CHAPTER 2 Cellular Chemistry objective B To describe the structure of atoms. A An atom is composed of three kinds of elementary particles: protons, neutrons, and electrons. Tey Particles are characterized by their weights(or masses)and their electric charges(table 2.2) The units for measuring weight and charge of the particles are such that a"normal"carbon atom has a weight of exactly 12, and an electron has a charge of -1 Protons and neutrons are bound in the nucleus of the atom. The number of protons in the nucleus is called the atomic number (z). The atomic number is the same for all atoms of a given chemical element. Each chemical element has a consistent number of protons in the nucleus of each of its atoms Surrounding the nucleus are precisely Zelectrons, making the atom as a whole electrically neutral. Electrons orbit the nucleus, much as the plan ets of the solar system orbit the sun. However, because electrons have properties of waves as well as particles, it is more useful to speak of energy levels occupied by the electrons. If these energy levels are imagined as organized into successive shells, then the chemical properties of the element can be explained in terms of the distribution of the z electrons among the shells. TABLE 2.2 Subatomic Particles, Weights, and Charges PARTICLE (SYMBOL) WEIGHT(APPROXIMATE) CHARGE Proton(p*t) Neutron(n) Electron(e) l/1840 2.2 Sketch structures for hydrogen(Z= 1), carbon(Z= 6), and potassium(Z= 19) The shells of an element are often represented by concentric circles around the nucleus(fig. 2. 1). The capacities of the first four shells are 2, 8, 8, and 18 electrons. The atom is built by one electron at a time, with a given shell entered only if all interior shells are full Carbon(C) Figure 2.1 Atomic representation of energy levels, or shells 2.3 What are isotopes? toms of a given element (all containing the same number [Z] of protons)but with different numbers of neutrons are said to be isotopes of the element. For example, in addition to the standard six neutron variety of carbon, there exist seven-neutron and eight-neutron varieties. The atomic weight of an element, as given in the periodic table of chemical elements, is the average of the weights of all the isotopes of the element. For example, the weight of six-neutron carbon is presented as 12.0000; however, the atomic weight of carbon is 12.01115. Because the number of neutrons in the nucleus tends to be close to the num- ber of protons, it follows from the information given in table 2. 2 that the atomic weight of an element is roughly 2Z. This rule does not hold up as well for larger atoms, but it is a fairly good estimate in the smaller atoms. Because the various isotopes of an element have a common electron shell structure, they behave identically in ordinary chemical reactions. However, the difference in weight often creates difference in stability and other properties
Objective B To describe the structure of atoms. An atom is composed of three kinds of elementary particles: protons, neutrons, and electrons. Particles are characterized by their weights (or masses) and their electric charges (table 2.2). The units for measuring weight and charge of the particles are such that a “normal” carbon atom has a weight of exactly 12, and an electron has a charge of 1. Protons and neutrons are bound in the nucleus of the atom. The number of protons in the nucleus is called the atomic number (Z). The atomic number is the same for all atoms of a given chemical element. Each chemical element has a consistent number of protons in the nucleus of each of its atoms. Surrounding the nucleus are precisely Z electrons, making the atom as a whole electrically neutral. Electrons orbit the nucleus, much as the planets of the solar system orbit the sun. However, because electrons have properties of waves as well as particles, it is more useful to speak of energy levels occupied by the electrons. If these energy levels are imagined as organized into successive shells, then the chemical properties of the element can be explained in terms of the distribution of the Z electrons among the shells. 20 CHAPTER 2 Cellular Chemistry Survey TABLE 2.2 Subatomic Particles, Weights, and Charges PARTICLE (SYMBOL) WEIGHT (APPROXIMATE) CHARGE Proton (p) 1 1 Neutron (n0 ) 10 Electron (e) 1/1840 1 2.2 Sketch structures for hydrogen (Z 1), carbon (Z 6), and potassium (Z 19). The shells of an element are often represented by concentric circles around the nucleus (fig. 2.1). The capacities of the first four shells are 2, 8, 8, and 18 electrons. The atom is built by one electron at a time, with a given shell entered only if all interior shells are full. 1p 6p 19p 6n 20n Hydrogen (H) Carbon (C) Potassium (K) Figure 2.1 Atomic representation of energy levels, or shells. 2.3 What are isotopes? Atoms of a given element (all containing the same number [Z] of protons) but with different numbers of neutrons are said to be isotopes of the element. For example, in addition to the standard six neutron variety of carbon, there exist seven-neutron and eight-neutron varieties. The atomic weight of an element, as given in the periodic table of chemical elements, is the average of the weights of all the isotopes of the element. For example, the weight of six-neutron carbon is presented as 12.0000; however, the atomic weight of carbon is 12.01115. Because the number of neutrons in the nucleus tends to be close to the number of protons, it follows from the information given in table 2.2 that the atomic weight of an element is roughly 2Z. This rule does not hold up as well for larger atoms, but it is a fairly good estimate in the smaller atoms. Because the various isotopes of an element have a common electron shell structure, they behave identically in ordinary chemical reactions. However, the difference in weight often creates a difference in stability and other properties.
CHAPTER 2 Cellular Chemistry Isotopes have important medical uses. Although all isotopes of a particular element behave iden- tically in chemical reactions, some are radioisotopes, whose radioactivity can be detected by radi- ographic instruments. Radioisotopes are frequently used by radiologists and oncologists to iagnose and treat diseases. Through injection or ingestion, a physician may introduce a radioiso- tope into the body of a patient and then track the movement, cellular uptake, tissue distribution, or excretion of the isotope in the bod Objective To describe the structure and bonds of molecule Molecules are structures composed of atoms held together by attractive forces called bonds. Ionic bonds form when atoms give up or gain electrons and become either positively or negatively charged. These charged atoms are called ions, and those with negative charges are attracted strongly to those with positive charges. Covalent bonds form when atoms share electrons. Chemical reac- tions occur when molecules form, are broken, or rearrange their component atoms. In chemical notation, ubscripts denote how many atoms of each element are in one molecule of the compound 2.4 Compute the molecular weight of water(H,O), carbon dioxide(Co, ), and glucose(CH12 o2 The molecular weight (MW) is the sum of the weights of the atoms composing the molecule(table 2.3) TABLE 2.3 The Molecular Weight of Water Carbon Dioxide, and Glucose Water(h,o) atomic weight ofH= 1 atomic weight ofO= 16 1×16=16 MW=18 Carbon dioxide(co, atomic weight of 2×16=32 MW=44 Glucose(C6H12 O2) atomic weight of c= 12 6×12=72 atomic weight ofH=1 atomic weight ofo= 16 6×16=96 MW=180 2.5 What types of bonds hold atoms together in molecules? lonic bonds. An ion is a charged atom that results from the loss or gain of one or more electrons from the atoms outer shell, causing it to lose its electrical neutrality. Atoms that gain electrons acquire an overall negative charge and are called anions. Atoms that lose electrons acquire an overall positive charge and are called cations An ionic bond is the electrical attraction that exists between an anion and a cation It is not as strong as a covalent bond in which electrons are shared rather than transferred. The Nacl (sodium chloride) molecule is held together by ionic bonding(fig 2.2). Like most ionic compounds, NaCi has a very high melting point because the molecules have a strong attraction for each other. lonic bonds dissociate easily in water Sodium atom Chlorine atom Sodium atom Chloride anion odium chloride molecule(NaCl) Figure 2.2 The formation of an ionic bond in the Nacl molecule
Isotopes have important medical uses. Although all isotopes of a particular element behave identically in chemical reactions, some are radioisotopes, whose radioactivity can be detected by radiographic instruments. Radioisotopes are frequently used by radiologists and oncologists to diagnose and treat diseases. Through injection or ingestion, a physician may introduce a radioisotope into the body of a patient and then track the movement, cellular uptake, tissue distribution, or excretion of the isotope in the body. Objective C To describe the structure and bonds of molecules. Molecules are structures composed of atoms held together by attractive forces called bonds. Ionic bonds form when atoms give up or gain electrons and become either positively or negatively charged. These charged atoms are called ions, and those with negative charges are attracted strongly to those with positive charges. Covalent bonds form when atoms share electrons. Chemical reactions occur when molecules form, are broken, or rearrange their component atoms. In chemical notation, subscripts denote how many atoms of each element are in one molecule of the compound. 2.4 Compute the molecular weight of water (H2O), carbon dioxide (CO2), and glucose (C6H12O6). The molecular weight (MW) is the sum of the weights of the atoms composing the molecule (table 2.3). CHAPTER 2 Cellular Chemistry 21 Survey TABLE 2.3 The Molecular Weight of Water, Carbon Dioxide, and Glucose Water (H2O) atomic weight of H 1 2 1 2 atomic weight of O 16 1 16 16 MW 18 Carbon dioxide (CO2) atomic weight of C 12 1 12 12 atomic weight of O 16 2 16 32 MW 44 Glucose (C6H12O6) atomic weight of C 12 6 12 72 atomic weight of H 1 12 1 12 atomic weight of O 16 6 16 96 MW 180 2.5 What types of bonds hold atoms together in molecules? Ionic bonds. An ion is a charged atom that results from the loss or gain of one or more electrons from the atom’s outer shell, causing it to lose its electrical neutrality. Atoms that gain electrons acquire an overall negative charge and are called anions. Atoms that lose electrons acquire an overall positive charge and are called cations. An ionic bond is the electrical attraction that exists between an anion and a cation. It is not as strong as a covalent bond in which electrons are shared rather than transferred. The NaCl (sodium chloride) molecule is held together by ionic bonding (fig. 2.2). Like most ionic compounds, NaCl has a very high melting point because the molecules have a strong attraction for each other. Ionic bonds dissociate easily in water. Sodium atom (Na) Chlorine atom (Cl) Sodium atom Chloride anion Sodium chloride molecule (NaCl) Figure 2.2 The formation of an ionic bond in the NaCl molecule
CHAPTER 2 Cellular Chemistry Covalent bonds. Sometimes atoms share their electrons instead of transferring them completely. They may share one, two, or three pairs of electrons. Such a sharing of electrons between two atoms is called a covalent bond. Covalent bonds are extremely strong. A shared pair is indicated by a short line drawn between the chemical symbols. For instance, in the oxygen molecule, O,, two pairs of electrons are shared (fig. 2.3), and so the molecule may be indicated as o=o xygen atom Oxygen atom Oxygen molecule Figure 2.3 The formation of a covalent bond in the O, molecule Hydrogen bonds. When hydrogen forms a covalent bond with another atom, such as oxygen, the hydro- gen atom often gains a slight positive charge as the larger oxygen atom exerts a stronger pull on the shared electron pair. The now slightly positive hydrogen atom has an affinity for the slightly negative oxygens of other molecules of the same compound, and this attraction is called a hydrogen bond (fig. 2. 4). It is not a bond that forms new molecules, but rather a weak""bond"between molecules. Hydrogen bonding is not nearly as strong as covalent or ionic bonding, but it plays an important role in determining the properties of water and many other compounds that are vital to life. Covalent bond Figure 2. 4 The configuration of hydrogen bonds between water molecules Water is a unique and special compound for many reasons. It covers about 70% of the Earths surface and is the only compound that exists in all three states(solid, liquid, and gas)in the normal temperature range of nature. It accounts for most of the body mass of every organism and has the special properties of surface tension, adhesion, cohesion, and capillary action. These properties, as well as water's char- acteristic boiling and freezing points, are due to the hydrogen bonding between water molecules. Water is known as the universal solvent and serves as the medium for nearly all biochemical reactions. In our bodies, the delicate homeostatic balance of nearly every substance depends on the presence and prope ties of water
Covalent bonds. Sometimes atoms share their electrons instead of transferring them completely. They may share one, two, or three pairs of electrons. Such a sharing of electrons between two atoms is called a covalent bond. Covalent bonds are extremely strong. A shared pair is indicated by a short line drawn between the chemical symbols. For instance, in the oxygen molecule, O2, two pairs of electrons are shared (fig. 2.3), and so the molecule may be indicated as OO. 22 CHAPTER 2 Cellular Chemistry Oxygen atom Oxygen atom Oxygen molecule Figure 2.3 The formation of a covalent bond in the O2 molecule. Hydrogen bonds. When hydrogen forms a covalent bond with another atom, such as oxygen, the hydrogen atom often gains a slight positive charge as the larger oxygen atom exerts a stronger pull on the shared electron pair. The now slightly positive hydrogen atom has an affinity for the slightly negative oxygens of other molecules of the same compound, and this attraction is called a hydrogen bond (fig. 2.4). It is not a bond that forms new molecules, but rather a weak “bond” between molecules. Hydrogen bonding is not nearly as strong as covalent or ionic bonding, but it plays an important role in determining the properties of water and many other compounds that are vital to life. H H H Water molecule Covalent bond Oxygen Hydrogen H H H Figure 2.4 The configuration of hydrogen bonds between water molecules. Water is a unique and special compound for many reasons. It covers about 70% of the Earth’s surface and is the only compound that exists in all three states (solid, liquid, and gas) in the normal temperature range of nature. It accounts for most of the body mass of every organism and has the special properties of surface tension, adhesion, cohesion, and capillary action. These properties, as well as water’s characteristic boiling and freezing points, are due to the hydrogen bonding between water molecules. Water is known as the universal solvent and serves as the medium for nearly all biochemical reactions. In our bodies, the delicate homeostatic balance of nearly every substance depends on the presence and properties of water
CHAPTER 2 Cellular Chemistry Objective d To understand the concept of moles A mole(mol)is a unit of measurement, just like a liter or a meter. It is a unit of weight, and it always contains 6.022 X 102 molecules. A mole of water therefore contains 6.022 X 102 mol- ecules of water, and a mole of helium contains exactly 6.022 X 102 helium atoms. A mole of any substance is equal to the same number of grams as the molecular weight of the substance. 2.6 How many grams do 2 moles of table salt(NaCl) weigh? molecular weight of NaCl= 23+ 35=58 116 mol 2.7 How many water molecules are in I mL (milliliter)of water? I mL H,O=I g molho= 18 1 mol ho= 6.022 x 102Ho molecules 6.022×102 molecules (mL) 3.34×1023 molecule Objective E To define the terms mixture, solution, suspension, and colloidal suspension When two or more substances combine without forming bonds with each other, the result is a rvey mixture. Solutions are mixtures in which the molecules of all the combined substances are dis- tributed homogeneously throughout the mixture. Solutions include solids dissolved in liquid, as with salt water and metals dissolved in each other. as in metal all Ispension Is a mixture in which particles of one substance are suspended in another substance but not evenly distributed down to a molecular level. The particles in a suspension will settle out of the mixture, like the dust settling out of the air in a room, but the particles of a colloidal suspension are so small that they do not settle out 2. 8 What is a solvent?A solute? Solutions are the most important kind of mixtures in organic chemistry, and most biological solutions consist of some solid substance dissolved in water. In this case. water serves as the solvent of the solu tion, and the substance, be it a salt, sugar, or protein, is the solute. A practical definition of a solvent is that it is the substance of any solution present in greatest proportion, often water. All other substances are con- sidered solutes. The distinction becomes less useful in solutions such as metal alloys, which may have equal amounts of two or more substances 2.9 How are concentrations in solution measured? Concentrations of solute in a solution may be measured in several ways, and the most appropriate way is determined by case or need. For example, it is sometimes most useful to measure the percentage of the solute in the solution Molality is a measure of the moles of solute per kilogram of solvent. Molarity(M) s a measure of the moles of solute per liter of solution. Molarity is by far the most frequently used meas- urement for biological solutions Objective F To describe acids, bases, and the ph scale In any sample of water, a certain minuscule proportion of water molecules exists in an ionized form, rvey as H*(hydrogen ions)and OH-(hydroxide ions ) In pure water, the number of H*equals the num ber of oH-, and the concentration of each is 10-7M. Chemical substances that, when added to water solutions increase the concentration of h+ are called acids: those that increase the concen- tration of OH are called bases. The acidity or basicity of a solution is expressed as a value on the pH scale, which is a number derived from the logarithm of the concentration of hydrogen ions
Objective D To understand the concept of moles. A mole (mol) is a unit of measurement, just like a liter or a meter. It is a unit of weight, and it always contains 6.022 1023 molecules. A mole of water therefore contains 6.022 1023 molecules of water, and a mole of helium contains exactly 6.022 1023 helium atoms. A mole of any substance is equal to the same number of grams as the molecular weight of the substance. 2.6 How many grams do 2 moles of table salt (NaCl) weigh? molecular weight of NaCl 23 35 58 2.7 How many water molecules are in 1 mL (milliliter) of water? 1 mL H2O 1 g 1 mol H2 O 18 g 1 mol H2O 6.022 1023 H2O molecules Objective E To define the terms mixture, solution, suspension, and colloidal suspension. When two or more substances combine without forming bonds with each other, the result is a mixture. Solutions are mixtures in which the molecules of all the combined substances are distributed homogeneously throughout the mixture. Solutions include solids dissolved in liquid, as with salt water, and metals dissolved in each other, as in metal alloys. A suspension is a mixture in which particles of one substance are suspended in another substance but not evenly distributed down to a molecular level. The particles in a suspension will settle out of the mixture, like the dust settling out of the air in a room, but the particles of a colloidal suspension are so small that they do not settle out. 2.8 What is a solvent? A solute? Solutions are the most important kind of mixtures in organic chemistry, and most biological solutions consist of some solid substance dissolved in water. In this case, water serves as the solvent of the solution, and the substance, be it a salt, sugar, or protein, is the solute. A practical definition of a solvent is that it is the substance of any solution present in greatest proportion, often water. All other substances are considered solutes. The distinction becomes less useful in solutions such as metal alloys, which may have equal amounts of two or more substances. 2.9 How are concentrations in solution measured? Concentrations of solute in a solution may be measured in several ways, and the most appropriate way is determined by case or need. For example, it is sometimes most useful to measure the percentage of the solute in the solution. Molallty is a measure of the moles of solute per kilogram of solvent. Molarity (M) is a measure of the moles of solute per liter of solution. Molarity is by far the most frequently used measurement for biological solutions. Objective F To describe acids, bases, and the pH scale. In any sample of water, a certain minuscule proportion of water molecules exists in an ionized form, as H (hydrogen ions) and OH (hydroxide ions). In pure water, the number of H equals the number of OH, and the concentration of each is 107 M. Chemical substances that, when added to water solutions, increase the concentration of H are called acids; those that increase the concentration of OH are called bases. The acidity or basicity of a solution is expressed as a value on the pH scale, which is a number derived from the logarithm of the concentration of hydrogen ions. CHAPTER 2 Cellular Chemistry 23 Survey 2 mol 116 58g mol g ⎛ ⎝ ⎜ ⎞ ⎠ ⎟ = (mL) 1g mL 1mol 18g ⎛ molecu ⎝ ⎜ ⎞ ⎠ ⎟ ⎛ ⎝ ⎜ ⎞ ⎠ ⎟ 6 022 10 × 23 . les mol molecules ⎛ ⎝ ⎜ ⎞ ⎠ ⎟ = × 3 34 1023 . Survey Survey
CHAPTER 2 Cellular Chemistry 2.10 Demonstate how the pH scale works The pH of a substance is determined by taking the negative logarithm of the H* concentration of a solu- tion. Because water has a hydrogen ion concentration of 10-M, its pHis 7. As H+ concentration increases, the negative logarithmic value decreases, and vice versa. Therefore, basic solutions have a pH higher than 7, and acidic solutions have a pH lower than 7. 2.11 What is a strong acid?A weak acid? Strong acids are acids that dissociate completely in water; in other words, every one of the acid molecules loses its proton in the water solution. Examples of strong acids are hydrochloric acid (HCi) and sulfuric acid(H, SO,). Weak acids are acids that only partially dissociate; in other words, some but not all of the molecules lose their protons in the water solution. Mole for mole, strong acids generally change the ph of a solution more significantly than do weak acids. However, weak acids and the salts they form are extremely important in organic chemistry, as they are the basis of buffers 2.12 Define the term salt Salts are ionic compounds formed from the residue of an acid and the residue of a base. When an acid loses its proton, and a base loses a hydroxyl group(OH ) the remaining ions of the molecules, if both are pres- ent in the solution, will sometimes bind to each other forming a salt. The reaction of HCl (an acid) with NaoH (a base)to form table salt(Nacl) is an example HC+NaOH→H2O+NaCl objecti ive G To define buffer. a abuffer is a combination of a weak acid and its salt in a solution that has the effect of stabiliz- ing the pH of the solution. If a solution contains a buffer, its pH will not change dramatically even when strong acids or bases are added when acid is added to the solution it is neutralized by the salt of the weak acid. When base is added to the solution, it is neutralized by the weak acid 2.13 What is the pH of blood, and how is it maintained at a constant level? Blood has a pH of 7. 4, which means it is slightly more basic than water. Blood maintains its pH in home ostasis(steady state)by means of the bicarbonate buffer system, which is regulated by the amount of carbon dioxide dissolved in the blood. The acid of the buffer system is carbonic acid, H, COz, which forms from carbon dioxide and water. The salt is sodium bicarbonate. which exists in solution as bicar- bonate ions. HcO 2.14 List the most important buffer systems in the body and indicate their locations. See table 2. 4 TAbLe 2. 4 Buffer Systems and Their locations Bicarbonate buffer Blood, extracellular fluid(most easily adjusted body buffer) Protein buffe All tissues(most plentiful body buffer ctive H To distinguish between inorganic and organic compound und ey small molecules. Organic compounds always contain carbon and are held together by covalent bonds. Organic compounds are usually large, complex molecules. Both inorganic and organic com- pounds are important in biochemistry, the study of chemical processes that are essential to life
2.10 Demonstate how the pH scale works. The pH of a substance is determined by taking the negative logarithm of the H concentration of a solution. Because water has a hydrogen ion concentration of 107 M, its pH is 7. As H concentration increases, the negative logarithmic value decreases, and vice versa. Therefore, basic solutions have a pH higher than 7, and acidic solutions have a pH lower than 7. 2.11 What is a strong acid? A weak acid? Strong acids are acids that dissociate completely in water; in other words, every one of the acid molecules loses its proton in the water solution. Examples of strong acids are hydrochloric acid (HCl) and sulfuric acid (H2SO4). Weak acids are acids that only partially dissociate; in other words, some but not all of the molecules lose their protons in the water solution. Mole for mole, strong acids generally change the pH of a solution more significantly than do weak acids. However, weak acids and the salts they form are extremely important in organic chemistry, as they are the basis of buffers. 2.12 Define the term salt. Salts are ionic compounds formed from the residue of an acid and the residue of a base. When an acid loses its proton, and a base loses a hydroxyl group (OH), the remaining ions of the molecules, if both are present in the solution, will sometimes bind to each other, forming a salt. The reaction of HCl (an acid) with NaOH (a base) to form table salt (NaCl) is an example: Objective G To define buffer. A buffer is a combination of a weak acid and its salt in a solution that has the effect of stabilizing the pH of the solution. If a solution contains a buffer, its pH will not change dramatically even when strong acids or bases are added. When acid is added to the solution, it is neutralized by the salt of the weak acid. When base is added to the solution, it is neutralized by the weak acid itself. 2.13 What is the pH of blood, and how is it maintained at a constant level? Blood has a pH of 7.4, which means it is slightly more basic than water. Blood maintains its pH in homeostasis (steady state) by means of the bicarbonate buffer system, which is regulated by the amount of carbon dioxide dissolved in the blood. The acid of the buffer system is carbonic acid, H2CO3, which forms from carbon dioxide and water. The salt is sodium bicarbonate, which exists in solution as bicarbonate ions, HCO3 . 2.14 List the most important buffer systems in the body and indicate their locations. See table 2.4. 24 CHAPTER 2 Cellular Chemistry HCl + NaOH H O + NaCl 2 acid base water salt → Survey TABLE 2.4 Buffer Systems and Their Locations Bicarbonate buffer Blood, extracellular fluid (most easily adjusted body buffer) Phosphate buffer Kidneys, intracellular fluid Protein buffer All tissues (most plentiful body buffer) Objective H To distinguish between inorganic and organic compounds. Inorganic compounds do not contain carbon (exceptions include CO and CO2 ) and are usually small molecules. Organic compounds always contain carbon and are held together by covalent bonds. Organic compounds are usually large, complex molecules. Both inorganic and organic compounds are important in biochemistry, the study of chemical processes that are essential to life. Survey