Nelson Biology 12 Quiz #1 Chapter 1 PDF

UNIT 1 Biochemistry OVERALL EXPECTATIONS analyze technological applications of enzymes in some industrial processes, and evaluate technological advances in the field of cel...

UNIT 1 Biochemistry OVERALL EXPECTATIONS analyze technological applications of enzymes in some industrial processes, and evaluate technological advances in the field of cellular biology investigate the chemical structures, functions, and chemical properties of biological molecules involved in some common cellular processes and biochemical reactions demonstrate an understanding of the structures and functions of biological molecules, and the biochemical reactions required to maintain normal cellular function BIG IDEAS Technological applications that affect biological processes and cellular functions are used in the food, pharmaceutical, and medical industries. Biological molecules and their chemical properties affect cellular processes and biochemical reactions. Biochemical compounds play important structural and functional roles in the cells of all living organisms. UNIT TASK PREVIEW In this Unit Task, you will select a molecule of interest to you. You will research its structure and function and build a model of it. The Unit Task is described in detail on page 108. As you work through the unit, look for Unit Task Bookmarks to see how the information you are learning relates to the Unit Task. 2 Unit 1 Biochemistry NEL 7923_Bio_Ch01.indd 2 3/27/12 5:11 PM  FOCUS ON STSE PROMOTING THE GOOD FATS Did you know that eating a tin of sardines may actually help you improve your marks in school? Sardines contain macromolecules, or large molecules, called omega-3 fatty acids. Omega-3 fatty acids are essential fatty acids. They are needed to maintain human health. Our cells, however, do not produce omega-3 fatty acids on their own, so we must obtain them through the foods we eat. Fatty fish, such as mackerel, salmon, and sardines, are rich sources of omega-3, as are nuts, oils, and other plant sources. In recent years, the popularity of omega-3 fatty acids has increased. Research has shown that omega-3 fatty acids play a significant role in brain function, cardiovascular health, and the produc- tion of healthy skin. In fact, the Heart and Stroke Foundation of Canada recommends eating fish such as sardines, mackerel, lake trout, herring, albacore tuna, and salmon at least twice a week to maintain optimal levels of omega-3 fatty acids. Foods such as eggs and orange juice may also contain omega-3 fatty acids. Other scientific research shows that eating a diet rich in omega-3 fatty acids reduces the risk of inflammation and may help to reduce the risk of chronic diseases, such as heart disease, cancer, and arthritis. The neurons in the human brain use omega-3 fatty acids for important cognitive and behavioural functions. Babies have a greater risk of developing vision and nerve problems if their mother lacks omega-3 fatty acids in her diet during pregnancy. Omega-3 fatty acids may also be useful for treating a number of conditions in adults, such as high cholesterol, high blood pressure, diabetes, depression, skin disorders, and asthma. People who do not consume enough omega-3 in their diet may suffer from fatigue, poor memory, dry skin, heart problems, mood swings or depres- sion, and poor circulation. It is important to obtain a balance of all nutrients, including omega-3 fatty acids, in your diet. All organisms use and rely on thousands of different types of molecules for the proper functioning of their cells. Some of these molecules can be obtained directly in the organism’s diet. Other molecules are synthesized from building blocks contained in their diet. In either case, the consumption of an appropriate mix of nutrient molecules is essential to meet the demands of cells. Questions 1. Were you already aware of the importance of omega-3 fatty acids in your diet? Are there other essential nutrients that you have heard about? In which foods can these nutrients be found? 2. The foods we eat provide us with a balance of carbohydrates, fats, and proteins. All of these nutrients are essential for a healthy body. How do you think carbohydrates, fats, and proteins compare, in terms of their chemical and physical properties? 3. Living organisms can make many “custom” molecules. For example, spiders produce silk molecules from custom proteins, while humans use another set of proteins to grow hair. What properties might these two types of proteins share? In what ways might they be different? 4. Do you think it is better to obtain omega-3 fatty acids from foods that contain them or from a supplement? Explain your reasoning. NEL Focus on STSE 3 7923_Bio_Ch01.indd 3 3/27/12 5:11 PM UNIT 1 ARE YOU READY? CONCEPTS SKILLS understand acid, base, and neutralization reactions determine the number of atoms in different molecules, explain the laws of conservation of mass and energy based on chemical formulas explain the role(s) of carbohydrates, proteins, and fats in determine the atomic mass of different compounds the body for cell membrane function balance chemical equations differentiate between diffusion and osmosis differentiate between combustion and neutralization reactions understand the role of carbon in biochemical molecules draw structural diagrams of the reactants and products in understand the role of nucleic acids a reaction draw Lewis diagrams Concepts Review 8. Electrons in a molecule will repel each other and move 1. Describe the difference between an atom and an ion. K/U as far away from each other as possible. How does this explain why a water molecule has a bent shape instead 2. What type of bond within a molecule holds the atoms of being linear (Figure 1)? T/I of the molecule together? K/U δ" 3. Compare the following: an element, a compound, and a mixture. K/U O 4. Describe the difference between an ionic bond and a covalent bond. K/U H H 5. (a) How many covalent bonds does each of the δ! δ! following type of atom typically form? (i) H (iii) C (ii) O (iv) N 104.5° (b) How is the number of covalent bonds formed Figure 1 by an atom related to its number of valence 9. (a) What is the law of conservation of mass? electrons? K/U T/I (b) How is the law of conservation of mass demonstrated 6. Match each term on the left with the most appropriate in biological systems such as food chains? K/U description on the right. K/U 10. (a) What is the law of conservation of energy? (a) pure substance (i) two or more substances (b) How is the law of conservation of energy demonstrated (b) mixture (ii) a substance made from only in biological systems such as food chains? K/U (c) solution one kind of particle 11. (a) Name the labelled cellular structures in Figure 2, (d) suspension (iii) substances in two different and describe their functions. phases that do not settle out (iv) one substance dissolved in (b) Does Figure 2 show an animal cell or a plant cell? another substance Explain your reasoning. K/U 7. (a) Does an acid or a base increase the concentration A of H1 ions in a solution? B (b) Does an acid or a base increase the concentration C of OH– ions in a solution? D (c) What is the result of a neutralization reaction E between an acid and a base? F (d) List several distinctive properties of acids and G of bases. (e) How is acidity related to the pH value of a H solution? K/U Figure 2 4 Unit 1 Biochemistry NEL  12. Examine the two equations below. How are they 17. Determine the number of atoms and the atomic mass related? T/I of each atom in each of the following molecules. Then photosynthesis: 6 H2O 1 6 CO2 S C6H12O6 1 6 O2 calculate the total atomic mass for each molecule. T/I (a) methane, CH4 cellular respiration: C6H12O6 1 6 O2 S 6 CO2 1 6 H2O (b) glucose, C6H12O6 13. Match each macromolecule on the left with the most (c) cysteine, C3H7NO2S appropriate description on the right. K/U 18. Write the balanced chemical equation for each of the (a) fats (i) directly involved with following reactions. T/I (b) proteins inheritance (ii) range from small sugar (a) CH4 1 O2 S CO2 1 H2O (c) carbohydrates molecules to large starch (b) CH2O2 S C2H2O3 1 H2O (d) nucleic acids molecules (c) C6H12O6 S C2H6O 1 CO2 (iii) responsible for long-term 19. Examine the equations below. T/I energy storage; also called (i) HCl 1 NaOH S NaCl 1 H2O lipids (ii) 2 C2H6 1 7 O2 S 4 CO2 1 6 H2O (iv) one or more folded and (a) Which equation represents a combustion reaction? coiled polypeptides; made of Explain your reasoning. amino acids (b) Which equation represents a neutralization 14. (a) Define osmosis and diffusion. reaction? Explain your reasoning. (b) Make a sketch to represent each process in (a). 20. Draw structural diagrams to represent the reactants Label your sketch. and products in the combustion reaction between (c) Explain why these processes are essential for methane, CH4, and oxygen, O2. T/I C cellular function. K/U C 21. How are cellular respiration and the combustion of 15. Draw a diagram to explain what happens to a cell in methane similar? K/U T/I each of the following types of solutions. K/U C 22. Draw the Lewis diagram for each molecule. C (a) isotonic (a) water, H2O (b) hypotonic (b) carbon dioxide, CO2 (c) hypertonic (c) ethane, C2H6 Skills Review CAREER PATHWAYS PREVIEW 16. Copy Table 1 into your notebook and complete it. (For atomic mass, round to the nearest whole number.) T/I Throughout this unit, you will see Career Links. Go to the Nelson Science website to find information about careers related to Table 1 biochemistry. On the Chapter Summary page at the end of each chapter, you will find a Career Pathways feature that describes the Name of Number Number of Atomic educational requirements for these careers, as well as some career- element Symbol of protons neutrons mass related questions for you to research. neon 20 Cl 17 28 31 52 110 NEL Are You Ready? 5 7923_Bio_Ch01.indd 5 3/27/12 5:11 PM  CHAPTER 1 The Biochemical Basis of Life What Types of Chemicals Make Up KEY CONCEPTS Living Things? After completing this chapter you will be able to Every second, millions of complex chemical reactions are taking place in your body. Your cells are continuously working and carrying out these reactions, describe the types of bonds that without any disruption in your daily activities. Your cells are mixtures of thou- are found in biological molecules sands of different chemical compounds, which are arranged into various cel- and interactions lular structures and perform a myriad of tasks. For example, more than a dozen describe the unusual properties chemical enzymes, such as DNA polymerase, DNA ligase, primase, helicase, and of water topoisomerase, are required to drive DNA replication in your cells. The process of compare the functional groups active cellular transport requires proteins, such as channel proteins, aquaporins, that contribute to the structure and and carrier proteins. Many of your cells transport chemicals that are used in other function of biological molecules parts of your body. For example, your red blood cells contain the chemical hemo- describe how the reactions globin, which is essential for transporting oxygen throughout your body. involved in biochemical Cells come in many types and have numerous functions. At their most processes can be subdivided into basic, they are tiny packages of self-replicating chemical processors. They are major classes and contribute to organized to work together and against each other to sustain life. Therefore, normal cellular function knowledge of the fundamental concepts of chemistry is necessary to under- stand how living systems function. Scientists have a deep understanding of the describe the major classes of chemical basis of life, but there are still many unsolved mysteries. biological molecules and their At the core of all cellular structures and processes are four main groups structural and functional roles of organic compounds. These compounds include a wide variety of carbo- understand that enzymes are hydrates, fats (lipids), proteins, and nucleic acids. These compounds must be involved in all biochemical either synthesized by the cell or consumed in foods, in forms that can either reactions controlled by the cell be used immediately by the cells or that can be dismantled and rearranged. As compare technologies that use and you will learn in this chapter, these four types of compounds are responsible manipulate enzymes to achieve a for thousands of functions and interactions performed by the cells. Biological desired product or process macromolecules can react with oxygen in the cells to provide chemical energy understand the potential in biological systems. Some molecules, such as omega-3 fatty acids, which are environmental and social contained only in certain foods, can help to fight off diseases and illnesses. impacts of various biochemical Collectively, the four biological macromolecules enable cells to function. technologies Understanding how these molecules carry out their roles will help you under- stand how your body functions from day to day. STARTING POINTS Answer the following questions based on your current knowledge. 2. The molecules of living things exhibit a very wide range You will have a chance to revisit these questions later, applying of chemical and physical properties. Brainstorm some concepts and skills you have learned in the chapter. of the properties that might be particularly useful for 1. (a) What kind of foods do you think of when you hear molecules that serve the following biological roles: the terms “carbohydrates,” “fats,” and “proteins”? (a) an energy storage molecule (b) How do you think carbohydrates, fats, and proteins (b) protection from predators compare in terms of their chemical and physical (c) a catalyst properties? (d) the building material for a cell membrane (c) Suggest ways that your body uses these nutrients, and explain how an excess of these nutrients could have a negative impact on your body. 6 Chapter 1 The Biochemical Basis of Life NEL 7923_Bio_Ch01.indd 6 3/27/12 5:11 PM  Mini Investigation Getting Physical with Biochemistry SKILLS Skills: Performing, Observing, Analyzing HANDBOOK A2.1 When living things use a biological substance to perform a 5. Record the state of each substance (solid or liquid) and task, the substance must have suitable properties. One of its solubility. the most important properties of a biological substance is A. How do you think the solubility of a substance affects its solubility. In this activity, you will examine the solubility and chemical reactivity in solution? T/I physical state of some common biological substances. B. Suggest reasons why some substances dissolve in water, Equipment and Materials: balance; test tubes; stirring rods; while others do not. How do you think the size of a 500 mL beaker; distilled water; hot water; samples such as molecule affects its solubility? T/I starch, glucose, sucrose, cellulose fibre, egg albumin, beeswax, C. How might the properties of wax be useful for living butter, olive oil, and table salt things? T/I 1. Measure 0.1 g of each substance. D. Which of the substances you tested formed a suspension, 2. Add each weighed sample to 10 mL of distilled water in a rather than a solution, in water? How do you know? T/I labelled test tube. Stir each mixture. E. Do you think blood is an example of a solution, a 3. Fill the beaker with hot water from the tap. You will use suspension, or both? Explain. K/U this as a water bath. F. Consider milk and other suspensions in nature that 4. Place the test tubes in the water bath, and stir each contain suspended proteins and fats. How do they differ mixture vigorously for 1 min. from solutions and why might suspensions be particularly valuable for living things? T/I NEL Introduction 7 7923_Bio_Ch01.indd 7 3/27/12 5:11 PM  1.1 The Fundamental Chemistry of Life As you learn about all living organisms, you will study everything from the micro- scopic subatomic level all the way to the macroscopic level of the organism. You will begin to understand how everything is uniquely adapted to provide a specific structure and function. You will soon see that the properties of life stem from a hier- archical arrangement of chemical parts. Matter makes up everything in the universe, including all living organisms. Matter is composed of elements. An element is a pure substance that cannot be broken down into simpler substances using ordinary chemical or physical techniques. The smallest particle of an element is an atom. Elements differ from one another in their atomic structure. Atoms often bind to each other chemically in fixed numbers and ratios to form molecules. For example, the oxygen gas we breathe is formed from the chemical combination of two oxygen atoms. A chemical compound is a stable combination of different elements that are held together by chemical bonds. In the study of biology, you will encounter a great variety of chemical compounds. The chemistry of life inside a cell is complex in terms of the sizes and shapes of molecules and their functions in chemical reactions. Even so, all organic (carbon- containing) compounds in living organisms are composed primarily of carbon, C; hydrogen, H; and oxygen, O. As well, they often include nitrogen, N. There are about 21 other elements found in living organisms, but these four elements make up 96 % of the weight of a living organism (Table 1). Most of the other 4 % is composed of only seven other elements: calcium, Ca; phosphorus, P; potassium, K; sulfur, S; sodium, Na; chlorine, Cl; and magnesium, Mg. These elements often occur as ions or in inorganic compounds within living organisms. The rest of the elements that are required by organisms are found in such small amounts (< 0.1 %) that they are called trace elements. Iodine, I, and iron, Fe, are examples of trace elements. A deficiency in any trace element can lead to health problems. Table 1 Percentage Composition of Selected Elements in Living and Non-living Things Seawater % Human % Pumpkin % Earth’s crust % oxygen 88.3 oxygen 85.0 oxygen 65.0 oxygen 46.6 hydrogen 11.0 hydrogen 10.7 carbon 18.5 silicon 27.7 chlorine 1.9 carbon 3.3 hydrogen 9.5 aluminum 8.1 sodium 1.1 potassium 0.34 nitrogen 3.3 iron 5.0 magnesium 0.1 nitrogen 0.16 calcium 2.0 calcium 3.6 sulfur 0.09 phosphorus 0.05 phosphorus 1.1 sodium 2.8 potassium 0.04 calcium 0.02 potassium 0.35 potassium 2.6 calcium 0.04 magnesium 0.01 sulfur 0.25 magnesium 2.1 carbon 0.003 iron 0.008 sodium 0.15 other elements 1.5 silicon 0.0029 sodium 0.001 chlorine 0.15 nitrogen 0.0015 zinc 0.0002 magnesium 0.05 strontium 0.0008 copper 0.0001 iron 0.004 iodine 0.0004 8 Chapter 1 The Biochemical Basis of Life NEL 7923_Bio_Ch01.indd 8 3/27/12 5:11 PM Atomic Structure Elements consist of individual atoms: the smallest units that retain the chemical and physical properties of a particular element. An atom is composed of three subatomic particles: protons, neutrons, and electrons. The number of protons in an atom defines its elemental identity (Figure 1). Protons and neutrons are located in the nucleus. 1p+ Electrons are located in the region surrounding the nucleus. Protons have a positive charge, neutrons have no charge, and electrons have a negative charge. An atom has no net charge because the number of protons is equal to the number of electrons. (a) hydrogen The atomic number of an element is equal to the number of protons in the ele- ment. For example, carbon has an atomic number of 6, so all carbon atoms have 6 protons in their nucleus. A carbon atom also has 6 electrons. The number of neu- trons in the nucleus can vary, as you will learn later in this section. The mass number of an atom is the total number of protons and neutrons in the nucleus. Electrons are not included in this number because the mass of an electron 6p+ is negligible compared with the mass of a proton or neutron. The mass of an atom 6n0 is therefore determined by the number of protons and neutrons it contains. Table 2 lists the atomic numbers and mass numbers of the most common elements in living organisms. The atomic symbol of an element is sometimes shown with the element’s atomic number and mass number (Figure 2). (b) carbon Figure 1 The basic atomic structure of Table 2 Atomic Number and Mass Number of the Most Common Elements in Living Organisms (a) hydrogen and (b) carbon Mass number of the mass number Element Symbol Atomic number most common form 39 19 K hydrogen H 1 1 atomic number carbon C 6 12 Figure 2 Potassium, K, has 19 protons and 20 neutrons in its nucleus. nitrogen N 7 14 oxygen O 8 16 sodium Na 11 23 magnesium Mg 12 24 phosphorus P 15 31 sulfur S 16 32 chlorine Cl 17 35 potassium K 19 39 calcium Ca 20 40 iron Fe 26 56 iodine I 53 127 Isotopes and Radioisotopes All atoms of the same element have the same number of protons, but they may have different numbers of neutrons in the nucleus. This means that atoms with the same atomic number can have different atomic masses. Isotopes are different forms of the isotope a form of an element that differs same element, with different atomic masses. Because isotopes of the same element in its number of neutrons have the same number of protons and electrons, they behave exactly the same in a chemical reaction. NEL 1.1 The Fundamental Chemistry of Life 9  RADIOISOTOPES AND RADIOACTIVE TRACERS The nuclei of some isotopes of an element are unstable and tend to break down, or decay, giving off particles of matter that can be detected as radioactivity. The decay radioisotope a radioactive isotope of process transforms an unstable, radioactive isotope—called a radioisotope—into an an element atom of another element. For example, both hydrogen and carbon have three isotopes, which all behave the same in a chemical reaction (Figure 3). A carbon isotope with the atomic mass 12 (called carbon-12 or 12C) has 6 neutrons, 6 electrons, and 6 pro- tons. Carbon-12 accounts for 99 % of all the carbon in nature. The isotope 13C has 7 neutrons, 6 electrons, and 6 protons. Like 12C, 13C is a stable isotope. A third iso- tope, 14C, has 8 neutrons, 6 electrons, and 6 protons. 14C is an unstable radioisotope of carbon. It decays, giving off particles and energy. As it decays, one neutron splits into a high-energy electron and a proton. The isotope then has 7 neutrons, 7 electrons, and 7 protons. This is characteristic of the most common form of the element nitrogen. Thus, the decay of 14C transforms the carbon atom into 14N, a nitrogen atom. 1 2 3 12 13 14 H H (deuterium) H (tritium) C C C 1 proton 1 proton 1 proton 6 protons 6 protons 6 protons 1 neutron 2 neutrons 6 neutrons 7 neutrons 8 neutrons atomic number = 1 atomic number = 1 atomic number = 1 atomic number = 6 atomic number = 6 atomic number = 6 mass number = 1 mass number = 2 mass number = 3 mass number = 12 mass number = 13 mass number = 14 (a) nuclei of the different isotopes of hydrogen (b) nuclei of the different isotopes of carbon Figure 3 Comparison of the nuclei of different isotopes of (a) hydrogen and (b) carbon Radioactive decay continues at a steady rate, with a constant proportion of radio- isotope atoms breaking down during a given time interval. The rate of decay of a radioisotope is independent of chemical reactions or environmental conditions, such as temperature or pressure. The radiation from decaying isotopes may damage molecules in living cells, thus harming the organism. However, some radioisotopes are useful in geological and biological research because of their steady rate of decay. They provide scientists with information about the age of organic materials, rocks, and fossils. As well, radioisotopes have numerous medical applications and uses in instrumentation to elucidate the structures of unknown compounds. Radioisotopes generally behave the same way in cells as non-radioactive isotopes of the same element. However, because radioisotopes give off a radioactive signal as they decay, they are easily detectable in a cell. Radioactive tracers are radioisotopes that are used to follow a specific chemical through a chemical reaction. Using the particles emitted by a radioisotope as a signal, scientists and doctors can trace the path of the radioisotope as it moves through the cells to different locations in the body. In this way, radioisotopes have found many applications in biological, chemical, and medical research. Scientist Melvin Calvin, a pioneer in the study of photosynthesis, used 14C-labelled molecules to determine the sequence of reactions in photosynthesis. Radioisotopes are used to study many biochemical reactions and to perform basic tech- niques, such as DNA sequencing. Since most biological compounds contain carbon and hydrogen, 14C and 3H (tritium) are commonly used as tracers in biological research. Radioisotopes are also used in the relatively new field of nuclear medicine to help with the diagnosis and treatment of diseases. For example, the thyroid gland pro- duces hormones that affect growth and metabolism. This gland, located in front of the trachea, is the only organ of the body that actively absorbs iodine. If a patient’s symptoms indicate an abnormal level of thyroid hormone output, the physician may inject a small amount of radioactive iodine-131 into the patient and then use a pho- tographic device to scan the thyroid gland. The radioactivity produces an image that is similar to an X-ray, which helps to identify the possible causes of the condition. CAREER LINK WEB LINK 10 Chapter 1 The Biochemical Basis of Life NEL 8159_Bio_Ch01.indd 10 4/2/12 1:40 PM Electron Arrangements The arrangement of electrons determines the chemical properties of an atom, because only electrons are usually directly involved in a chemical reaction. Recall that the number of electrons in an atom is equal to the number of protons in the nucleus. Since the electrons carry a negative charge that is exactly equal in magnitude but opposite to the positive charge of the protons in the nucleus, the atom is electrically neutral. Electrons move around the atomic nucleus in specific regions, called orbitals. An orbital orbital a region of space that is occupied is a region of space that one or two electrons can occupy. Even though one or two elec- by electrons located around the nucleus trons may occupy a given orbital, the most stable and balanced condition occurs when of an atom the orbital contains two electrons. Electron orbitals are grouped into energy levels, which are sometimes called energy shells. These energy shells are numbered 1, 2, 3, and so on, indicating their relative distance from the nucleus. The lowest energy shell of an atom is closest to the nucleus. A maximum of two electrons can occupy the lowest energy shell. The second and third energy shells hold up to eight and 18 electrons, respectively. The first electron shell, the 1s orbital, is a single spherical orbital (Table 3). Hydrogen, for example, has only a 1s electron orbital, containing one electron. Similarly, helium has only a 1s electron orbital, but it contains two electrons. Atoms with more than two electrons have higher energy levels. The shell at the second energy level consists of a 2s orbital and three 2p orbitals. The 2s orbital is spherical in shape. Each 2p orbital shape looks like two balloons tied together. The three 2p orbitals bisect in the centre at right angles to each other, giving the orbitals their overall shape. The orbital that is occupied by an electron is what determines the energy level of the electron. The farther away the electron is from the nucleus, the greater its energy. The balloon-like 2p orbitals contain electrons that are farther away from the nucleus than the electrons in the 2s orbital, and thus hold the electrons with a higher energy level. In large atoms, some higher-energy electrons occupy d and f orbitals, which have even more complex shapes. Table 3 Types of Electron Orbitals 1s orbital 2s and 2p orbital Neon: 1s, 2s, 2p Electron 1s orbital 2s orbital 2p orbitals orbitals (2e!) (2e!) 2pz orbital (2e!) 2py orbital (2e!) 2px orbital (2e!) Electron- shell diagrams Valence electrons are the electrons in an atom’s outermost energy shell, or valence valence electron an electron in the shell. Atoms with an outermost energy shell that is not completely filled with elec- outermost energy level or shell of an atom trons tend to be chemically reactive atoms. Atoms with a completely filled outermost energy level are chemically inactive, or inert. For example, hydrogen has a single electron in the 1s shell, its outermost shell. Hydrogen is a highly reactive element. The helium atom is chemically inert because it has two electrons in its 1s shell, NEL 1.1 The Fundamental Chemistry of Life 11  Table 4 Valence Electron Shell of its outermost shell. (Table 4). Neon has all eight positions in its outer shell occupied. Common Biological Elements and Helium It, too, is an inert and highly stable element (Table 3). Since an unfilled valence shell is less stable than a filled valence shell, atoms with Lewis an incomplete outer shell have a strong tendency to interact with other atoms. They Atomic dot Valence may gain, lose, or share enough electrons to complete their outermost shell. All the number Element diagram shell elements in living organisms have unfilled outermost shells (Table 4) and can there- fore participate in chemical reactions with other atoms. 1 hydrogen H 1s1 Some atoms with outer energy shells that are almost empty or almost full gain or 6 carbon 2s 2 2p 2 lose electrons and form stable charged ions. For example, sodium has two electrons C in its first energy shell, eight in its second energy shell, and only one valence electron in its third energy shell. This outermost single electron is weakly held and readily 7 nitrogen 2s 2 2p 3 N lost to another atom. The sodium atom is left with a completely filled outer shell, becoming a positively charged ion, Na1. In contrast, chlorine has seven electrons 8 oxygen 2s 2 2p 4 in its outermost shell and readily accepts another electron to fill its outer shell com- O pletely and become a negatively charged ion, Cl–. 2 helium He 1s 2 Atoms can also become more stable when they share electrons in such a way that their valence orbitals are filled. This sharing of electrons in the valence shells of atoms creates what are called hybridized electron orbitals. In a hybridized electron orbital, there is a direct overlap of the valence electron orbitals of the two atoms, so the orbital is a combi- nation of two different orbitals. Sharing electrons is the most common way for atoms to bond and form biological molecules. The sharing of electrons by C, H, O, and N under- lies the formation of countless chemical bonds that hold biological molecules together. Chemical Bonds Atoms of inert elements, such as helium, neon, and argon, occur naturally in single-atom form. Atoms of reactive elements, however, combine with each other to form compounds. These atoms form a stable attraction to one another called a chemical bond. Four types of chemical bonds are important in biological molecules: ionic bonds, covalent bonds, and two types of intermolecular forces. You will learn about these bonds in this section. You will also learn about polar molecules, which influence how biological molecules interact. Ionic Bonds ionic bond a bond that results from the An ionic bond forms between atoms that have lost or gained electrons to become attraction between two oppositely charged charged (Figure 4). Atoms that have lost or gained electrons are called ions. Ions of atoms or molecules opposite charge—one positive and the other negative—are strongly attracted to one another, and this attraction leads to an ionic bond. A positively charged ion is called a cation an ion that has a positive charge cation, and a negatively charged ion is called an anion. Ions are very strongly attracted to water molecules. As a result, ionic compounds tend to dissociate and dissolve in anion an ion that has a negative charge water, forming hydrated ions. 1 Cl" and 1 Na! Cl" Na! Cl" Na Cl ! " Na! Cl Na! Cl Na " " ! Cl" Na Cl" Na Cl ! ! " Na! Cl Na! Cl Na " " ! ! Cl" Cl" Na Cl" Na ! Figure 4 Sodium, Na, and chlorine, Cl, atoms form an ionic bond to become NaCl. The sodium atom loses an electron, and the chlorine atom gains an electron. Sodium chloride crystals are cubic in shape. 12 Chapter 1 The Biochemical Basis of Life NEL 7923_Bio_Ch01.indd 12 3/27/12 5:12 PM Covalent Bonds Covalent bonds form when atoms share one or more pairs of valence electrons. The for- – mation of molecular hydrogen, H2, from two hydrogen atoms is the simplest example + + of the sharing mechanism (Figure 5). In H2, two hydrogen atoms share two electrons – and, as a result, both fill their valence electron shell. The strength of a covalent bond depends on the electronegativity of the atoms involved in the sharing of the electron pair. Figure 5 Covalent bond between two Electronegativity is the measure of an atom’s attraction for additional electrons. hydrogen atoms. The atoms share their In molecular diagrams, a dash or a pair of dots represents a single pair of shared electrons to fill the valence shell of electrons in a covalent bond. For example, a molecule of hydrogen is represented as each atom completely, creating a stable H–H or H:H. The number of covalent bonds that an atom can form is usually equal to molecule of hydrogen, H2. the number of additional electrons needed to fill its valence shell. The shared orbitals that form covalent bonds extend between atoms at specific angles and directions, electronegativity the measure of an giving covalently bound molecules distinct three-dimensional forms. atom’s attraction to shared electrons A carbon atom has four electrons in its valence shell and can therefore bond with four hydrogen atoms to form the compound methane, CH4 (Figure 6). Each H hydrogen atom shares its one electron with the carbon atom, completely filling the H C H carbon atom’s valence shell. The carbon atom shares its four valence electrons with H the four hydrogen atoms, completely filling the hydrogen atoms’ valence shells. The bonds between the carbon atom and hydrogen atoms in the compound methane are Figure 6 Lewis dot diagram showing arranged in four symmetrical orbitals, so the hydrogen atoms are 109.58 from each the structure of methane, CH4 other. The overall arrangement of this molecule is tetrahedral. The number and tet- rahedral arrangement of the bonds around a carbon atom allow carbon atoms to link together in more complicated biological compounds (Table 5). Table 5 Orbital and VSEPR models for Methane, Carbon Dioxide, and Water Name and Bonding Non-bonding formula valence valence (shape) Orbital model VSEPR model electron pairs electron pairs methane, CH4 H 4 0 H (tetrahedral) C C H H H H 109.5˚ H H water, H20 2 2 O (angular) H O H H 104˚ H The electrons attempt to move as far away from one another as possible, creating an angle between the bonds. You can predict the arrangement of the bond angles around an atom using the valence shell electron pair repulsion (VSEPR) theory developed by Canadian chemist Ronald J. Gillespie. The VSEPR theory states that because electrons are negatively charged, valence electron pairs repel one another and move as far apart from one another as possible. For example, in a water molecule, there are two valence electron spaces available on the oxygen atom (Table 5). The valence electrons from two hydrogen atoms fill these spaces, forming the molecule H2O. The VSEPR theory also takes into account non-bonding electron pairs of an atom. Oxygen has two electron pairs or four non-bonding electrons in its valence shell. The negative charge of these non-bonding electrons repels the pairs that make up the O–H bonds, so the O–H bonds arrange themselves at an angle of 1048 from each other. NEL 1.1 The Fundamental Chemistry of Life 13  Polar Molecules Although all covalent bonds involve the sharing of valence electrons, they differ widely in the degree of sharing that takes place. The more electronegative an atom is, the more strongly it attracts electrons. Electronegativity is influenced by the atomic number and polar covalent bond a bond between the distance between the valence electrons and the nucleus of an atom. Electronegativity two atoms, made up of unequally shared increases as the distance between the electrons and the nucleus decreases. Oxygen, for electrons example, has a very high electronegativity because two additional electrons can occupy the valence orbitals that are very close to the oxygen’s nucleus. In contrast, hydrogen Table 6 Electronegativity Differences of is less electronegative because, although an additional electron can also get close to its Selected Bonds nucleus, the very small size of the nucleus provides a much weaker force of attraction. Electronegativity The unequal sharing of electrons between two atoms with different electronegativity Bond difference (∆EN ) results in a polar covalent bond. The difference in electronegativity between atoms in common bonds is shown in Table 6. Remember that most biological compounds con- H-H 0.0 tain carbon, hydrogen, and oxygen, and many also contain nitrogen. Due to their rela- C-H 0.4 tively high electronegativities, both oxygen and nitrogen form polar bonds with atoms of most other elements (Table 7). Carbon and hydrogen differ in electronegativity by C-N 0.5 only 0.4. Such a small difference results in bonds that can be considered non-polar. N-H 0.8 The atom that attracts the valence electrons more strongly carries a partial nega- tive charge, which results in the other atom carrying a partial positive charge. Because C-O 0.9 the atoms carry partial charges, the whole molecule may have a non-uniform charge O-H 1.2 distribution. This is the polarity of the molecule. For example, the oxygen atom in water forms a covalent bond with two hydrogen atoms. The oxygen atom attracts the shared electrons much more strongly than the hydrogen atoms (Figure 7(a)). Therefore, the Table 7 Common Polar and Non-polar bonds between oxygen and hydrogen are strongly polar. The more electronegative Bonds in Biological Molecules oxygen pulls the electrons closer, so the oxygen atom has a partial negative charge. Because the water molecule has an asymmetrical shape, with the oxygen atom posi- Polar bonds Non-polar bonds tioned at a bend in the molecule, the partial charges on the atoms create an unequal C-O C-C distribution of charges on the molecule as a whole. In contrast, molecules such as carbon dioxide, CO2, have an equal distribution of charges (Figure 7(b)). Carbon O-H C=C dioxide contains two double bonds with oxygen, making its three-dimensional C=O C-H arrangement appear linear. The bonds are polar, but the symmetry of the molecule results in a balanced distribution of the charges. Overall, the molecule is non-polar. N-H d H d d polarity partial positive or negative O d d H C O H charge at the ends of a molecule H H H water methanol (a) polar molecules d d d O C O O O carbon dioxide oxygen (b) non-polar molecules Figure 7 Comparison of (a) polar and (b) non-polar biological molecules. The lowercase Greek δ! δ" δ! δ" letter delta (d) is used to indicate a partial charge. The symbol d1 indicates partially positive atoms, and the symbol d2 indicates partially negative atoms. δ! Polar molecules attract and align themselves to other polar molecules and tend to Figure 8 Methanol, CH3OH, is pulled be soluble in water. Polar molecules, including water molecules, tend to exclude non- toward water molecules. The partial polar molecules, such as oils and fats. As a result, non-polar molecules have very low positive charge on the hydrogen atom in solubility in polar liquids. Figure 8 shows how a molecule of methanol will align itself CH3OH is pulled by the partial negative with a water molecule. The partial negative charge of water’s oxygen atom is attracted charge in the oxygen atom in water. to the end of the methanol molecule that has a partial positive charge. 14 Chapter 1 The Biochemical Basis of Life NEL 7923_Bio_Ch01.indd 14 3/27/12 5:12 PM  Intermolecular Forces In addition to the intramolecular forces that exist within a molecule, there are forces of attraction between molecules, or intermolecular forces. Intermolecular forces are intermolecular force the force of known as van der Waals forces. They are extremely important because they influence attraction between two molecules the physical properties, such as the solubility, melting point, and brittleness, of a van der Waals forces very weak substance. Van der Waals forces act between similar molecules, as well as between attractions between two molecules, or different types of molecules. The strength of these forces is dependent on the size, parts of molecules, when they are close shape, and polarity of molecules. together HYDROGEN BONDS A hydrogen atom that is covalently bonded to a strongly electronegative atom in one hydrogen bond the attractive force molecule, such as oxygen or nitrogen, can become attracted to a strongly electronega- between a partially positively charged tive atom in a different molecule. The resulting attractive force is called a hydrogen bond hydrogen atom and a partially negatively (Figure 9). Hydrogen bonding is evident in the attraction of the hydrogen atom in charged atom in another molecule CH3OH and the partially negative oxygen atom in H2O (Figure 8). Hydrogen bonds may form between atoms in the same or different molecules. Hydrogen bonds are the strongest and most biologically significant form of van der Waals forces. Individual hydrogen bonds are weak compared with ionic and covalent δ! bonds, but they can be very significant when they occur in large numbers. Although hydrogen bond most of the strongest bonds in living organisms are covalent, weaker hydrogen bonds δ" H are crucial to the function of cells and cellular processes. The large size of biological O molecules offers many opportunities for hydrogen bonding, both within and between H molecules and with surrounding water molecules. A large number of hydrogen bonds are collectively strong, and they lend stability to the three-dimensional structure of large molecules, such as proteins. The hydrogen bonds that exist between water molecules are responsible for many of the properties that make water a uniquely important molecule for all living organ- Figure 9 Each water molecule forms isms. Some of these properties include a very high heat capacity; high melting and hydrogen bonds with up to four boiling points; and cohesion, adhesion, and surface tension. Trees depend on cohe- neighbouring molecules. sion to help transport water through xylem tissue up from their roots. The adhesion of water to the xylem cell walls of a plant helps to counteract the downward pull of gravity. Hydrogen bonds also give water an unusually high surface tension, causing it to behave as if it were coated with an invisible film (Figure 10). The weaker attractive force of hydrogen bonds makes them much easier to break than covalent or ionic bonds, especially when there is an increase in temperature, which increases the movement of the molecules. Hydrogen bonds begin to break extensively as temperatures rise above 45 8C. At 100 8C, hydrogen bonds in water are rapidly overcome. OTHER VAN DER WAALS FORCES Other van der Waal forces are even weaker and result from the momentary attrac- tions of the electrons of one molecule to the nuclei of another molecule. These forces Figure 10 The hydrogen bonds in water develop between all molecules, but they are only significant where other, stronger give it an unusually high surface tension. bonds are not prominent. Such is the case between non-polar molecules or between regions of slightly positive and slightly negative charges within a single molecule. Although an individual bond that has formed as a result of van der Waals forces is weak and transient, the formation of many bonds of this type can stabilize the shape of a large molecule, such as a fat. The size and shape of a molecule influences the number and total strength of van der Waals forces of attraction. A larger molecule has larger forces of attraction. For example, small methane molecules, CH4, are gases at room temperature because the very weak van der Waals forces are unable to hold the molecules together. In contrast, much larger non-polar octane molecules, C8H18, are liquid at room tem- perature because of the cumulative effect of many more van der Waals forces between the larger molecules (Figure 11(a) and (b), next page). Also, linear molecules can align more easily with other molecules, and therefore the van der Waal forces are NEL 1.1 The Fundamental Chemistry of Life 15 8159_Bio_Ch01.indd 15 4/2/12 1:37 PM  stronger. Polar molecules, such as hydrogen chloride, HCl, can also experience an attractive force between the positive and negative ends of the interacting molecules (Figure 11(c)). These influences of size and shape also apply to interactions that involve hydrogen bonds. The effect is very significant in long linear cellulose mol- ecules, which have numerous OH functional groups and are able to form very strong solid fibres. Globular shaped molecules such as starches have fewer accessible atoms for van der Waals forces and therefore tend to form less rigid solids. !" !" Cl Cl H H H H Cl Cl Cl H (a) methane, CH4 (g) (b) octane, C8H18 (l) (c) hydrogen chloride, HCl Figure 11 Non-polar molecules, such as (a) methane and (b) octane, can experience small attractive van der Waals forces, which hold the molecules together. (c) Polar molecules, such as hydrogen chloride, also experience van der Waals forces between their positively and negatively charged regions. Chemical Reactions Thousands of different chemical reactions occur inside cells, but they all have one thing in common. All chemical reactions involve the breaking and formation of chem- ical bonds, thereby changing the arrangements of atoms and ions. In the simple reac- tion between hydrogen gas, H2, and oxygen gas, O2, bonds within the molecules are broken and then new bonds are formed between oxygen and hydrogen to form water. 2 H2 1 O2 S 2 H2O There are four major types of chemical reactions that are common in bio- logical processes: dehydration, hydrolysis, neutralization, and redox reactions. dehydration reaction a chemical reaction Dehydration reactions (also called condensation reactions) consist of the removal of an in which subunits of a larger molecule are 2OH and an 2H from two reactant molecules. The 2OH and 2H form a water mol- joined by the removal of water; also called ecule, while the two reactant molecules are joined together (Figure 12(a)). Dehydration a condensation reaction reactions are the most common method used by cells to join smaller molecules and assemble extremely large macromolecules, such as complex carbohydrates and proteins. H O H H O H H O H H O H H C O H H O C C H H C O C C H H C O H H O C C H H C O C C H H H H H H H H H (a) dehydration reaction H (a) dehydration reaction H H O H O H H H O H O H O H H O H H O H H O H H C O C C H H C O H 1 H O C C H H C O C C H H C O H 1 H O C C H H H H H H H H H (b) hydrolysis reaction (b) hydrolysis reaction Figure 12 (a) Dehydration is the removal of –OH and –H from two reactant molecules to create water and a new bond. (b) Hydrolysis is the cleaving of a bond by the addition of water (as –OH and –H), splitting a larger molecule. 16 Chapter 1 The Biochemical Basis of Life NEL 7923_Bio_Ch01.indd 16 3/27/12 5:12 PM  Hydrolysis reactions are the reverse of dehydration reactions. Water acts as a reactant hydrolysis reaction a chemical reaction to split or “lyse” a larger molecule. In living organisms, hydrolysis breaks down large in which water is used as a reactant molecules into smaller subunits. A bond in the reactant molecule is broken, and the to split a larger molecule into smaller –OH and –H from a split water are attached, resulting in two products (Figure 12(b)). subunits Neutralization reactions occur between acids and bases to produce salts. Water is neutralization reaction a reaction in also often produced in these reactions. An example of a neutralization reaction is the which an acid and a base combine to reaction between hydrochloric acid and sodium hydroxide. The products of these create a salt and water two reactants are water and sodium chloride (common table salt). Redox reactions are the fourth type of chemical reaction. During a redox reaction redox reaction an electron transfer (named for “reduction” and “oxidation”), electrons are lost from one atom and gained reaction by another atom. oxidation Xe– + Y X + Ye– reduction (a)you will oxidation As learn, entire atoms may be transferred during a redox reaction. The term oxidation refers to the loss of electrons. The result is an oxidized molecule or atom. oxidation a reaction in which a molecule oxidation The oxidation of one molecule is always linked to the reduction of another molecule. loses electrons Xe–term The + reduction Y – the +gainYeof electrons. In a redox reaction, the oxidizing refers to X reduction a reaction in which a molecule agent is the molecule or atom being reduced. The reducing agent is the molecule gains electrons CH O2 oxidized. + 2 being or 4atom CONote 2 + energy reduction 2 H2“oxidation” that the+term O also refers to the transfer of entire (a) hydrogen atoms (and their electrons) from less electronegative atoms to more electronegative atoms, as shown in the combustion of methane below. LEARNING TIP reduction oxidation Redox Reactions (b) A helpful pneumonic for remembering the roles of oxidation and reduction is CH4 + 2 O2 CO2 + energy + 2 H2O “LEO the lion says GER”: Loss of Electrons 5 Oxidation Gain of Electrons 5 Reduction reduction (b) During oxidation reactions in biological systems, the electrons involved are more strongly attracted to the oxidizing agent. For example, during the oxidation of methane, CH4, the electrons in the original C–H bonds are strongly attracted to oxygen. As the reaction proceeds, the weaker forces of attraction between the C and H atoms are overcome and the atoms separate. They are then pulled toward and form strong bonds with the oxygen atoms. The products now contain much stronger C5O bonds and O–H bonds, compared with the C–H bonds in methane. Thus, redox reac- tions involve electrons moving from where they are weakly held to where they are more strongly held. As you will learn in Chapter 3, redox reactions are responsible for most of the energy transfer within cells. NEL 1.1 The Fundamental Chemistry of Life 17 8159_Bio_Ch01.indd 17 5/2/12 7:20 AM  1.1 Review Summary Isotopes are different forms of the same element, with different numbers of neutrons. A radioisotope is an unstable isotope that decays to release particles. Valence electrons are the electrons in the outermost electron shell of an atom. There are four types of chemical bonds in biochemistry: ionic, covalent, and hydrogen bonds, and weak van der Waals forces. Electronegativity is a measure of an atom’s attraction for electrons. Differences in electronegativity result in bond polarity. Intermolecular forces are attractive forces between molecules. Dehydration is the removal of –OH and –H from two reactant molecules to form a larger molecule and water. Hydrolysis occurs when a bond in a large molecule is broken, and water is added to the resulting subunits. Oxidation is the loss of electrons, and reduction is the gain of electrons. The oxidation of one molecule or atom is always linked to the reduction of another molecule or atom. This is called a redox reaction. Questions 1. One atom has 6 protons and a mass number of 13. 8. How do polar covalent bonds and non-polar Another atom has 6 protons and a mass number covalent bonds differ? K/U of 15. K/U 9. In a bond between nitrogen and hydrogen (N-H), (a) Identify each of the atoms. which atom will the electrons be closer to? Explain (b) Explain why there is a difference in the mass your reasoning. T/I numbers. 10. Oxygen plays a major role in biological molecules. 2. (a) List the three common isotopes of hydrogen. Explain how oxygen plays a role in polarity, bond (b) What are radioisotopes? shape, and redox reactions. K/U (c) How are radioisotopes used in scientific 11. (a) In what ways do hydrogen bonds produce research and medicine? K/U attractive forces between molecules? Include a 3. An atom has eight electrons, and six of these labelled diagram to illustrate your answer. electrons are valence electrons. K/U C (b) How do hydrogen bonds influence the physical (a) Draw an electron shell model for this atom. properties of water? T/I C Which shells are occupied? 12. Describe dehydration and hydrolysis. How are these (b) What is the name of the element? two types of reactions related? Draw a labelled 4. How do bonding arrangements in a molecule affect diagram to support your answer. K/U C the shape of the molecule? K/U 13. (a) Describe reduction and oxidation. 5. Compare ionic bonds with covalent bonds. K/U T/I (b) Can a reduction reaction occur independently of an oxidation reaction, or vice versa? Why or 6. How can the atomic composition and shape of a why not? K/U molecule affect its polarity? K/U 7. (a) What effect do the polarity, size, and shape of a molecule have on the physical properties of the molecule? (b) How do these factors influence intermolecular forces? K/U T/I 18 Chapter 1 The Biochemical Basis of Life NEL 7923_Bio_Ch01.indd 18 3/27/12 5:12 PM  Water: Life’s Solvent 1.2 Every time we feel thirsty, we are reminded about how much our bodies depend on water for survival. In fact, all living organisms depend on water. Up to 60 % of human body weight comes from water. About 70 % of the brain, 90 % of the lungs, and 22 % of bone tissue is water. Virtually all cellular processes occur in water, since cellular components are dissolved, suspended, and surrounded by water. Without water, we would not exist. Water is a ubiquitous substance—all living organisms contain water, and many kinds of organisms live directly in water (Figure 1). It is both simple in its structure and complex in its functions. More substances dissolve in water than in any other liquid solvent. The reason for the excellent dissolving ability of water is related to its polarity. In this section, you will explore the properties of water. You will learn how it acts as the universal solvent and how it plays a role in so many chemical reactions. Figure 1 Jellyfish are about 95 % water. They have one of the highest ratios of Properties of Water water content to body mass in the animal kingdom. Water is the most abundant liquid on Earth and is known as the “universal solvent.” Water molecules are special because of their size, shape, polar structure, and ability to associate with each other through hydrogen bonding. Hydrogen bonds form readily between water molecules in both liquid water and ice. In liquid water, each water mol- ecule forms an average of 3.4 hydrogen bonds with its neighbouring water molecules. This bonding forms an arrangement known as the water lattice (Figure 2(a)). The water lattice is a unique feature of water. Most molecules that are the size of water, such as H2, O2, CO2, HCl, and H2S, are gases at room temperature. In liquid water, the hydrogen bonds that hold the lattice together constantly break and reform, allowing water molecules to slip past one another and reform the lattice in new positions. This gives liquid water its fluid properties. In ice, the water lattice is a rigid crystalline structure. Each water molecule in ice forms four hydrogen bonds with its neighbouring water molecules. The rigid ice water lattice spaces the water molecules farther apart than they are in the liquid water lattice (Figure 2(b)). Because of the greater spacing, water has the unusual property of being about 10 % less dense in its solid state than in its liquid state. Imagine what Earth would be like if ice sank to the bottom of oceans, lakes, and streams. (a) (b) Figure 2 (a) Hydrogen bonding forms the liquid water lattice. Each water molecule makes an average of 3.4 bonds with its neighbours. (b) Hydrogen bonding forms the ice water lattice. Each water molecule bonds to four of its neighbours, creating a greater distance between the water molecules in ice. As a result of its stabilizing hydrogen bond lattice, water has a high specific heat capacity. Specific heat is the amount of thermal energy that is required to increase specific heat the amount of thermal the temperature of a given quantity of water by one degree Celsius. As thermal energy required to raise the temperature energy flows into a sample of water, much of it is absorbed by the process of breaking of a given quantity of a substance by 1 8C NEL 1.2 Water: Life’s Solvent 19 7923_Bio_Ch01.indd 19 3/27/12 5:12 PM  hydrogen bonds. Therefore, the temperature of water increases relatively slowly as air water surface H2O thermal energy is added. As a result, a lot of thermal energy and a relatively high tem- adhesion the attraction between different perature are needed to break enough bonds in water for it to boil. The high boiling kinds of molecules point of water ensures that it is in a liquid state from 0 8C to 100 8C. Without its hydrogen bond lattice, water would boil at −81 8C. If the boiling point of water were −81 8C, most of the water on Earth would be in the gaseous state. We would not be (a) able to drink it, swim in it, or have it inside our cells. The hydrogen bond lattice of water results in water molecules staying close together—a property called cohesion. Surface tension is related to the concept of cohesion. Surface tension is the measure of how difficult it is to stretch or break the surface of a liquid (Figure 3(a)). Water molecules on the surface of a body of water can form hydrogen bonds on all sides, except the side that faces the air. This creates an imbalance in bonding, which produces a force that places the surface water mole- cules under tension and makes them more resistant