protons (p+) have a single positive charge and neutrons (n0) have no charge. Each proton or neutron weighs approximately 1 atomic mass unit (amu). The atomic mass of an element is approximately equal to its total number of protons and neutrons. Around the nucleus are one or more clouds of electrons (e−), tiny particles with a single negative charge and very low mass. Electrons swarm about the nucleus in concentric regions called electron shells (energy levels). The more energy an electron has, the farther away from the nucleus its orbit lies. Each shell holds a limited number of electrons. The elements known to date have up to seven electron shells, but those ordinarily involved in human physiology do not exceed four. Electrons of the outermost shell, called valence electrons, determine the chemical bonding properties of an atom
11. Define chemical bond and explain how the octet rule can be used to predict how atoms will behave chemically. Define electron shell, valence shell, and electronegativity.
Chemical bonding describes a variety of interactions that hold atoms together in chemical compounds. The octet rule is a chemical rule of thumb that reflects the observation that main group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. Electrons swarm about the nucleus in concentric regions called electron shells (energy levels). The more energy an electron has, the farther away from the nucleus its orbit lies. Each shell holds a limited number of electrons. The elements known to date have up to seven electron shells, but those ordinarily involved in human physiology do not exceed four. Electrons of the outermost shell, called valence electrons, determine the chemical bonding properties of an atom. Electronegativity is an atom's tendency to attract electrons to itself in a chemical bond.
16. Describe energy transfer in chemical reactions by explaining the following: c. chemical energy
Chemical energy is potential energy stored in the bonds of molecules. Chemical reactions release this energy and make it available for physiological work.
14. Define chemical reaction and explain the significance of the following terms to chemical reactions: reactant, products, metabolism.
Chemical reactions occur when chemical bonds between atoms are formed or broken. The substances that go into a chemical reaction are called the reactants, and the substances produced at the end of the reaction are know as the products. All the chemical reactions in the body are collectively called metabolism.
1. Define chemistry and matter.
Chemistry is the study of matter - it's composition, properties and transformations. Matter is anything that has mass and takes up space. Matter can be either naturally occurring or man-made.
13. Describe and give examples of ionic compounds and covalent compounds. Explain the dissociation (ionization) of an ionic compound.
Compounds are classified as ionic or molecular (covalent) on the basis of the bonds present in them. Ionic compounds are made up of electrically charged atoms or molecules as a result of gaining or losing electrons. Ions of opposite charges form ionic compounds, usually as a result of a metal reacting with a nonmetal. For example, when each sodium atom in a sample of sodium metal (group 1) gives up one electron to form a sodium cation, Na+, and each chlorine atom in a sample of chlorine gas (group 17) accepts one electron to form a chloride anion, Cl−, the resulting compound, NaCl, is composed of sodium ions and chloride ions in the ratio of one Na+ ion for each Cl− ion. Covalent, or molecular, compounds generally result from two nonmetals reacting with each other. The elements form a compound by sharing electrons, resulting in an electrically neutral molecule.
12. Describe and give examples of the three major types of chemical bonds: b. covalent bonds (including single, double, triple, polar and nonpolar covalent bonds)
Covalent bond in which electrons are more attracted to one nucleus than to the other, resulting in slightly positive and negative regions in one molecule. May be single or double. A polar covalent bond, in which electrons orbit one nucleus significantly more than the other, as represented by the asymmetric cloud. This results in a slight negative charge (δ-) in the region where the electrons spend most of their time, and a slight positive charge (δ+) at the other pole.
7. Define and explain the significance of atomic number and mass number.
Each element is identified by an atomic number, the number of protons in its nucleus. The periodic table of the elements arranges the elements in order by their atomic numbers. The atomic mass of an element is approximately equal to its total number of protons and neutrons.
17. Describe the different types of chemical reactions: synthesis reactions (anabolism), decomposition reactions (catabolism), exchange reactions, reversible reactions.
In exchange reactions, two molecules exchange atoms or groups of atoms; AB+CD⟶AC+BD For example, when stomach acid (HCl) enters the small intestine, the pancreas secretes sodium bicarbonate (NaHCO3) to neutralize it. The reaction between the two is NaHCO3+HCl⟶NaCl+H2CO3 We could say the sodium atom has exchanged its bicarbonate group (─HCO3) for a chlorine atom. Reversible reactions can go in either direction under different circumstances. For example, carbon dioxide combines with water to produce carbonic acid, which in turn decomposes into bicarbonate ions and hydrogen ions: see pic. Catabolism (ca-TAB-oh-lizm) consists of energy-releasing decomposition reactions. Such reactions break covalent bonds, produce smaller molecules from larger ones, and release energy that can be used for other physiological work. Anabolism (ah-NAB-oh-lizm) consists of energy-storing synthesis reactions, such as the production of protein or fat. Anabolism is driven by the energy that catabolism releases, so endergonic and exergonic processes, anabolism and catabolism, are inseparably linked.
8. Define and distinguish among the following: ion, cation, anion, molecule, and compound.
Ions are atoms or molecules which have gained or lost one or more valence electrons, giving the ion a net positive or negative charge. Cations are ions with a net positive charge. Anions are ions with a net negative charge. Because they have opposite electrical charges, cations and anions are attracted to each other. Cations repel other cations and anions repel other anions.
5. Define isotope and distinguish between stable and radioactive isotopes.
Isotopes are atoms of the same element that have different numbers of neutrons but the same number of protons and electrons. The difference in the number of neutrons between the various isotopes of an element means that the various isotopes have different masses. Although different isotopes of an element exhibit identical chemical behavior, they differ in physical behavior. Many of them are unstable and decay (break down) to more stable isotopes by giving off radiation. Unstable isotopes are therefore called radioisotopes, and the process of decay is called radioactivity Stable isotopes do not decay into other elements.
12. Describe and give examples of the three major types of chemical bonds: b. covalent bonds (including single, double, triple, polar and nonpolar covalent bonds)
It's formed by the sharing of electrons. For example, two hydrogen atoms share valence electrons to form a hydrogen molecule, H2 A single covalent bond is the sharing of a single pair of electrons. It is symbolized by a single line between atomic symbols, for example, H─H
9. Distinguish between molecular formula and structural formula.
Molecular formulas describe the exact number and type of atoms in a single molecule of a compound. The constituent elements are represented by their chemical symbols, and the number of atoms of each element present in each molecule is shown as a subscript following that element's symbol. A structural formula is used to indicate not only the number of atoms, but also their arrangement in space.
8. Define and distinguish among the following: ion, cation, anion, molecule, and compound.
Molecules are chemical particles composed of two or more atoms united by a chemical bond. The atoms may be identical, as in nitrogen (N2), or different, as in glucose (C6H12O6). Molecules composed of two or more elements are called compounds. Oxygen (O2) and carbon dioxide (CO2) are both molecules, because they consist of at least two atoms; but only CO2 is a compound, because it has atoms of two different elements.
10. Define isomer.
Molecules with identical molecular formulae but different arrangements of their atoms
15. Define energy and explain the difference between potential energy and kinetic energy.
Potential energy is energy contained in an object because of its position or internal state but that is not doing work at the time. Kinetic energy is energy of motion, energy that is doing work. Chemical energy is potential energy stored in the bonds of molecules. Chemical reactions release this energy and make it available for physiological work. Heat is the kinetic energy of molecular motion. The temperature of a substance is a measure of rate of this motion, and adding heat to a substance increases molecular motion. Electromagnetic energy is the kinetic energy of moving "packets" of radiation called photons. The most familiar form of electromagnetic energy is light. Electrical energy has both potential and kinetic forms. It is potential energy when charged particles have accumulated at a point such as a battery terminal or on one side of a cell membrane;
22. Explain what is meant by the concentration of a solution, and understand the difference between concentrations expressed as a percentage or as moles per liter.
Solutions are often described in terms of their concentration—how much solute is present in a given volume of solution. Percentage. This is the weight of solute as a percentage of solution volume (weight per volume, w/v) or volume of a liquid as a percentage of total solution volume (volume per volume, v/v). For example, a common intravenous fluid is D5W, which means 5% w/v dextrose in distilled water. Ethanol is often used as a 70% v/v solution. Molarity. One mole of a chemical is the number of grams equal to its molecular weight, and molarity (M) is a measure of the number of moles of solute per liter of solution. This reflects not merely the weight of solute in the solution, but the number of molecules per volume. It is the number of molecules, not their total weight, that determines the physiological effect of a solution, so molarity is often the most meaningful measure of concentration. Body fluids are usually quantified in millimolar (mM) concentrations, since they are much less than 1 molar.
19. Explain the importance of the water molecule to living systems by describing the following characteristics of water: solvent, medium for chemical reactions, heat capacity, lubricant. Contrast hydrophilic and hydrophobic.
Solvency is the ability to dissolve other chemicals. Water is called the universal solvent because it dissolves a broader range of substances than any other liquid. Substances that dissolve in water, such as sugar, are said to be hydrophilic (HY-dro-FILL-ic); the relatively few substances that do not, such as fats, are hydrophobic (HY-dro-FOE-bic). Virtually all metabolic reactions depend on the solvency of water. Biological molecules must be dissolved in water to move freely, come together, and react. The solvency of water also makes it the body's primary means of transporting substances from place to place. The chemical reactivity of water is its ability to participate in chemical reactions. Not only does water ionize many other chemicals such as acids and salts, but water itself ionizes into H+ and OH−. These ions can be incorporated into other molecules, or released from them, in the course of chemical reactions such as hydrolysis and dehydration synthesis. The high heat capacity of water also makes it a very effective coolant. When it changes from a liquid to a vapor, water carries a large amount of heat with it. One milliliter of perspiration evaporating from the skin removes about 500 cal of heat from the body. This effect is very apparent when you are sweaty and stand in front of a fan.
16. Describe energy transfer in chemical reactions by explaining the following: b. activation energy and catalysts
The energy needed to get the reaction started, supplied by the match, is called the activation energy Without catalysts, some chemical reactions proceed slowly because of the high activation energy needed to get molecules to react. A catalyst facilitates molecular interaction, thus lowering the activation energy and making the reaction proceed more rapidly. Catalysts are substances that temporarily bind to reactants, hold them in a favorable position to react with each other, and may change the shapes of reactants in ways that make them more likely to react. By reducing the element of chance in molecular collisions, a catalyst speeds up a reaction. It then releases the products and is available to repeat the process with more reactants. The catalyst itself is not consumed or changed by the reaction. The most important biological catalysts are enzymes, discussed later in this chapter.
16. Describe energy transfer in chemical reactions by explaining the following: d. law of conservation of energy
The law of conservation of energy states that energy can neither be created nor destroyed - only converted from one form of energy to another.
18. Briefly distinguish between inorganic compounds and organic compounds.
The main difference is in the presence of a carbon atom; organic compounds will contain a carbon atom (and often a hydrogen atom, to form hydrocarbons), while almost all inorganic compounds do not contain either of those two atoms. While most inorganic compounds do not contain carbon, there are a few that do. Carbon monoxide and carbon dioxide, for example, each contain carbon atoms, but the amount is not large enough to form strong bonds with the oxygen present in the molecule. Due to the small amount of carbon and the weak bonds it forms, scientists have long classified those molecules as inorganic. Organic compounds will include things like the nucleic acids, found in DNA, lipids and fatty acids found in the cells of living organisms, proteins and enzymes that are necessary for cellular processes to take place, and more. Meanwhile, inorganic compounds include the salts, metals, and other elemental compounds.
20. Describe the components of a solution by explaining the difference between solute and solvent.
When one substance dissolves into another, a solution is formed. A solution is a homogeneous mixture consisting of a solute dissolved into a solvent . The solute is the substance that is being dissolved, while the solvent is the dissolving medium. Solutions can be formed with many different types and forms of solutes and solvents.
12. Describe and give examples of the three major types of chemical bonds: b. covalent bonds (including single, double, triple, polar and nonpolar covalent bonds)
When shared electrons spend approximately equal time around each nucleus, they form a nonpolar covalent bond. the strongest of all chemical bonds. Carbon atoms bond to each other with nonpolar covalent bonds.
12. Describe and give examples of the three major types of chemical bonds: b. covalent bonds (including single, double, triple, polar and nonpolar covalent bonds)
double covalent bond is the sharing of two pairs of electrons. In carbon dioxide, for example, a central carbon atom shares two electron pairs with each oxygen atom. Such bonds are symbolized by two lines—for example, O═C═O Often occurs between carbon atoms, between carbon and oxygen, and between carbon and nitrogen.
12. Describe and give examples of the three major types of chemical bonds: c. hydrogen bonds
is a weak attraction between a slightly positive hydrogen atom in one molecule and a slightly negative oxygen or nitrogen atom in another. Water molecules, for example, are weakly attracted to each other by hydrogen bonds. Hydrogen bonds also form between different regions of the same molecule, especially in very large molecules such as proteins and DNA. They cause such molecules to fold or coil into precise three-dimensional shapes. Hydrogen bonds are represented by dotted or broken lines between atoms: ─C═O⋯H─N─. Hydrogen bonds are relatively weak, but they are enormously important to physiology.
24. Define pH and understand the importance of hydrogen ion concentration to the pH concept. Understand how the pH scale can be used to describe a solution's acidity or alkalinity. Distinguish between the following: a. proton donor and proton acceptor
pH is a measure of hydrogen ion concentration, a measure of the acidity or alkalinity of a solution. The pH scale usually ranges from 0 to 14. Aqueous solutions at 25°C with a pH less than 7 are acidic, while those with a pH greater than 7 are basic or alkaline. A pH level of 7.0 at 25°C is defined as "neutral" pH only has meaning in an aqueous solution (in water). Many chemicals, including liquids, do not have pH values. If there's no water, there's no pH. For example, there is no pH value for vegetable oil, gasoline, or pure alcohol. an acid is any proton donor, and a base is any proton acceptor.