HTHS 1110 Unit 02 Textbook PDF

2 The Atomic Level of Organization Objective 1 Objective 8 Parts of the Atom States of Matter Objective 2 Objective 9 Photons Diffusion Objective 3 Objective 10 Isotopes Energy & the Chemical...

2 The Atomic Level of Organization Objective 1 Objective 8 Parts of the Atom States of Matter Objective 2 Objective 9 Photons Diffusion Objective 3 Objective 10 Isotopes Energy & the Chemical Bond Objective 4 Electrons as the Objective 11 Source of Chemical Special Properties of Water Properties Objective 12 Objective 5 Acids, Bases & the Concept The Periodic Table of pH Objective 6 Objective 13 Noble Gases in the Solutions Periodic Table Objective 14 Objective 7 Functional Groups & Atomic Number & Polyatomic Ions Atomic Mass in the Periodic Table Objective 15 Chemical Equations 2-1  Objective 1: Parts of the Atom 1 Name the elementary particles that make up atoms: neutrons, protons, and electrons. Describe the charge, mass and relative location of neutrons, protons, and electrons. Atoms are made up of neutrons, protons, and electrons. You should know the basic properties of each one and where they’re located. Note that it’s not important to know the exact mass of these particles: the masses round to one or zero and those are the values we’ll use here. You should recognize that electrons are what give atoms their chemical properties. This is due to the arrangement of electrons, a concept that will run throughout this Unit (really, throughout the whole course). We’ll start the process of figuring out those electron configurations by defining a few basic terms in this objective. You should also understand that our conceptions of the nucleus, protons, and electrons are really just metaphors: stories that scientists tell themselves based on the observable data. The only purpose of these metaphors or stories is to help us model how the world works. An important process in Unit 2 is to understand the stories so that you can understand how the human body works from the atomic level up. All matter is made up of atoms. An atom is a particle that contains all the characteristics of an element. That is, one atom acts chemically like two hundred, or two thousand, or two billion atoms. An atom used to be an imaginary thing. Now, with certain kinds of imaging, scientists can actually “see” atoms directly. Atoms make up the smallest unit of matter that retain the properties of their associated element. A pure element (like the carbon in a diamond, or helium gas in a balloon) consists of atoms, all of which have identical chemical properties. 2-1 The most common elements in the human body (and the symbols which represent them) are: ꞏ oxygen (O) Element Life (Humans) Atmosphere Earth's Crust ꞏ carbon (C) Oxygen (O) 65% 21% 46% ꞏ hydrogen (H) Carbon (C) 18% trace trace ꞏ nitrogen (N) Hydrogen (H) 10% trace 0.1% The graph indicates Nitrogen (N) 3% 78% trace approximate percentages of elements in living organisms (humans) compared to the nonliving world. These are the elements that we will focus on in this course. Because elements are made by fusion reactions in the sun, the smaller, lighter elements are more common than the larger ones here on earth. We will list the elements that you need to know a bit later but we’ll start with the “top four”. Properties of atoms that we will learn in this Unit come from both theoretical and experimental models of atoms and elements. The diameter of an atom is about 10–10 m (0.1 nm, 0.0000001 mm, or 0.0000000001 m). An old name for this unit is an Ångstrom (Å). The diameter of a molecule is about 10 times that, or 1 nm. For this reason, when scientists build devices that work at the molecular (nanometer) scale, it’s called nanotechnology. There are things smaller than atoms: subatomic particles. Physicists have lots of subatomic particles, but for our purposes, we will consider atoms to be made up of three elementary subatomic particles: ꞏ neutrons ꞏ protons ꞏ electrons We’ll consider each of these in turn. 2-2 The nucleus is in the center of an atom and contains the massive (i.e. heavy) particles that make up atoms. The mass of atoms is measured by a unit that is (naturally enough) called the atomic mass unit (amu) or Daltons (Da). A neutral atom of a particular element has an equal number of electrons and protons (as we’ll see later, this means the + and – charges cancel out, hence the term “neutral”) but the number of neutrons can vary. This variation changes the atom’s mass but not its chemical properties. Atoms which vary in the number of neutrons are called isotopes. We’ll take a closer look at isotopes in Objective 3. Neutrons are located in the nucleus. They have no charge, but have a mass of 1.009 Da. For the purpose of this course, we round that to 1 Da. The only time neutrons are apparent is when atoms break apart in a process called nuclear fission. Nuclear fission is the basis for nuclear weapons, nuclear reactors, and nuclear medicine. Protons are the other particles found in the nucleus. They carry a positive charge. Protons have a mass of 1.007 Da, but for our purposes, the mass is 1 Da. Like neutrons, they are sequestered in the nucleus and cannot participate in chemical reactions. Also like neutrons, the only time they are apparent is in nuclear fission. We indicate the number of protons (called the atomic number) with a subscript number in front of the element symbol: 8 O 6 C 1 H 7 N 2-3 Electrons are small and light. (The weight is 9 × 10–28 g, or 0.0005 Da, 1/2000 the weight of a neutron or proton, but you do not have to learn it.) They have a negative charge. They are distant from the nucleus, so they are hanging out looking for things to do, like contributing to the chemical properties of the atom. Electrons have been conceived as a number of different ideas: as particles, as waves, or as packets of negative charge (quanta) with a certain probability of existing in a specific place. That is, according to quantum theory, we can never really know where an electron is. Rather, we can just say where it’s more likely or less likely to be. We will see later that atoms readily gain or lose electrons. A chemical reaction can alter the position (or probability function) of electrons, but a chemical reaction does not change anything about the nucleus—by definition, only a nuclear reaction can do that. In a neutral atom, the number of packets of positive charge (protons) and the number of packets of negative charge (electrons) are exactly equal: all the positives cancel out all the negatives and the net charge of a neutral atom is zero. 2-4 But the number of electrons can change. When the number of electrons is no longer equal to the number of protons, we call this an ion. An ion has either more positive charges or more negative charges. A cation has more protons than electrons (i.e. an excess of positive charges). An anion has more electrons than protons (i.e. an excess of negative charges). We indicate the excess charges with a superscript after the element’s symbol: 8 O 8 protons, 2 electrons: O6+ 8 protons, 10 electrons: O2– 6 C 6 protons, 2 electrons: C4+ 6 protons, 10 electrons: C4– 1 H 1 proton, 0 electrons: H1+ (usually represented as H+) 1 proton, 2 electrons: H1– (usually represented as H–) 7 N 7 protons, 2 electrons: N5+ 7 protons, 10 electrons: N3– (Note that the number of protons is equal to the atomic number, as explained earlier. Also note that ions “like” to have zero, 2, or 10 electrons. This will be discussed later.) 2-5 Electrons sort themselves into energy levels. Electrons closer to the nucleus are at a lower energy level, while electrons further away from the nucleus are at a higher energy level. In one model (called the Bohr or planetary model), these energy levels are represented by concentric circles. Because quantum theory states that electrons are a probability distribution of the location of negative charges, a more complete understanding of electrons is to represent them as circular or pear-shaped probability clouds called orbitals. Orbitals can hold from one to seven pairs of electrons. Electrons are paired by their spin (up or down); “up” electrons like to be next to “down” electrons. 2-6  Objective 2: Photons 2 State the correct order of energy (wavelength) in the electromagnetic spectrum: visible light, ultraviolet, X-rays, gamma rays. Describe how the wavelength and energy of photons are related. Along with the three elementary particles we learned in Objective 1, there is another particle-wave we need to consider called a photon. Essentially, a photon is a packet of pure energy that is created whenever an atom sheds excess energy, for example when the nucleus rearranges or splits apart. Photons are important for an understanding of radiation medicine, because in this context the word “radiation” means the same thing as a photon. In this objective, students will master the basic properties of photons, with a focus on those used in human medicine. Photons are particles of light. Photons also behave as waves. It’s another weirdness of quantum physics that photons are both a particle and a wave. This seems weird at first, but we know of a lot of things that are a superposition of two states: you might be both an American and a college student. Just like you can be two things at once, and what you are depends on what questions we ask, so photons behave as particles in some experiments and waves in others. As particles, they carry a certain energy. For example, a bullet is a pretty small piece of metal, but it moves at a very high speed so it carries a lot of energy and can do much damage. In classical physics, MASS X VELOCITY = MOMENTUM Einstein’s Theory of Relativity says photons have zero mass but a measurable momentum. That’s strange, because anything multiplied by zero should be zero, so if mass = 0, then mass × velocity should equal zero. 2-7 Einstein also said the velocity of photons (“speed of light”) is a constant: 3 × 109 m/sec. Photons arise from an excess amount of energy that results through at least three processes that involve rearrangement of subatomic particles: ꞏ nuclear rearrangement ꞏ nuclear fission (splitting) ꞏ electron rearrangement 2-8 Visible light, ultraviolet light, X-rays and gamma rays are all “made” of photons but with different momentum (energy) and different wavelengths. A wavelength is the distance between the crest of waves. The height (amplitude) of the wave is the intensity of the signal (e.g a bright light vs a dim light). 2-9 The wavelength is what makes a difference between different forms of electromagnetic radiation. The wavelength is also related to the energy carried by the wave-particle. High energy photons have short wavelengths; low energy photons have long wavelengths. (Waves crashing on the beach several times a second have a lot more energy than a few waves a minute.) Visible light has a wavelength in the range of about 400 to 700 nm (0.4 to 0.7 μm). The longest of these, 700 nm, is also the lowest energy (red). The shortest of these, 400 nm, is also the highest energy (violet). From longest to shortest wavelength (lowest to highest energy), the colors are red, orange, yellow, green, blue, [indigo] and violet. The mnemonic for this sequence is ROY G. BiV who is said to be the “discoverer” of light. In the visible spectrum, the color violet has the highest energy and shortest wavelength. Beyond violet in the visible light spectrum violet is the invisible spectrum, starting with ultraviolet, and then X-rays and gamma rays. Ultraviolet (UV) has enough energy to cause first-degree burns to tissue (a sunburn) or to cause scar tissue to form in the tissues of the eye (cornea or lens). X-rays are photons with a high enough energy to penetrate living tissue, so they are used for medical diagnostics. When X-rays strike film, they produce a black color, so dense tissue like bone (which stops X-rays) appears white under X-ray. Gamma rays are photons with enough energy to penetrate tissue, but they can also damage living tissue. For that reason, they are used to destroy tissue that cannot be reached by surgical means. 2-10 Notice in this image that the shorter the wavelengths are, the higher the energy of the electromagnetic wave. On the low end of the energy spectrum are radio waves with the crest of the wavelengths far apart. On the high end of the spectrum are gamma rays. Look how close together the wavelengths are. The closer the wavelengths, the higher the energy. 2-11  Objective 3: Isotopes 3 Define isotope. Describe how different isotopes of the same element are formed. Distinguish between stable isotopes and radioisotopes. Summarize and discuss the forms of energy which are emitted by radioisotopes: positron emission; alpha (α), beta (β), and gamma (γ) radiation. In this objective, we are answering the theoretical question: what if we keep the number of protons and electrons in an atom the same, but change the number of neutrons? The charge won’t change, because neutrons have no charge, but the mass of the nucleus will change. That will change certain properties of the atom. The resulting atom, with a different mass but the same charge and chemical properties, is called an isotope of that element. We indicate the mass of an isotope with a superscript number in front of the element name: 16 O 14 C 3 H 12 C 1 H 14 N 13 C 2 H 15 N Hydrogen-1 (1H) is the simplest atom. Hydrogen has an atomic number of 1, so we know that it always has one proton in its nucleus. For 1H, one is also the Isotopes of Hydrogen mass number, so there are zero neutrons in the nucleus. The number of protons and electrons is always equal for a non- charged atom, so there is also one electron. Because the number and arrangement of electrons determines the chemical properties of the atom, and because a neutral atom has the same number of electrons as it does protons, we can predict the chemical properties of an element by its atomic number (number of protons). We’ll see how this prediction works when we get to the periodic table in Objective 5. Remember, the number of neutrons found in the nucleus can vary without changing the chemical properties of the element. When the number of neutrons varies, we call these isotopes of the element. Because this changes the atomic mass (superscript before the symbol), the three isotopes of hydrogen are represented like this: 1H, 2H, 3H. 2-12 2-13 For example, hydrogen always consists of one proton and one electron. If we add a neutron, so that it consists of one neutron, one proton and one electron, the isotope is called hydrogen-2 (2H) or deuterium, which is sometimes represented as D. Deuterium plus oxygen makes up so-called heavy water (D2O), a compound used as a tracer in metabolic studies. If we add one more neutron, the resulting hydrogen atom, with two neutrons, one proton and one electron, is called tritium (3H). Tritium is unstable and a tritium atom will fall apart, on average, after 12.3 years. We call this the half-life of tritium because over a 12.3 year period, there is a 50% probability a single tritium atom will fall apart. Another way of saying this is that half of the tritium atoms will fall apart in 12.3 years. When we create unstable isotopes of an atom, these are called radioisotopes. 2-14 Unstable isotopes are called radioisotopes. Not all radioisotopes decay (fall apart) in the same way. Different forms of radioisotope decay are used in human medicine. Alpha (α) particles are the same as helium nuclei. Helium is atomic number 2 and has an atomic mass of 4, but in an alpha particle the two electrons are missing. Beta (β) particles are electrons expelled at high energy from radioactive atoms. For example, 3H and 14C are beta emitters. Common isotopes of carbon are 12C, 13C, and 14 C. 14C is unstable and falls apart, on average, every 5730 years. Like 3H, 14C emits an electron (beta particle or β) in the process of falling apart. The decay of 14C is the basis for radiocarbon dating in archeology. A gamma (γ) particle (or gamma ray) is a high-energy packet of light energy (photon). Photons of lower energy (from highest to lowest) are x-rays, ultraviolet light, and visible light. One common mode of decay is by nuclear fission. In this case, more neutrons are found in the nucleus than can be supported, and the nucleus is unstable. One familiar kind of fission is the decay of 235U. This is the kind of uranium used in some nuclear weapons and reactors. When a neutron hits a 235U nucleus, the nucleus breaks apart, creating two new atomic nuclei and three neutrons. These three neutrons then collide with more 235U nuclei in a chain reaction, like the positive feedback loops we studied in Unit 1. 2-15 2-16 Proton beams are used to treat prostate, brain and lung cancers. Neutron beams have also been used experimentally in cancer treatment. Particles called positrons, which are like electrons in all respects except they carry a positive charge, are used in positron-emission tomography (PET). (We call this antimatter: the positron is an antimatter electron. Antimatter is made up of antineutrons, antiprotons, and positrons.) The magnetic properties of hydrogen nuclei are used as the basis for nuclear magnetic resonance imaging, an old name for magnetic resonance imaging (MRI). Technetium (Tc), atomic number 43, is the only element between atomic numbers 1 and 83 that does not occur naturally on Earth. Because it is not used anywhere in the body, it will not interact with the patient’s tissues. For this reason, it is used in medical diagnostic procedures. If we are referring to an atomic number, then that shows as a subscript: 43Tc. Remember the atomic number is equal to the number of protons (43 protons). In nuclear medicine, we usually find atomic symbols with the mass number also indicated. For example, technetium-99m (the m stands for “metastable”, a concept which will not be explained here), or 99MTc, is the form of Tc used in medicine. 2-17 The superscript “99” before the element name tells us that 99 is the mass number (43 protons + 56 neutrons = 99) of the artificially created atom. In the periodic table we’ll see in Objective 5, the atomic mass is in parentheses, indicating there are no technetium atoms in nature. Other atoms commonly used in nuclear medicine include 47Ca (20 protons + 27 neutrons) to study bone loss and the location of tumor cells, 131I (53 protons + 78 neutrons) to destroy thyroid tissue, and 133Xe (54 protons + 79 neutrons) for respiratory and cerebral blood flow studies. 2-18  Objective 4: Electrons as the Source of Chemical Properties 4 Illustrate how the number and distribution of electrons is related to the chemical properties of an element. In this objective, we are asking students to answer one seemingly simple question: what is chemistry? One seemingly simple answer is: electrons are chemistry. So our job at this point is to understand what the relationship between electrons and chemistry entails. Tom Cech, who was Dr. Hutchins’ teacher at the University of Colorado, won the Nobel Prize in 1989. He said: Chemistry happens because molecules and atoms collide together with the right orientation and velocity. Because those collisions involve the relatively large electron clouds, and not the relatively small and internal nuclei, it is the interaction of electrons with the collision of atoms that results in chemistry. Electrons are arranged in energy levels called orbitals, and are paired together by their up or down spin. We use orbital diagrams or orbital notation to indicate the energy levels of all the electrons in an atom of a given element. Each of these energy levels or orbitals wants to be completely filled. Once the orbital is filled, the electron cloud is in its lowest energy state (represented by a lower level on the diagram). 2-19 Aufbau is a German word that means “building up”. For this reason, the rule for filling orbitals is called the Aufbau Principle. Just like a multi-story building starts with a foundation, we fill the lowest energy orbital (the 1s) first; then the next level up (2s); then the next (2p); and so on. In the orbital diagrams shown here, the lowest energy level is Electrons have spin. Remember we said chemistry is a metaphor? This lowest on the is a picture which illustrates what no diagram. human has ever "seen", which is the spin states of the electron. In orbital diagrams, spin states are represented as an "up arrow" or "down arrow". An atom of an element achieves its lowest energy state (most favorable configuration) by placing its electrons in a way that fills all its orbitals from the lowest to the highest energy. The example shown here is nitrogen, atomic number 7. Nitrogen has 7 protons (by definition) so a neutral nitrogen atom has 7 electrons: there are 7 packets of positive charge and 7 packets of negative charge for a net charge of zero. This orbital diagram shows the arrangement of the 7 electrons in nitrogen. The first orbital we fill If this nitrogen atom adds with two electrons (one up, one down) is the three electrons with "down" 1s orbital. The second orbital we fill with two spin, its 2p orbital will be electrons is the 2s orbital. Now we’ve used filled and it will be in the lowest possible energy 4 of our 7 alloted electrons and we have 3 state. Since electrons left over. The up electron spins always go in have a -1 charge, we say first, and the 2p orbital has 3 parking spots nitrogen has a valence of to put them in, so we indicate 3 up electrons -3 (3 x -1 = -3). in those three circles. 2-20 Like filling a parking garage from the bottom floors up, we want to fill the “parking spaces” in each “floor” or orbital for maximum efficiency. Chemistry is about completing those orbital sets. So if we have one open “parking spot” we want to find an electron to fill it, and if we have one car taking up an entire floor we want to give it away to get the most efficient “parking configuration”. In real terms, an atom has less energy if the entire set of orbitals (1s, 2s, 2p, 3s, etc.) is filled. Sodium, chlorine, and many other atoms are rarely found in the body as uncharged atoms. Rather, in water, the sodium chloride (NaCl, table salt) molecule breaks into two charged atoms: Na+ and Cl–. The number of protons doesn’t change because they are buried in the nucleus; the number of electrons changes, resulting in a cation or anion. 11 Na has 11 protons and prefers to have 10 electrons: +11 – 10 = +1 17 Cl has 17 protons and prefers to have 18 electrons: +17 – 18 = –1 Cations and anions in the blood are called electrolytes (English: “electricity” + Greek: “released”) because they carry electrical charge. When atoms like oxygen, sulfur, selenium, fluorine, chlorine, or iodine are present as ions (either in a compound or in solution) we refer to them as oxide, sulfide, selenide, fluoride, chloride, or iodide. For example, NaCl is called sodium chloride. 2-21  Objective 5: The Periodic Table 5 Classify the arrangement of elements in a periodic table. Define the chemical properties which result in the arrangment of elements in a periodic table: valence and electronegativity. State the relationship between valence and the position of an element on the periodic table. Categorize the electronegativity of elements and describe how electronegativity and periodic table position are related. Elements are listed in the periodic table in order of atomic number (number of protons and electrons in a neutral atom). As we progress along a row from left to right, and between rows from top to bottom, the atomic numbers increase in order. The periodic table consists of a series of boxes or blocks with essential information about the element: the atomic number, the chemical symbol, and either the mass number or atomic mass. When writing an element with its atomic number or mass number outside of the periodic table, however, scientists put the atomic number as a subscript before the symbol and the mass number as a superscript before the symbol. For example, hydrogen (symbol H) is atomic number 1 so it is shown as: 1 H The atomic mass of hydrogen can be 1, 2, or 3, so these three isotopes (protium, deuterium, and tritium, respectively) are written: 1 H, 2H, 3H In 1869, Mendeleev discovered that known elements could be arranged in a periodic table. If one arranges elements in numerical order by atomic number, leaving blanks in some rows, elements with similar properties can be grouped into columns. Elements in one column share common chemical properties. 2-22 Elements in the same column (group) of the periodic table have similar chemical properties, just like days of the week have the same properties. Elements in the same column (group) of the periodic table have the same number of valence electrons, which is what gives them similar chemical properties. Remember chemistry is what happens when electrons interact. 2-23 These are called groups and are numbered from 1 to 18. For example, in the first column, sodium (Na) and potassium (K) have many chemical properties in common, such as which elements they like to combine with. Sometimes, the groups have names: Group 1 (H, Li, Na, K, and chemically similar elements) are called the alkali metals. Group 2 (Be, Mg, Ca, and chemically similar elements) are called the alkaline earth metals. Group 17 (F, Cl, Br, I) are the halogens. Group 18 (He, Ne, Ar, Kr, Xe, Rn) are the noble gases. Metals are on the left 2/3 of the periodic table (yellow blocks). Non-metals are on the far right side of the table (green blocks). For example, calcium (Ca) is the most common metal in the human body (atomic number 20, 1.5% of body mass) while oxygen is the most common non-metal (atomic number 8, 65.0%). There are elements that share metal and non-metal properties and are called metalloids (lilac blocks). The only metalloid atoms found in trace quantities in the human body are boron (B) and silicon (Si). The metalloids form a dividing line between the metals below and to the left and the non-metals above and to the right. 2-24 The far-right column (group 18) Hydrogen (H) Silicon (Si) contains atoms that do not combine Helium (He) Phosphorus (P) with other atoms. For this reason, they Lithium (Li) Sulfur (S) are called noble gases. (Remember Beryllium (Be) Chlorine (Cl) that oxygen combines with itself to Boron (B) Argon (Ar) form O2 molecules, oxygen gas. Argon Carbon (C) Potassium (K) is 1% of the atmosphere and is the Nitrogen (N) Calcium (Ca) most common of the noble gases, Oxygen (O) Iron (Fe) but it won’t even combine with itself, Fluorine (F) Selenium (Se) so argon gas is simply Ar.) Because Neon (Ne) Bromine (Br) the noble gases don’t combine with Sodium (Na) Krypton (Kr) anything, they are not present in the Magnesium (Mg) Iodine (I) human body. Still, they have their Aluminum (Al) Xenon (Xe) uses; we will revisit the noble gases in Objective 6. For this course, you need to know the names and symbols (but not the atomic numbers or masses) for the elements shown at the right. Note that in most cases, the symbols are the first, first and second, or first and third letters in the element name. The exceptions are sodium (Na), potassium (K), and iron (Fe). This is because in the Renaissance when these elements were named, they were given Latin names (natrium, kalium, and ferrum, respectively). A blood condition is called by the suffix –emia. Hypernatremia is too much sodium in the blood, while hyponatremia is too little. Hyperkalemia is too much potassium in the blood, while hypokalemia is too little. Now that we know the names of the elements, let’s turn to their arrangement in the periodic table. The arrangement depends on their chemical properties, so that elements with similar chemical properties are stacked on top of each other in groups. How do we describe the chemical properties of an element? If we learn to interpret the orbital diagrams that were introduced in Objective 4, then we can use them to figure out chemical properties. An example is potassium. 2-25 The right figure shows all the available orbitals. But for the first 20 elements, the ones we’re focused on, we only need the 1s, 2s, 2p, 3s, 3p, and 4s orbitals. Those are the ones we’ll focus on in order to simplify these diagrams. We show those in the below figure. Another way of representing orbitals is shown here. Notice the diagram with colored circles and this shorthand notation have the same basic information, but we feel it’s easier for beginners to understand the diagrams so we’ll tend to use these going forward in this Unit. Atoms collide with the right orientation and velocity. Upon colliding, they either break or form chemical bonds. Chemical reactions, the breaking and forming of chemical bonds, are the basis for all biological systems. Remember that only the electrons are involved in chemical reactions. Atoms form chemical bonds by giving, taking, or sharing electrons. Remember from Objective 4 that atoms are most stable when their outer shells are filled. As an example, take the noble gas krypton (18Kr). Because it’s a noble gas, we know it has an optimal number of electrons (i.e. the electron orbitals are arranged in the lowest possible energy state). A chlorine atom (17Cl) can achieve a stable 18 electrons by adding one to become Cl–. A potassium (19K) atom can reach 18 electrons by giving one away to become K+. 2-26 Another way that atoms can achieve filled shells, and the lower energy state that results, is by sharing electrons. The sharing of electrons produces very stable interactions between atoms, what is called a covalent chemical bond. For example, one carbon atom (C4+) shares 4 electrons with two oxygen (O2–) atoms to form carbon dioxide, CO2. We will discuss sharing or trading of electrons to form chemical bonds in the Energy & the Chemical Bond objective. 2-27 The number shown here in superscript following the element symbol is referred to as the valence of the atom. In general, the valence is the same as the number of electrons that will be taken to form anions (expressed as a negative number) or the number of electrons which will be given away to form cations (expressed as a positive number). The uppermost orbital, in terms of energy, is called the valence orbital and participates in chemical reactions. The bond formed between two or more atoms that share electrons is called a covalent bond. Let’s look at the elements with 3, 4, 5, or 6 electrons in the neutral atom: 3Li, 4Be, 5B, and 6C. Each of these can reach the stable electron configuration represented by 2He by shedding electrons: 3Li wants to shed 1 electron, 4Be wants to shed 2 electrons, 5B wants to shed 3 electrons, and 6C wants to shed 4 electrons (but, importantly, it can also pick up 4, as we’ll see later). The position of carbon in the orbital diagrams, and therefore in the periodic table, is what makes this element the basis for life on Earth. 2-28 11 Na, 12Mg, and 13Al also want to shed electrons to achieve the stable configuration represented by 10Ne. This is shown in the electron orbital diagrams and Bohr atoms represented here. 2-29 Now you should be comfortable with the concept of shedding electrons to reach stable configurations of electrons. Can an element gain electrons to achieve the same goal? Sure. Taking 10Ne as the ideal again, we see how 6C, 7N, 8O, and 9F gain 4, 3, 2, or 1 electron (respectively) to achieve the stable 10Ne configuration, with the 1s, 2s, and 2p orbitals completely filled with electron spin pairs. 2-30 Periodicity in the Periodic Table: Valence This diagram shows the relationship between sodium’s valence, its electron configuration in the shell model, and its electron configuration in the orbital model. Sodium is atomic number 11. That means there are 11 protons in the nucleus. This number cannot change by chemical means. What can change is the number of electrons. In the neutral sodium atom, there are 11 protons (packets of positive charge) and 11 electrons (packets of negative charge). +11 –11 = 0. In the sodium ion, sodium reaches its low-energy configuration by giving away one electron. That lone electron is in the outermost shell in the shell model (top) and in the 3s orbital in the orbital model (bottom). Sodium ions still have those 11 protons (+11) but now have only 10 electrons (–10): +11 –10 = +1. We say the valence of sodium is +1. 2-31  We’ve already seen how different groups in the periodic table want to take, or give away, electrons to reach a stable configuration. For example, hydrogen (H), lithium (Li), sodium (Na), and potassium (K) in Group 1, at the far left of the periodic table, all want to give away one electron to reach a stable configuration. Compare their orbital diagrams: each has one “extra” electron to give away. 2-32 Near the right edge of the periodic table, in Group 17, we find the halogens, which are highly chemically reactive. This is because they like to steal electrons to complete their top orbitals and reach a low energy state. Note how, in each case, there is one “open parking spot” where a single electron (arrow) can be added to complete the orbital and achieve that low energy state. 2-33 Magnesium (Mg) and calcium (Ca) are important elements in the human body, regulating a wide variety of cellular processes. Calcium is also an important component of bones (Unit 9). They are both in group 2, and they both have similar chemistry. How many electrons do they want to give away? Magnesium has 12 protons (+12) and when ionized has 10 electrons (–10), so it has a valence of +12 –10 = +2. Calcium has 20 protons (+20) and when ionized has 18 electrons (–18) so it also has a valence of +20 –18 = +2. 2-34 The Canvas course includes a worksheet and formative exercises that allow you to see these patterns more clearly. The tendency of Group 1 elements to reach a low energy state when there is one more proton than electron is indicated as +1. The tendency of Group 17 elements to have one fewer proton, since they take an extra electron, is indicated as –1. Remember that electrons have a negative charge, so for example, fluorine has 9 packets of positive charge (+9) and when it is ionized, has 10 packets of negative charge (–10) and so overall has a charge of +9 –10 = –1. Sodium has 11 packets of positive charge (+11) and when it is ionized, also has 10 packets of negative charge (–10) and so overall has a charge of +11–10 = +1. The red numbers at the top of each column (group) indicate the most common valence, and therefore the chemical properties of that group. 2-35 It’s relatively simple to figure out how the elements in Groups 1-2, and 13-17 are going to act, and that’s why those have big red numbers at the top. (Carbon, in group 14, is not labeled but it either gives away 4 electrons (C4+) or takes 4 electrons (C4–) since it is between the +3 and –3 columns of the periodic table. The ability of carbon to either take or give away electrons is the basis for the chemistry of living things, organic chemistry, which we will discuss in Unit 3 and for the rest of the course. In the middle of the periodic table, we find the transition elements, where the filling of electron shells becomes quite complicated. As you can see by squinting and looking at the valence numbers for each block, these elements have valences that are all over the place and don’t follow any regular pattern. The only one we’re concerned with is iron (26Fe) because it likes to hold onto oxygen. Iron exists in a +2 and +3 form (ferrous and ferric, respectively) and the +2 (ferrous) form is found caged in a molecule called heme which acts as an oxygen carrier in the body. 2-36 Periodicity in the Periodic Table: Electronegativity Electronegativity is another chemical property that helps us sort elements in the periodic table. If we remove group 18, the noble gases, because they don’t participate in chemistry, then electronegativity increases as we move towards the higher group numbers (17, on the right) and upper periods (rows). Thus, fluorine (9F), the lightest element in group 17, has the highest electronegativity of any element. At the other end of the periodic table, potassium (19K) in group 1 has one of the lowest electronegativities of all the elements. Electronegativity is the power of an atom when in a molecule to attract electrons to itself. — Linus Pauling (Nobel Prize in Chemistry, 1954) 2-37 In orbital terms, this means there is a low-energy orbital which is almost, but not quite, full. This is why group 16 and 17 elements have the highest electronegativities. Note we are not concerned with exact numbers, nor do you have to learn them. Different sources give different numbers. Rather, we want you to know the most electronegative elements (F > O > Cl > N > Br > S) because their electronegativity gives them a unique chemistry that we’ll focus on in later objectives. 2-38  Objective 6: Noble Gases in the Periodic Table 6 Summarize the properties of noble gases and describe the orbital configurations that result in a noble gas. The far-right column of the periodic table contains elements that don’t participate in chemistry at all. What configuration of electrons results in this lack of chemical properties? Noble gases, as described earlier, don’t participate in chemical reactions—not even with themselves. For example, argon is 1% of the atmosphere but we are unaware of it when we breathe because it doesn’t do anything except take up space. The chemical properties of an element depend on its electron orbital configuration. Filled orbitals represent a low-energy configuration, and to displace any electrons requires an input of energy. This is why the highly electronegative group 17 elements (F, Cl, Br, I) are so chemically reactive: they only lack one electron to achieve a filled-shell, low-energy state. 2-39 Noble gases are already in that low-energy state. For example, helium (2He) has only one orbital, the 1s. The 1s orbital can hold two electrons, and that’s how many a helium atom has: 2 protons, a variable number of neutrons (usually 2) and 2 electrons. Neon (10Ne) has 2 electrons in the 1s orbital, 2 in the 2s orbital, and 6 in the 2p orbital. Thus, both level 1 (lowest energy) and level 2 (next-lowest) are filled. (Even though you don’t need to know this way of representing orbitals, chemists sometimes write this configuration as 1s22s22p6.) Argon (18Ar) has 2 electrons in the 1s orbital, 2 in the 2s orbital, 6 in the 2p orbital, 2 in the 3s orbital, and 6 in the 3p orbital. Thus, its first-, second-, and third-level orbitals are all completely filled (1s22s22p63s23p6). Krypton (36Kr), xenon (54Xe), and radon (86Rn) all have filled shells as well. All are in group 18 with He, Ne, and Ar, and all are noble gases. Radon is a radioactive gas which does not participate in chemical reactions. When uranium and radium decay, they emit alpha particles. Radon gas can accumulate in houses and expose the residents to radioactivity in the form of α radiation, which increases the incidence of cancer, especially lung cancer. 2-40  Objective 7: Atomic Number & Atomic Mass in the Periodic Table 7 Given a block of the periodic table, identify the atomic number and describe its significance. Identify the atomic mass and describe how it is derived. In this objective, students gain experience in figuring out how the atomic mass found in the periodic table is derived. The atomic number is the number of protons. Remember that in a neutral atom, the number of protons and electrons is equal, but the electrons come and go, so the proton number is a reliable indicator of the chemical properties of that element. Mass number is the sum of protons and neutrons. Electrons are so light that they do not contribute to the mass number and are ignored. The atomic mass (or atomic weight) is the average mass of all naturally occurring combinations of protons and neutrons in the nucleus. That is, the atomic mass depends on what isotopes of the element are found in nature. 2-41  If we use an atomic butterfly net (not yet invented, but we can imagine it) and capture 100 atoms of hydrogen, we find that one out of 100 H atoms in nature is hydrogen 2, deuterium (2H). This is how we calculate the atomic mass (or atomic weight) that is shown in the periodic table. If 99 atoms are 1H, with a mass of 1 Da each, and 1 atom is 2 H, with a mass of 2 Da, then the total mass of the 100 hydrogen atoms is 101 (99+2) and the average mass is 1.01 Da. If the math freaks you out, notice there are 100 hydrogen atoms in the figure above. Count the number of large particles (red protons and purple neutrons) and divide by 100 to get the average mass. Carbon exists as 12C and 13C. There are about 99 12Cs for every one 13C, so its atomic mass is 12.01. Nitrogen exists as 14N or 15N, and there are about 99 14Ns for every 15N. That gives an atomic mass of 14.01. 2-42 The most common isotopes of iron are 54Fe, 56Fe, 57Fe and 58Fe. What do the numbers 54, 56, 57, and 58 represent? How does that compare to 26, the atomic number for iron? How many neutrons are present in each case? (Hint: 54 massive particles are found in the nucleus of 54Fe, and the number of protons is always 26, so the number of neutrons is 54 – 26 = 28.) Wikipedia says that the proportions of iron isotopes are 5.8% of 54Fe, 91.7% of 56Fe, 2.1% of 57Fe and 0.3% of 58Fe. Given this information, you can check and see if it agrees with the block of the periodic table. Since the most common isotope by far is 56Fe, we expect to get a number close to 56, and we do. By rounding the atomic mass, we can easily determine the most common isotope of the element, with a couple of exceptions that you will not be tested on. The atomic mass of hydrogen is 1.01, and the most common isotope of hydrogen is 1H, as we’ve already explained. The atomic mass of carbon is 12.01, and the most common isotope of carbon is 12C. The atomic mass of nitrogen is 14.01, and the most common isotope of nitrogen is 14 N. The atomic mass of iron is 55.85, and the most common isotope of iron is 56Fe. 2-43  Objective 8: States of Matter 8 Name the three states of matter found in the human body, and describe characteristics associated with each. Describe the relationship of molecules to one another in a solid, liquid, or gas. Our goal here is to understand how the submicroscopic arrangement of molecules is related to the macroscopic states of matter: gas, liquid, or solid. We will also develop a relationship between temperature and states of matter. All biological materials exist in one of three states: gas, liquid, or solid. Gases have widely spaced particles that are flying around, banging off of each other. The particles don’t interact much beyond these collisions. This gives gases their special properties we’ll discuss later. Gases have an indefinite shape and indefinite volume. That is, gases are very compressible. Liquids contain particles that are interacting with one another, but much more weakly. A liquid changes its shape to match the container it’s in, so it’s said to have an indefinite shape. However, the volume is definite, and liquids are not very compressible. At the submicroscopic level, liquids contain particles that are colliding with each other but even during those collisions they maintain some sort of interaction with each other. 2-44  Solids are substances in which the particles are tightly associated with each other. Solids have a definite shape and definite volume and are not very compressible (which is why we sit on solid substances and not other kinds). The same molecules make up solids, liquids, and gases. They change states depending on the relationship between the molecules. That relationship depends on temperature and pressure. As pressure increases, the relationship between molecules becomes closer and closer until they are in a rigid array. This conception of how molecules are related to each other in the three forms of matter is formally called the kinetic molecular theory. For water, this rigid array the solid form, is called ice. In solid ice, water molecules are vibrating, but they do not move past each other. That means that ice maintains its shape and you can stand on it, for example when ice fishing. Temperature is a way of describing the motion of particles. The higher the temperature, the more molecular motion that takes place. The more molecular motion that takes place, the higher the temperature. The lowest possible temperature in the universe, called absolute zero, is the temperature at which there is no molecular motion at all. Since there can’t be any less motion than nothing, absolute zero is truly absolute zero. This happens at –273.15°C which is also described as 0 Kelvin or 0 K. The Kelvin scale is the same as the Celsius scale, it just starts at –273.15°. Thus, the surface temperature of the human body is 37°C or 310 K. You will want to spend some time practicing how to interpret this graph. 2-45 Follow the dotted line which indicates atmospheric pressure (101 kPa) on the phase diagram. At 0°C, water transitions from the solid form to the liquid form. The molecules of water are more loosely related to each other, forming temporary bonds that break and reform easily. Because the molecules can slide past each other, they no longer adopt a defined shape but fill the container they’re put into. At 100°C, water transitions from the liquid form to the gaseous form (water vapor). Now the molecules of water are bouncing off each other like little billiard balls on speed. There is no relationship between the water molecules except for when they collide with each other (about 10 billion times per second!). Understanding the relationship of molecules of O2, N2, Ar, and CO2 in the atmosphere is essential for our understanding of how the respiratory system works, which we’ll address in Unit 17. Understanding that billions of molecular collisions per second occur even in liquids is fundamental for an understanding of our next concept, diffusion. Objective 9: Diffusion 9 Explain how molecules diffuse. Illustrate diffusion with examples from biological systems. Discuss how diffusion is related to increasing entropy (the Second Law of Thermodynamics). Because molecules move, and collide with other molecules, they distribute themselves throughout a container or closed space. The formal term for this property of molecules is diffusion. As molecules collide, they distribute themselves more randomly (i.e. with increasing disorder). The formal term for this disorder is entropy. Diffusion and entropy are related concepts. As particles move, they bounce around and distribute themselves all over. This is one aspect of a property called entropy: things become more disorganized in our houses. If we want to decrease entropy (i.e., increase organization) we have to do work—we have to expend energy to clean the house. The idea that disorder or randomness (formally, entropy) is increasing is called The Second Law of Thermodynamics. 2-46 This means that molecules in a liquid or gas, which are bouncing off each other, will eventually distribute themselves all over the container (i.e., increase the entropy of the closed system in the beaker). When we put ink in water, it eventually moves throughout the water and does not stay in one place. This is called diffusion. If there is no barrier, substances (like ink) always move from where they are at high concentration to where they are at low concentration. Given time, socks in your house will go from where they’re at high concentration (the sock drawer) to where they are at lower concentration (everywhere else). Understanding diffusion is essential for much of Unit 4, so take time now to make sure you are familiar with the concept. To take one specific example, diffusion moves molecules of the hormones 17β-estradiol (one of the estrogens) and dihydrotestosterone (one of the androgens) from the blood into the cells of the body, where these hormones can exert their effects on the DNA of the cell, permanently altering the cell’s fate. This is a large part of understanding the endocrine system in Unit 14. 2-47 A final principle of thermodynamics that we will use many times in the course is that energy is always transferred from where it is high to where it is low. That is, we can heat a pot of water by putting it on the stove, but we can’t cool it down by putting it on a “reverse stove”. Similarly, there’s no “reverse microwave” to cool down the baby formula you accidentally overheated. You can increase the number of collisions between the molecules that make up milk, but you can’t directly slow them down. (If we put the baby bottle in a refrigerator or ice bath, the heat is transferred from the hot milk through the walls of the container into the cooler material. Similarly, a refrigerator works by heating your house, transferring heat from the inside of the refrigerator to the air in the kitchen. An air conditioner then cools your house by heating up the atmosphere around your house.) 2-48  Objective 10: Energy & the Chemical Bond 10 Illustrate how energy is stored in, and released from, chemical bonds. Explain the relationship of electrons in covalent bonds, polar covalent bonds, ionic bonds, and hydrogen bonds. Classify the relationship between the energy in these bonds and the strength of the bond. Chemical properties, as we saw earlier, are the result of the arrangement of electrons into energy levels. Now we will consider how chemical bonds hold onto that energy, and how it converts between two forms: the energy of motion (temperature, kinetic energy) and the energy which is present but not (yet) visible (bond energy, potential energy). Bond energy can result from the sharing of electrons, the partial sharing of electrons, or the giving or taking of electrons. Each of these situations results in a different type of bond which we will discuss in this objective as well. Energy comes in two forms: potential and kinetic. Potential energy, like a coiled spring, is stored energy that is not (yet) able to do work. Kinetic energy is the energy of movement. When you are riding in a ski lift to the top of a mountain, you are moving: this is kinetic energy. When you pause at the top and look down, you have turned the kinetic energy of the ski lift into potential energy. When you begin to slide down the hill, that potential energy is turned back into kinetic energy. Energy is conserved. It probably won’t surprise you to know that this concept is called The Law of Conservation of Energy. This Law says chemical energy can change forms (from kinetic to potential, and back again) but it cannot be gained or lost. (You may have heard of Einstein’s famous equation: E = mc2. The reason this was revolutionary is that it violates the Law of Conservation of Energy. Nuclear energy, as contrasted with chemical energy, can be converted to matter and back into energy again. Hence the nuclear bomb.) Energy in chemical bonds is potential energy. When it is released to do work (like lifting a dumbbell) it is converted to kinetic energy. Since this is what cells of your body spend most of their time doing, we will study the general rules for converting food to cellular chemical bonds and converting cellular chemical bonds to work in Units 3 and 5, and then amplify on those concepts throughout the course. 2-49 Chemical reactions that convert kinetic energy to potential energy of chemical bonds are called endergonic which is roughly equal to endothermic (endo–, inner; –therm, temperature). The potential energy is hidden within these chemical bonds (i.e. we say the reaction absorbs energy). In biology, endergonic or endothermic reactions are roughly equal to anabolic reactions, our name for chemical reactions which form chemical bonds. Chemical reactions which release the potential energy of molecules from their bondage are called exergonic or exothermic. Remember that the kinetic molecular theory says increasing temperature increases the kinetic energy of molecules, so this type of chemical reaction converts potential energy to kinetic energy. Our lab director, Pamela, and the male firefly would like me to remind you here that in rare circumstances, the excess energy from a chemical reaction is given off as visible photons rather than as heat. This happens in Pamela's glowstick demo and in bioluminescent ("living light-emitting") animals such as the firefly pictured here. In biology, exergonic or exothermic reactions are roughly equal to catabolic reactions, our name for chemical reactions which break chemical bonds. In most of the chemical reactions we’ll study, the energy is turned into work. For example, the breaking of chemical bonds allows muscles to contract, and a weightlifter can raise a barbell over his head in a “clean and jerk” maneuver. Notice that the Greek word εργου (pronounced “ergon”) is work, which is where exergonic comes from. When iron particles rust, they give off heat in an exothermic reaction, which is why powdered iron is used in handwarmers, and why the handwarmers last longer if they are inside a glove (where there is less oxygen) than out on a table. 2-50 The three Laws of Thermodynamics tell us that no reaction is 100% efficient. In a biological system, the chemical bond energy that isn’t turned into work is turned into heat. Heat production is a side-effect which has been harnessed by our body. For example, if we are cold, we contract our muscles without doing work; this produces waste heat, which helps us to stay warm. That’s why we shiver when we’re cold. It’s also why heat management becomes a problem if we do work in hot weather. In that case, we may have a harder time shedding the waste heat—we might overheat. You’re probably sick of me reminding you of the Tom Cech quote: chemistry happens because molecules and atoms collide together with the right orientation and velocity. But we need to use that idea again. Remember that in the kinetic molecular theory, molecules and atoms are colliding. Sometimes, their collisions have the right orientation and have just enough energy (velocity) that their electrons have no choice to interact, and sometimes, rearrange themselves. These electron rearrangements are where chemical bonds are formed: where we convert kinetic energy (velocity) to potential energy (a In an endothermic reaction, chemical bond). energy must be used to build a chemical bond. Notice the energy of the products is higher than the energy of the reactants. Endothermic reactions are storing energy in the chemical bonds that are built. In an exothermic reaction energy is released as a bond is broken. Notice that activation energy is still needed to break the bond. As energy is released, the energy of the products is lower than the energy of the reactants. 2-51 Either the orientation of molecules or their velocity requires more energy than will eventually be represented in the resulting chemical bond. This energy, which you can think of as a kind of “hill” the electrons in the molecules need to “climb”, is called the activation energy. We will see ways of lowering the activation energy in Unit 3. We can lower the activation energy (i.e. by making sure the molecules are correctly oriented) but we can’t change the potential energy in the chemical bonds of the reactants, nor can we change the potential energy in the chemical bonds of the products. There are different kinds of covalent bonds, with different strengths. Covalent bonds result when atoms share electrons. For example, when two hydrogen atoms share both of the electrons they hold in common, it’s called a covalent single bond. Carbon (6C) always makes four bonds. It can complete its shell by releasing four electrons (to be like He), or by capturing four electrons (to be like Ne): 2He (6 – 4 = 2) or 10 Ne (6 + 4 = 10) are the nearest noble gases with completed shells. For example, if C shares four electrons with four H atoms, four covalent single bonds are formed, as in methane (CH4). Equal sharing of electrons, in quantum terms, means there’s an equal probability of finding the electron near the carbon nucleus or near the hydrogen nucleus. If H shares its single electron with another H, a covalent single bond is formed as we saw previously. A single bond is usually indicated by a single line: H—H A covalent double bond is when four electrons are shared equally. For example, each O atom wants to gain two electrons to complete its shell; when each oxygen atom shares two electrons, for a total of four electrons shared between two atoms, it can achieve the same goal. A covalent double bond is symbolized by a double dash: O=O O=C=O A covalent triple bond is when six electrons are shared equally. For example, each nitrogen atom wants to gain three electrons to complete its shell; when each nitrogen atom shares three electrons, for a total of six electrons shared between two atoms, they can achieve the same goal. A triple bond is symbolized by three lines: N≡N 2-52 In polar covalent bonds like those in H2O, the oxygen nucleus is more highly electronegative so it has greater electron pulling ability. In quantum terms, there is a higher probability of finding the electron near the oxygen nucleus than near the hydrogen nucleus. The difference in probability is observed as a partial negative (higher probability of finding a negatively-charged electron) or partial positive charge (lower probability of finding a negatively-charged electron) in different parts of the molecule. In science, the lower-case Greek letter delta (δ) is used to mean “partial” so a partial positive charge is symbolized δ+ and a partial negative charge is symbolized δ–. 2-53 In HCl, there is a greater chance of finding the electron near the highly electronegative Cl nucleus so the Cl end of the molecule is labeled δ– while there is a smaller chance of finding the electron near the H end of the molecule so that end is labeled δ+. Ionic bonds result from the complete removal or addition of electrons to an atom. For example, there is a difference of 2.2 units between Na and Cl. The highly electronegative Cl atom just straight-up takes the electron; now 17Cl has 17 protons and 18 electrons for a net charge of –1 (symbolized Cl–). The weakly electronegative Na atom cannot hold onto its valence electron (nor does it want to, since it can complete an energy level by giving it away) so 11Na now has 11 protons and 10 electrons for a net charge of +1 (symbolized Na+). 2-54 In a crystal of Na+Cl–, there is an opposite-charge attraction between the Na+ and Cl– ions, and it’s this charge attraction that holds the ions together. In other words, they are joined by an ionic bond. Hydrogen bonds result when a hydrogen atom gets into a relationship with oxygen, nitrogen, or sulfur, each of which results in an electronegativity difference of 0.5 to 1.7, the range for formation of polar covalent bonds within the molecule. Remember that polar covalent bonds result in a δ– near the more electronegative molecule (O, N, or S) and a δ+ near the H. These partial positive and negative charges attract each other like ionic bonds, but in a weaker sort of way. Of the bonds we’ve discussed, covalent bonds are the strongest. Of covalent bonds, triple are stronger than double, and double covalent bonds are stronger than single covalent bonds. Ionic bonds are the next strongest type. Hydrogen bonds are the weakest type of bonds that we’ve discussed. Summarizing bond strength: triple covalent > double covalent > single (polar or non-polar) covalent > ionic > hydrogen 2-55 2-56  Objective 11: Special Properties of Water 11 Describe the special properties of water: hydrogen bonding and surface tension. Water forms from a relationship between the highly electronegative atom oxygen and the more weakly electronegative and really kinda wimpy hydrogen atom. Since hydrogen only has one proton, it has a hard time holding onto its electrons. This results in the special properties of water. Water is 60% of the human body for a reason: these special properties make water the ideal substrate for chemical reactions. Water has special hydrogen bonding properties. Water is thought of as a near-universal solvent. A wide variety of substances dissolve in water. 2-57 A property called surface tension results from the hydrogen bonds between water. Water molecules are strongly attracted to each other and seem to "pull" together like an elastic sheet. Surface tension can cause delicate baby lungs to collapse if we don’t administer a surfactant—a biological soap that breaks hydrogen bonds. Surface tension also leads to capillary action, when blood seems to “leap” into a narrow hematocrit tube in defiance of the Law of Gravity. Surfactants disrupt the hydrogen bonds between water molecules and thereby prevent capillary action. Most substances are denser as solids than liquids. Water is different. Water is less dense as a solid than a liquid—ice floats. Frozen water does not allow its molecules to rotate and so it does not “pack” as tightly as liquid water where the water molecules can rotate and find a close-packing configuration. Ionic bonds are easily disrupted by water, which of course has charges of its own in its polar covalent bonds. This is why salt (NaCl) easily dissolves in water; as it does so, the Na+ ion surrounds itself with water molecules and the δ– O end of water orients towards the positively charged Na+. Similarly, Cl– attracts the δ+ H molecules in water and Cl– is surrounded by a shell of water with its molecules oriented the other way. These oriented water molecules surrounding ions in solution are called hydration shells. The hydration shells of Na+ and Cl– are discussed and shown in Objective 13. When water (H2O, which is also written HOH) is split to form H+ and OH– ions, the resulting charged ions contribute to pH, a property that we will discuss next. 2-58  Objective 12: Acids, Bases & the Concept of pH 12 Define: acid and base. Describe how acids donate hydrogen ions. Explain the formation of hydronium ions. Describe how bases accept hydrogen ions. Summarize what makes an acid or base strong or weak. Describe the concept of pH. Given a pH value, be able to state whether it is acidic, neutral, or basic. Understand the reasons why a biological fluid might be acidic or basic. Discuss how buffers moderate changes in pH. One consequence of hydrogen’s weakness in holding onto electrons is that sometimes it just gives up and wanders off by itself as a free proton. That sort of compound is called an acid. Sometimes molecules attract those free protons. Those molecules are called bases or alkalis. The tendency of a molecule to donate or accept H+ is formally called the pH, a scale that runs from 0 (strongest acid) to 14 (strongest alkali). Pure water is neither acid nor base (i.e. is neutral) and has a pH exactly halfway between 0 and 14, at 7.00. Substances which “sponge up” hydrogen ions and reduce the changes in pH which would otherwise result are called buffers, and we’ll see how that “sponging” occurs. Substances like HCl share characteristics of polar covalent and ionic compounds. There is a significant probability, approaching 100% or 1.0, that H+Cl– will split apart into H+ and Cl– ions. Similarly, there is a near-100% probability that H2SO4 will split into H+ and HSO4– and that HNO3 will split into H+ and NO3–. Therefore, all these compounds are called acids. You may already know these acids by name: ꞏ HCl is hydrochloric acid (used to clean concrete or add chlorine ions to swimming pools) ꞏ H2SO4 is sulfuric acid (used in lead-acid car batteries) ꞏ HNO3 is nitric acid A chemist named Lewis then decided that defining an acid was nice and easy: according to Lewis, acids are hydrogen ion donors. 2-59  In water (aqueous) solutions, virtually all the hydrogen ions react with water to form hydronium. We will use "hydrogen ion concentration" to mean "hydronium ion concentration" throughout this course. pH is a measure of the hydrogen ion concentration in a solution. But, as Tina Turner said, “we never do anything nice and easy”. Also, Lewis’ definition leaves a lot to be desired. Therefore, two chemists named Brønsted and Lowry tried a different definition, and that’s the one that is more useful in biochemistry. The Brønsted-Lowry definition of an acid starts with the donation of a H+ but doesn’t end there. H+ then reacts with H2O to form hydronium ion (H3O+): H+ + H2O → H3O+ In the Brønsted-Lowry conception of bases, a base is a hydrogen ion acceptor. Bases have a full or partial extra electron which attracts and holds a positively-charged hydrogen ion (proton). Bases are also called alkalis and basic solutions are called alkaline. 2-60 A twist in the acid-base story is that acids and bases can be strong or weak. You know this already. Only a cruel person would offer to put battery acid (H2SO4) on your tongue, but a chef might offer to put citric acid (C6H8O7) on your tongue. (Citric acid is the acid that gives citrus fruits like lemons their tang.) H2SO4 is a strong acid. Citric acid is a weak acid. Strong acids donate almost all their hydrogen ions. Weak acids donate only a small proportion of their hydrogen ions. Similarly, you might drink baking soda (NaHCO3) in water to settle a sour stomach, but you probably wouldn’t want to drink the drain cleaner lye (NaOH) to settle your stomach. Severe tissue damage would result if you did. NaHCO3 is a weak base. NaOH is a strong base. Strong bases accept almost all hydrogen ions in the solution. Weak bases accept only a small proportion of the available hydrogen ions. So if acids are hydrogen ion donors, what makes them strong or weak? Strong acids like HCl completely fall apart in water to donate all their hydrogen ions, while weak acids like citric acid only donate 1 H+ for every 1408 molecules of citric acid. Similarly, strong bases like NaOH accept all the hydrogen ions they can find, while only 1 in 4700 NaHCO3 molecules accept a hydrogen ion. 2-61 You do not need to memorize this table. But notice the relative strength of some common acids: sulfuric acid, hydrogen chloride (stomach acid), phosphoric acid, acetic acid. Also notice the relative strength of some common bases: hydroxide ion, carbonate ion, ammonia, bicarbonate ion (hydrogen carbonate ion). 2-62 We use a quantity called the pH to define how strong or weak an acid or base is. The pH scale runs from 0 to 14. The pH is formally defined as “the negative logarithm of the hydrogen ion concentration” but for our purposes, we don’t need to learn the formal definition. We do need to know which pH values represent acids and bases, and how the scale works. Do not memorize these values; rather use them to understand the pH scale. Strong acids completely dissociate in water to release all their hydrogen ions. Solutions of strong acids have a pH near zero. 2-63 Weak acids only partly dissociate in water. Lime juice, a solution of citric acid, has a pH of 2. Orange juice is even less acidic than lime juice, with a pH of about 4. About 1 in 100,000,000,000,000 water molecules spontaneously break apart into H+ and OH– ions. If you care to do the math (we assume you don’t) and apply the definition of pH, that means that pure water has a pH of 7. A pH of 7 is called neutral and it is the only value that is neither acid nor base. Blood is a weak base. It has a pH of 7.4, just slightly above 7. Baking soda in water has a pH of 8.3, more basic than blood. Bleach (sodium hypochlorite) has a pH of 12.6, which is quite basic, and lye (NaOH) is the most basic substance we commonly encounter, with a pH of 14. That means all the NaOH molecules in water break apart into Na+ and OH– ions. Weak acids are used in the body to kill bacteria, viruses, and fungi. Stomach acid (pH about 2-3) not only kills bacteria but helps break down swallowed food into small molecules, because hydrogen ions participate in a lot of chemical reactions that our digestive system finds useful. Vaginal fluid (pH 3.8-4.5) is mildly acidic, which helps maintain the health of this organ. 2-64 Canned tomatoes must have a pH below 6, according to USDA home canning guidelines; this helps inhibit the growth of deadly Clostridium botulinum, the bacterium that causes botulism. It’s prudent to test the pH of your home-canned tomatoes with pH papers. In the lab, it’s more common to use a pH meter to measure the pH of a solution. When we place a strong acid such as H+Cl– into water, the H+Cl– dissociates into H+ and Cl–. The resulting H+ then reacts with water (H2O) to form hydronium ion (H3O+), meeting our definition of acid. 2-65 When we place a strong base such as the ionic compound Na+OH– into water, it dissociates into Na+ and OH– ions. The OH– ions act as H+ acceptors; the reaction between H+ and OH– forms water: H+ + OH– → HOH (H2O) Thus, OH– meets our definition of base. Weak acids and weak bases are important components of a special type of solution called a buffer. A buffer is a substance which acts as a “sponge” to hold and inactivate H+ and/or OH–. In this way, it stabilizes the concentration of H+ and/or OH–, which in turn stabilizes the pH. This figure shows the creation of a buffer by mixing a weak acid and a weak base. 2-66  Now when we add acid to this buffer mixture, the H+ ions associate with the weak base anion (A–) to form the weak acid HA. This keeps the H+ ions from reacting with water to make hydronium (H3O+). The less hydronium we make, the less the pH changes. And when we add base to this buffer mixture, the weak acid HA donates its H+ to the OH– ions we added, making water and the weak base A–. The more hydrogen acceptor OH– we remove in this way, the less the pH changes. The most important biological buffer is the system of carbon dioxide, carbonic acid, and bicarbonate (CO2/H2CO3/HCO3–). Note that when CO2 is dissolved in water (H2O), carbonic acid is formed: H2O + CO2 ↔ H2CO3 ↔ H+ + HCO3– 2-67 Each of the double-headed arrows is a step where H+ can be donated or accepted. Note that all parts of this complicated chemical equation have 2 H, 1 C and 3 O so this is just a rearrangement of the same atoms. The pH of blood is carefully maintained by homeostatic mechanisms to lie between 7.35 and 7.45, and in healthy individuals is almost always exactly 7.40. This pH is kept constant by the carbonic acid—bicarbonate buffer system mentioned earlier. The lungs and kidneys work together to maintain the pH of blood as close to 7.40 as possible. Objective 13: Solutions 13 Compare the molecular properties of a solution, colloid, and suspension. Explain how to prepare a solution of a given molarity. Describe the property of osmolarity. We’ve already seen how ionic compounds such as salt dissolve in water. At the submicroscopic level, sodium ions and chloride ions are surrounded by water molecules and the resulting ions are too small to scatter light. The salt disappears into the spaces between water molecules. Sucrose dissolves the same way, but without breaking apart. It simply associates with water molecules. We need a simple way to describe how many of these submicroscopic particles are found hanging out in between water molecules. This is called the concentration of the solution. We also want to know the concentration of particles in water, because particles take up room and thereby reduce the number of water molecules per unit volume. This is called the osmolarity of the solution, which is when we don’t care what type of particles we’re dealing with, just how much room they occupy. 2-68 2-69 Solutions, Colloids & Suspensions A solution consists of small molecule-sized particles surrounded by water. For example, when NaCl is dissolved in water, Na+ and Cl– in the ionic crystal are each separated and surrounded by the opposite partial charges (a group of δ– O atoms in the H2O molecule pointing towards the Na+ and a bunch of δ+ H atoms in the H2O molecule pointing towards each Cl– ion). Similar hydration shells are formed around other + or – ions. These individual ions or polyatomic ions (see Objective 14) are too small to alter the path of visible light, so the solution appears clear (that is, salt water and pure water look the same). 2-70 This same process of surrounding molecules with hydration shells occurs whenever a particle (such as the purple cation here) is capable of interacting with the partial charges of water. For example, the hydration shell formed by the polyatomic ion SO42– is shown below. The yellow S atom is surrounded by four red O atoms, which in turn attract the partial positive charges of multiple water molecules. In a solution of small molecules such as sugar, the sugar or other small molecules remain intact and becomes surrounded by a hydration shell formed by hydrogen bonds. Note this can only happen if the small molecule is capable of forming hydrogen bonds (i.e. has plentiful O, N, or S). For example, table sugar (sucrose) is C12H22O11; with 11 oxygens capable of hydrogen bonding. It is very soluble in water. Oleic acid, a component of olive oil, is C18H34O2; with only 2 oxygens capable of hydrogen bonding. Because it is predominantly C and H (i.e. a hydrocarbon), it is not very soluble in water. We will look at this further in Unit 3. As for individual ions surrounded by a hydration shell, a water-soluble, small molecule surrounded by a hydration shell does not scatter light. A solution of table sugar appears clear, the same as pure water. 2-71 A colloid is formed when the molecules are big enough to scatter light. For example, the proteins in milk are large enough molecules that they can interfere with the travel of visible light. For this reason, even skim milk appears cloudy. When particles are the size of cells, and can be separated by a tabletop centrifuge, we call the resulting mixture a suspension. Blood is a suspension; as we will see in Unit 15, if you centrifuge blood the red and white blood cells separate from the colloid (large proteins) in a salt solution. The colloidal proteins in an NaCl solution are called blood plasma. Molarity: Measuring the Strength of a Solution with One Component In the emergency department, if tissue-typed, cross-matched blood or fresh-frozen plasma is not available, and we need to quickly and safely replace the volume of blood lost in trauma, we use sterile saline, which is 0.9% NaCl. That means that 0.9 g of NaCl have been dissolved in 100 g of water. (The Latin word per cent, symbolized %, literally means “out of 100”.) Because 100 g of water has a volume of 100 mL (which is the same as 100 cm3 or 100 cc), to make a 0.9% NaCl solution, someone at the factory added 0.9 g of NaCl to 100 mL of pure water. Scientists use 10.94 g sucrose in 100 mL water (10.94% sucrose) for similar purposes, as discussed below. For most applications in biochemistry and medicine, we don’t want to know the percentage of NaCl or sucrose. We want to know how many molecules of NaCl or sucrose are in 1 liter (1 L) of solution. How is this even possible? 2-72 Luckily, Amadeo Avogadro worked this out for us. He developed a relationship between the molecular weight of a substance and the number of molecules of the same substance. If we simply add together the atomic masses of Na (22.99 g) and Cl (35.45 g), we get the molecular weight of NaCl (22.99 g + 35.45 g = 58.44 g). That means if we weigh out 58.44 g of NaCl, we know exactly how many molecules of NaCl we have. That number is called a mole and is exactly 6.02 × 1023. Avogadro’s number, or 6.02 × 1023, is a number without a unit, a so-called dimensionless number. That seems strange at first until we realize we’ve come across dimensionless numbers before. For example, a dozen is dimensionless; whether we get a dozen eggs or a dozen bagels or a dozen pencils, we have 12 objects. 2-73 If we get a gross of pencils or firecrackers, we have 12 × 12 = 144 of those objects. A mole works just the same. A mole of NaCl is 58.44 g, and contains 6.02 × 1023 NaCl molecules. A mole of sucrose is 342.3 g and contains 6.02 × 1023 molecules. A mole of hydrogen gas (H2) is 2.002 g (twice the atomic weight of hydrogen) and contains 6.02 × 1023 molecules of hydrogen gas. And so forth. Now to make a solution where we know the number of molecules in a given volume, all we need do is add a defined mass of the substance (such as 58.44 g NaCl) and bring the total volume to 1 L with water. Now we have a 1 molar (abbreviated 1 M) NaCl solution. (Another, identical, unit is moles/L.) Each liter of this solution contains 6.02 × 1023 NaCl molecules. A liter of a 1 M sucrose solution contains 6.02 × 1023 sucrose molecules. It’s a lot easier to measure the volume of a molar solution than it is to count molecules. Osmolarity: Measuring the Strength of a Solution with Many Components One more twist. For blood replacement, and many other purposes, it’s not important to know what molecules are in solution, just how many there are. That’s why you can replace blood with 0.9% NaCl (0.15 M NaCl): there are as many particles in a liter of this solution as there are in a liter of blood. When we count particles regardless of what the particles are, the quantity is called the osmolarity (Osm) of the solution. Blood plasma is 0.3 Osm/L. 0.9% NaCl is 0.3 Osm/L. 10.94% sucrose is 0.3 Osm/L. The tube on the left has a low water concentration. The tube on the right has a high water concentration. 2-74 We can think of this two ways. One way is to look at the concentration of water, which tells us which way water is moving. Since water is the only substance that can move, this is what we'll consider in the diagrams at the bottom of this image. At left, we see a situation where water concentration is higher inside the cell than outside; the green arrows show water moving out. In the center, the water concentration is equal on both sides. In this case, water still moves, but for every water molecule that moves in (yellow arrow) there is a water molecule that moves out (green arrow). Finally, at bottom right, we see a situation where the water concentration is higher outside than inside the cell. In that case, water moves into the cell, as shown by the yellow arrows. The complication arises because of the names people put on these three situations. They called them by the particle concentration, which is the opposite of the water concentration. The word "tonic" means "strength" (as in muscle tone), so the particle strength is high in a hypertonic solution and low in a hypotonic solution. When the particle concentration is high, the water concentration is low, and vice versa. So taking these left to right again, but looking at the top labels and images of red blood cells, we have: hypertonic (high particle concentration and low water concentration outside the cell, low particle concentration and high water concentration inside, water moves out of the cell) isotonic (same particle and water concentration inside and outside, no net water movement) hypotonic (low particle concentration and high water concentration outside the cell, high particle concentration and low water co