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Short and simple assignment, keep answers related to the information and documents provided Chemical Background for Biology Name _________________________ 10 Points Each 1. Define: Element. How is an element designated? ________________________________________________________________ ________________________________________________________________ ________________________________________________________________ ________________________________________________________________ 2. For each of the symbols, give the name of the element Symbol Element C ____________________ K ____________________ O ____________________ Ca ____________________ Mg ____________________ For each element, give the chemical symbol Element Symbol Hydrogen ____________________ Chlorine ____________________ Phosphorus ____________________ Gold ____________________ Flourine ____________________ 3. Define: Atomic Number, Mass Number. Atomic number: ________________________________________________________________________________________________________________________________. Mass number: ________________________________________________________________________________________________________________________________. 4. Give the atomic number and mass number for the following examples: Atomic number Mass number 8O16 __________ __________ 12Mg24 __________ __________ 19K39 __________ __________ 16S32 __________ __________ 20Ca40 __________ __________ 5. Give the number of protons, the number of neutrons, and the number of electrons for the above examples. Protons Neutrons Electrons 8O16 __________ __________ __________ 12Mg24 __________ __________ __________ 19K39 __________ __________ __________ 16S32 __________ __________ __________ 20Ca40 __________ __________ __________ 6. Explain the difference between a molecule and a compound. ________________________________________________________________ ________________________________________________________________ ________________________________________________________________ ________________________________________________________________. 7. Draw the Bohr diagram of each atom: 8O16 17Cl35 8. Define: Ion, Isotope. ________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________. 9. Define: Ionic bonding, Covalent bonding. ________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________. 10. Explain the difference between inorganic compounds and organic compounds. ________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________. Introduction to Chemistry Study Guide Chemical Level of Organization Introduction Chemistry Chemistry is the study of the composition, structure, and properties of substances, as well as their reactions with one another. Matter All living and nonliving things are made of matter . Matter – anything that takes up space and has mass. Elements The matter of the universe is composed of a limited number of basic substances called elements. Element – a substance that cannot be decomposed into simpler substances by ordinary chemical reactions. An element is a substance that consists only of atoms with the same number of protons (designated the atomic number), and therefore the same nuclear charge. Common examples of elements include carbon, hydrogen, iron, sodium, and chlorine. There are a total of 118 chemical elements. The first 94 of these elements are believed to occur naturally on the earth. The elements with atomic numbers 95 and above do not occur naturally, and are known only as a result of their synthesis in nuclear reactors or particle accelerators. Chemical Symbols Each element is designated by a chemical symbol of either one or two letters that stands for its English or Latin name. Thus H is the symbol for hydrogen, O for Oxygen, C for carbon, Cl for chlorine, Mg for magnesium, K for potassium (the Latin name is Kalium), Na for sodium (the Latin name is Natrium), etc. Exercise For each of the symbols, give the name of the element Symbol Element N ____________________ P ____________________ S ____________________ Ca ____________________ Mg ____________________ For each element, give the chemical symbol Element Symbol Iron ____________________ Iodine ____________________ Mercury ____________________ Lead ____________________ Flourine ____________________ Elements important in Living Organisms The most abundant elements in living organisms are carbon, hydrogen, oxygen, nitrogen, phosphorus, and sulfur. Other elements such as calcium, sodium, potassium, magnesium, and iron are not as abundant but are still essential for life. Trace elements are present in minute quantities yet are also essential for life. Examples include manganese, zinc, copper, and iodine. Atoms Matter is composed of tiny units known as atoms. Atom – the smallest unit of an element, not divisible by ordinary chemical means. A chemical symbol represents one atom of the element, e.g. N stands for a single atom of nitrogen. Atomic Structure Atoms are composed of three fundamental particles: protons, neutrons, and electrons. Fundamental Particles Particle Mass Charge (µ or Daltons) (electronic charge units) Electron 5.485 7990 x 10-4 ‒1 Proton 1.007 276 47 +1 Neutron 1.008 664 90 0 The protons and neutrons are found in the nucleus of the atom. The nucleus contains the positive charge and almost all the mass of an atom. Each proton carries an electronic charge of +1. The neutrons, as their name implies, have no charge. Protons and neutrons have roughly the same mass, which is close to one. Atomic Number Atomic number – the number of protons in the nucleus of an atom. The number of protons in the nucleus is unique for each element. The atomic number is usually written as a subscript immediately before the chemical symbol. For example, 1H indicates that the atomic number of hydrogen is one, i.e. its nucleus contains only one proton. Similarly, 8O indicates that oxygen nuclei contain eight protons. Mass Number The mass number is the total number of protons and neutrons in a nucleus. The mass number is commonly written as a superscript before the chemical symbol. For example, most atoms of oxygen contain eight protons and eight neutrons; the mass number is therefore 16 and the nucleus can be symbolized as O16 or, if we wish to show both the atomic number and the mass number as 8O16. Other Examples: 17Cl35 Chlorine 17 protons 18 neutrons 11Na23 Sodium 11 protons 12 neutrons 6C12 Carbon 6 protons 6 neutrons 7N14 Nitrogen 7 protons 7 neutrons 15P31 Phosphorous 15 protons 16 neutrons If we are given the atomic number and the mass number, we can determine the number of protons, the number of neutrons, and the number of electrons in a normal, neutral atom. For example in 17Cl35 the atomic number is 17 so there are 17 protons. The mass number, the total number of protons plus neutrons is 35. If we subtract the atomic number (17) from the mass number (35) we get the number of neutrons (18). We can also determine the number of electrons. In a normal neutral atom, the number of electrons equals the number of protons. So in chlorine, if there are 17 protons there are also 17 electrons. Isotopes Isotopes – forms of the same element which have the same number of protons, but which differ in the number of neutrons in the nucleus. All atoms of a particular element must have the same number of protons. If the number of protons changed, the atom would become a different element. However, the number of neutrons can be different for atoms of the same element. These different forms of the same element are known as isotopes. Because the number of neutrons is different for the isotopes of an element, so also is the mass number. Examples: There are three different forms of the element Hydrogen: Hydrogen 1 Hydrogen 2 Hydrogen 3 1H1 1H2 1H3 Hydrogen 1 has one proton in the nucleus and no neutrons, Hydrogen 2 has one proton and 1 neutron, and Hydrogen 3 has 1 proton and 2 neutrons. There are three isotopes of Oxygen: Oxygen 16 Oxygen 17 Oxygen 18 8O16 8O17 8O18 Oxygen 16 has 8 protons and 8 neutrons, Oxygen 17 has 8 protons and 9 neutrons, and Oxygen 18 has 8 protons and 10 neutrons. There are three naturally occurring isotopes of carbon: carbon 12, carbon 13, and carbon 14. Carbon 12 Carbon 13 Carbon 14 6C12 6C13 6C14 It is possible to create new isotopes that do not exist in nature using nuclear reactors. Some isotopes are radioactive. Their nuclei are unstable and tend to break down, emitting energy in the form of radiation. Importance of Radioisotopes Radioisotopes have been used as tracers to identify the steps in metabolic reactions. For example, the reactions of photosynthesis were worked out using radioactive carbon dioxide. Radiation is used to treat cancer. Radioisotopes are used to visualize structures in the body to locate disorders. For example, the coronary arteries of the heart can be made visible, allowing the detection of blockage. The Electrons The electrons are negatively charged particles that encircle the nucleus of the atom. Electrons have very little mass. Each electron has a charge of –1; its charge is exactly opposite to that of a proton. In a normal neutral atom, the number of electrons is equal to the number of protons. The positive charges of the protons cancel out the negative charges of the electrons, making the atom as a whole electrically neutral. Consequently, in a neutral atom, the atomic number represents both the number of protons inside the nucleus and the number of electrons circling around the nucleus. Examples: 17Cl35 17 protons 18 neutrons 17 electrons 19K39 19 protons 20 neutrons 19 electrons Orbitals The electrons travel around the nucleus of the atom in regions known as orbitals. The distance of an electron from the nucleus is a function of its energy; the higher the energy, the farther from the nucleus will be the probable location of the electron. The average energy levels of electrons in an atom correspond to a series of so-called electron shells, which can conveniently be represented by concentric circles located at specified distances from the nucleus. In an atom of oxygen, for example, there are two electrons in the first shell and six in the second shell. The Orbitals of electrons may have different shapes. In the first electron shell, this shape is always spherical (it is symbolized by s). In the second electron shell, both the spherical shape (s) and a dumbbell shape (symbolized by p) occur. Additional shapes occur in succeeding electron shells. It has been shown that there is a maximum number of electrons that each shell can contain. The first electron shell can contain a maximum of 2 electrons, the second shell can contain 8, the third shell 18, and the fourth shell 32, etc. Although the third and successive shells can hold more than eight electrons, they are in a particularly stable configuration when they contain only eight. For our purposes, then, the first shell can be considered complete when it holds two electrons and every other shell can be considered complete when it holds eight electrons. Electron Distribution and the Chemical Properties of Elements If the outer shell of electrons is complete, as it is, for example in Helium, which contains 2 electrons in its outer shell, or in neon, which contains 8 electrons in its outer shell, the element has very little tendency to react chemically with other atoms. Valence electrons – the electrons in the outer shell of an atom are known as valence electrons. The valance electrons are important in determining the chemical properties of elements and whether they will combine with one another. Electron Shell Diagrams We can represent the structure of atoms using Electron Shell or Bohr diagrams. In such a diagram, the number of protons and neutrons are indicated in a circle that represents the nucleus of the atom. The electrons are placed in Orbitals around the nucleus. Exercise Given the information below, construct Bohr diagrams of each atom. 19K39 12Mg24 15P31 Atomic Mass The atomic mass of an element is the ratio of its mass to one twelfth the mass of an atom of carbon-12, a unit known as a Dalton or µ. The atomic mass is calculated by averaging the atomic masses of all the chemical element’s isotopes, weighed by isotopic abundance and dividing it by one Dalton (µ), which is equal to 1.660538782 x 10−27 kg. Chemical Bonds Elements combine to form molecules and compounds. Molecule – consists of two or more atoms that have been bound together by chemical bonds A molecule is the smallest chemical unit of a substance that is capable of a stable, independent, existence. Compound – a compound is composed of two or more different kinds of elements joined together by chemical bonds. The difference between a molecule and a compound is that in a molecule the elements comprising the molecule may be the same or different. For example, H2, N2, O2, CH4, and C6H12O6 are all molecules. In the above list, CH4, and C6H12O6 also are compounds because they contain different kinds of elements. Water (H2O) is also a compound. However, H2, N2, and O2 are molecules of the element hydrogen, nitrogen, and oxygen respectively. They contain only one kind of element. Elements combine to form compounds. Example: Sodium + Chlorine → Sodium Chloride There are two types of Chemical Bonds: Ionic bonds and Covalent bonds. Ionic bonding – bonding by transfer of electrons. Covalent bonding – bonding by sharing of electrons. When atoms of elements combine, the atoms usually become more stable by completing their outer shells of electrons (2 electrons for the first shell, eight electrons for the outer shells). Ionic Bonding Ionic bonding is bonding by transfer of electrons. Electrons are transferred from the outer shell of one atom to the outer shell of a second atom. By this process both atoms usually attain stability by completely filling their outer shells with electrons. An example of ionic bonding is the reaction in which sodium combines with chlorine to form sodium chloride. Sodium + Chlorine → Sodium Chloride The substances on the left side of the equation, sodium and chlorine, are reactants. The arrow means yields. The substance(s) on the right side of the equation are products. The equation reads: sodium and chlorine react to yield sodium chloride. Using chemical symbols, 11Na23 + 17Cl35 → NaCl Using Bohr diagrams, Examining the Bohr diagrams for sodium and chlorine, we can see that the sodium atom contains 11 electrons. Two electrons are in the first shell, eight in the second, and there is 1 electron in the outermost shell. Sodium would attain greater stability if it had a complete outer shell of 8 electrons. Chlorine has 17 electrons. Two electrons are in the first shell, 8 electrons are in the second shell, and the remaining 7 electrons are in the third and outermost shell. Chlorine would attain a more stable arrangement by having a complete outer shell of eight electrons. Both sodium and chlorine could attain complete outer shells containing eight electrons if the single electron in the outer shell of sodium were transferred to the outer shell of chlorine. When sodium combines with chlorine, an electron is transferred from the outer shell of sodium to the outer shell of chlorine. This creates an ionic bond joining the two elements together. When sodium loses an electron by transferring it to chlorine, the sodium atom becomes electrically charged. Prior to the transfer, the sodium ion was electrically neutral. There were 11 protons and therefore 11 positive charges in the nucleus. 11 electrons with 11 negative charges canceled these. When the electron was transferred from sodium to chlorine, the electrical charges no longer equaled each other. There are 11 positive charges contributed by the 11 protons, but now there are only 10 negative charges from the 10 electrons. The sodium now has become a charged particle – a sodium ion . Ion – a charged particle. A charged atom or group of atoms. By reacting with sodium, the chlorine gains an electron. It now has 18 electrons and 17 protons. This gives it a charge of –1. The chlorine atom is now a chloride ion. A useful principle to learn is that opposite electrical charges attract one another. Like charges repel one another. Opposite electrical charges attract one another. The positively charged sodium ion is attracted to the negatively charged chlorine ion. This creates an ionic bond joining the two ions together. The product of the reaction is sodium chloride, that is, ordinary table salt. Positive ions are called cations . Negative ions are called anions . These ions are named according to their migration in an electrical field. If the ions are placed in an electrolytic cell and are in solution, the positive ions will migrate toward the negative terminal or cathode. They are therefore called cations. The negative ions will migrate toward the positive terminal or anode. They are therefore called anions. Covalent Bonding Covalent bonding is bonding by sharing of electrons. In forming a single covalent bond, two atoms each share one of their electrons with the other. These two shared electrons effectively fill an orbital in each atom and thus form a covalent bond between these two atoms. The atoms of the elementary gases hydrogen, oxygen, nitrogen, fluorine, and chlorine, form stable diatomic (consisting of two atoms) molecules by covalent bonding. A molecule is the smallest chemical unit of a substance that is capable of a stable, independent existence. Formation of H2 Two hydrogen atoms join together to form a molecule of hydrogen gas. When the atoms combine, each hydrogen atom shares one of its electrons with the other atom. Each hydrogen atom has one electron circling its nucleus. Its outer shell of electrons would be complete if it had two. In order to achieve this stable state, each hydrogen atom shares its electron with the other. Most of the time, the single electron will orbit around its own nucleus. However, part of the time, it will circle the nucleus of the other atom. This creates a strong covalent bond that holds the two atoms together. The sharing of the electrons produces the effect of complete outer shells containing two electrons. For at least some of the time, the two electrons may orbit one of the nuclei, giving it in effect, two electrons. At other times they will circle the other nucleus. In this way they attain stable outer shells. Formation of water The oxygen atom has six electrons in its outer shell. It would attain greater stability by gaining two electrons, thereby completing its outer shell with eight electrons. In forming a molecule of water, the oxygen atom combines with two atoms of hydrogen. Each of the hydrogen atoms shares one of its electrons with oxygen; together with the six electrons that oxygen already has, this gives oxygen a complete outer shell at least part of the time. Each hydrogen atom borrows one electron from oxygen, giving each a complete outer shell as these electrons orbit the hydrogen nuclei. By sharing two pairs of electrons the two hydrogen atoms and the oxygen atom are joined by covalent bonds. The electrons are not shared equally between the hydrogen atoms and oxygen atom. Because oxygen is a larger atom, it has a greater electronegativity. This means that it has a greater attraction for the electrons than hydrogen. As a result, the electrons spend more time circling around the oxygen than the hydrogen. This leads to a greater negative charge around the oxygen, and because the negatively charged electrons are pulled away from the positively charged hydrogen nuclei, a positive charge develops around the hydrogens. Water, thus, is a polar molecule i.e. a molecule with an unequal distribution of electrical charge. Weak chemical bonds Hydrogen Bonds Hydrogen bond – a weak chemical bond between a negatively charged nitrogen or oxygen atom and a positively charged hydrogen atom. Because opposite electrical charges attract one another, a negatively charged oxygen atom will be attracted to a positively charged hydrogen atom. This creates a weak chemical bond between the two atoms. For example, if two water molecules are adjacent to one another, the negative oxygen end of one water molecule is attracted to the positive hydrogen end of the other water molecule. These bonds create links between water molecules and are responsible for the forces of cohesion observed between water molecules. Another example of hydrogen bonds is found in DNA. In DNA, nitrogenous bases are stacked in pairs in the center of the DNA molecule between the two strands which comprise the molecule. Hydrogen bonds form between the negatively charged nitrogen atoms of nitrogenous bases and positively hydrogen atoms of adjacent nitrogenous bases. This creates hydrogen bonds that hold the two strands of DNA together. Chemical reactions Synthesis Reactions When two or more atoms, ions, or molecules combine to form larger molecules, the process is called a synthesis reaction. Example: 2H2 + O2 → 2H2O Decomposition Reactions In a decomposition reaction, large molecules are broken down to form smaller ones. Example: C6H12O6 + O2 → CO2 + H2O Exchange Reactions In an exchange reaction, the atoms of one compound switch places with atoms in another compound. HCl + NaHCO3 ⇌ NaCl + H2CO3 Reversible reactions The above reaction is also a reversible reaction. It can go in either direction. In the above reaction HCl can combine with NaHCO3 to form NaCl and H2CO3, or NaCl can combine with H2CO3 to form HCl and NaHCO3. Chemical Compounds Compounds can be divided into inorganic compounds and organic compounds. Inorganic Compounds Inorganic compounds – compounds which do not contain carbon. Carbon dioxide is an exception to the above definition. Carbon dioxide contains carbon. Nevertheless, it is considered to be an inorganic compound. This is because the properties of carbon dioxide are more like those of the other inorganic compounds. Inorganic compounds are usually small, ionically bonded molecules. They include water, many salts, such as NaCl, and many acids, such as HCl. Organic Compounds Organic compounds – compounds which contain carbon. Carbon dioxide is an exception to the above definition also. Although it contains carbon, It is classified as an inorganic rather than organic compound. Organic compounds are held together primarily by covalent bonds. They tend to be very large molecules. Organic compounds include carbohydrates, lipids, proteins, and nuclei acids. References McQuarrie D, Rock PA, Gallogly EB. General Chemistry. 4th ed. Mill Valley (CA): University Science Books; 2011. Introduction to Chemistry Study Guide draft 9 (Corrected 12/26/2018) PAGE 1 Chapter 2 Lecture Outline Understanding Biology THIRD EDITION Kenneth A. Mason Tod Duncan Jonathan B. Losos © 2021 McGraw Hill. All rights reserved. Authorized only for instructor use in the classroom. No reproduction or further distribution permitted without the prior written consent of McGraw Hill. Because learning changes everything.® The Nature of Molecules and the Properties of Water Chapter 2 ©Fuse/Gett y Images © McGraw Hill ‹#› All Matter is Composed of Atoms Matter has mass and occupies space All matter is composed of atoms Atoms are composed of subatomic particles © McGraw Hill ‹#› 3 Atomic Structure Atoms are composed of three types of subatomic particles Protons Positively charged particles Located in the nucleus Neutrons Neutral particles Located in the nucleus Electrons Negatively charged particles Found in orbitals surrounding the nucleus © McGraw Hill ‹#› 4 Figure 2.2 Access the text alternative for slide images. © McGraw Hill ‹#› Atomic number Number of protons equals number of electrons Atoms are electrically neutral Atomic number = number of protons Every atom of a particular element has the same number of protons Element Any substance that cannot be broken down to any other substance by ordinary chemical means © McGraw Hill ‹#› Atomic mass Mass or weight? Mass – refers to amount of substance Weight – refers to the force gravity exerts on a substance Sum of protons and neutrons is the atom’s atomic mass Each proton and neutron has a mass of approximately 1 Dalton © McGraw Hill ‹#› Electrons Negatively charged particles located in orbitals Neutral atoms have same number of electrons and protons Ions are charged particles – unbalanced Cation – more protons than electrons = net positive charge Anion – fewer protons than electrons = net negative charge © McGraw Hill ‹#› Isotopes Atoms of a single element that possess different numbers of neutrons Radioactive isotopes are unstable and emit radiation as the nucleus breaks up Half-life – time it takes for one-half of the atoms in a sample to decay © McGraw Hill ‹#› Figure 2.3 Access the text alternative for slide images. © McGraw Hill ‹#› Electron arrangement Key to the chemical behavior of an atom lies in the number and arrangement of its electrons in their orbitals Bohr model – electrons in discrete orbits Modern physics defines orbital as area around a nucleus where an electron is most likely to be found No orbital can contain more than two electrons © McGraw Hill ‹#› Figure 2.4 Access the text alternative for slide images. © McGraw Hill ‹#› Atomic energy levels Electrons have potential energy related to their position Electrons farther from nucleus have more energy Be careful not to confuse energy levels, which are drawn as rings to indicate an electron’s energy, with orbitals, which have a variety of three-dimensional shapes and indicate an electron’s most likely location © McGraw Hill ‹#› Figure 2.5 © McGraw Hill ‹#› Redox During some chemical reactions, electrons can be transferred from one atom to another Still retain the energy of their position in the atom Oxidation = loss of an electron Reduction = gain of an electron © McGraw Hill ‹#› Elements Periodic table displays elements according to valence electrons Valence electrons – number of electrons in outermost energy level Inert (nonreactive) elements have all eight electrons Octet rule – atoms tend to establish completely full outer energy levels © McGraw Hill ‹#› Periodic Table of the Elements © McGraw Hill ‹#› Figure 2.6b 90 naturally occurring elements Only 12 elements are found in living organisms in substantial amounts Four elements make up 96.3% of human body weight Carbon, hydrogen, oxygen, Organic molecules contain primarily CHON Some trace elements are very important Access the text alternative for slide images. © McGraw Hill ‹#› Chemical Bonds Molecules are groups of atoms held together in a stable association Compounds are molecules containing more than one type of element Atoms are held together in molecules or compounds by chemical bonds © McGraw Hill ‹#› Ionic bonds Formed by the attraction of oppositely charged ions by electrostatic force Ions form when the atom has a gain or loss of electrons Na atom loses an electron to become Na+ Cl atom gains an electron to become Cl− Opposite charges attract so that Na+ and Cl− remain associated as an ionic compound Electrical attraction of water molecules can disrupt forces holding ions together © McGraw Hill ‹#› Figure 2.8 Access the text alternative for slide images. © McGraw Hill ‹#› Covalent bonds 1 Form when atoms share 2 or more valence electrons Results in no net charge, satisfies octet rule, no unpaired electrons Strength of covalent bond depends on the number of shared electrons Many biological compounds are composed of more than 2 atoms – may share electrons with 2 or more atoms © McGraw Hill ‹#› Covalent bonds 2 Access the text alternative for slide images. © McGraw Hill ‹#› Electronegativity Atom’s affinity for electrons Differences in electronegativity dictate how electrons are distributed in covalent bonds Nonpolar covalent bonds = equal sharing of electrons Polar covalent bonds = unequal sharing of electrons © McGraw Hill ‹#› Hydrogen bonds Electropositive hydrogen from one polar molecule is attracted to an electronegative atom that is often oxygen Attraction produces hydrogen bonds Each individual bond is weak and transitory Cumulative effects are enormous Responsible for many of water’s important physical properties © McGraw Hill ‹#› Van de Waals Attraction Weak bond Non-directional attractive force called Van der Waals forces Form when two atoms are very close to one another Antibodies recognize the shape of an invading organism with this bond © McGraw Hill ‹#› Chemical reactions 1 Chemical reactions involve the formation or breaking of chemical bonds Atoms shift from one molecule to another without any change in number or identity of atoms Reactants = original molecules Products = molecules resulting from reaction © McGraw Hill ‹#› Chemical reactions 2 Extent of chemical reaction influenced by Temperature Concentration of reactants and products Catalysts Many reactions are reversible © McGraw Hill ‹#› Water Life is inextricably tied to water Single most outstanding chemical property of water is its ability to form hydrogen bonds Weak chemical associations that form between the partially negative O atoms and the partially positive H atoms of two water molecules © McGraw Hill ‹#› Figure 2.9 Access the text alternative for slide images. © McGraw Hill ‹#› Polarity of water Within a water molecule, the bonds between oxygen and hydrogen are highly polar Oxygen is much more electronegative than Hydrogen Partial electrical charges develop Oxygen is partially negative δ+ Hydrogen is partially positive δ− © McGraw Hill ‹#› Figure 2.10 Access the text alternative for slide images. © McGraw Hill ‹#› Figure 2.11 Surface tension of water Cohesion – water molecules stick to other water molecules by hydrogen bonding Surface tension due to hydrogen bonds © Hermann Elsenbeiss/National Audubon Society Collection/Science Source. © McGraw Hill ‹#› Figure 2.12 Adhesion – water molecules stick to other polar molecules by hydrogen bonding © McGraw Hill ‹#› Properties of water 1 TABLE 2.3 The Properties of Water Property Explanation Example of Benefit to Life Cohesion/Adhesion Hydrogen bonds cause water molecules to be attracted to other polar or charged species. Leaves pull water upward from the roots; seeds swell and germinate. High specific heat Hydrogen bonds absorb heat when they break and release heat when they form, minimizing temperature changes. Water stabilizes the temperature of organisms and the environment. High heat of vaporization Many hydrogen bonds must be broken for water to evaporate. Evaporation of water cools body surfaces. Lower density of ice Water molecules in an ice crystal are spaced relatively far apart because of hydrogen bonding. Because ice is less dense than water, lakes do not freeze solid, allowing fish and other life in lakes to survive the winter. Solubility Polar water molecules are attracted to ions and polar compounds, making these compounds soluble. Many kinds of molecules can move freely in cells, permitting a diverse array of chemical reactions. Hydrophobic exclusion Water repels hydrophobic compounds, forcing them to associate together. Biological membranes have bilayer structure with hydrophobic interior. © McGraw Hill ‹#› Properties of water 2 Water has a high specific heat A large amount of energy is required to change the temperature of water Water has a high heat of vaporization The evaporation of water from a surface causes cooling of that surface Solid water is less dense than liquid water Bodies of water freeze from the top down © McGraw Hill ‹#› Figure 2.13 Access the text alternative for slide images. © McGraw Hill ‹#› Properties of water 3 Water is a good solvent Water dissolves polar molecules and ions Water organizes nonpolar molecules Hydrophilic “water-loving” Hydrophobic “water-fearing” Water causes hydrophobic molecules to aggregate or assume specific shapes Water can form ions © McGraw Hill ‹#› Acids and bases Pure water [H+] of 10−7 mol/L Considered to be neutral Neither acidic nor basic pH is the negative logarithm of hydrogen ion concentration of solution © McGraw Hill ‹#› Acids and Bases - pH Acid Any substance that dissociates in water to increase the [H+] (and lowers the pH) The stronger an acid is, the more hydrogen ions it produces and the lower its pH Base Substance that combines with H+ dissolved in water, and thus lowers the [H+] (and raises the pH) © McGraw Hill ‹#› Figure 2.14 Access the text alternative for slide images. © McGraw Hill ‹#› Buffers Substance that resists changes in pH Act by Releasing hydrogen ions when a base is added Absorbing hydrogen ions when acid is added Overall effect of keeping [H+] relatively constant © McGraw Hill ‹#› Figure 2.15 Access the text alternative for slide images. © McGraw Hill ‹#› Biological buffers Most biological buffers consist of a pair of molecules, one an acid and one a base Access the text alternative for slide images. © McGraw Hill ‹#› End of Main Content © 2021 McGraw Hill. All rights reserved. Authorized only for instructor use in the classroom. No reproduction or further distribution permitted without the prior written consent of McGraw Hill. Because learning changes everything.® www.mheducation.com Accessibility Content: Text Alternatives for Images © McGraw Hill ‹#› Figure 2.2 - Text Alternative Return to parent-slide containing images. Hydrogen has one positively charged proton in the nucleus and one negatively charged electron in an orbital. Oxygen has 8 protons and 8 neutrons (with no charge) in the nucleus and 8 electrons, two in an inner orbital and 6 in the outer orbital. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.3 - Text Alternative Return to parent-slide containing images. The three isotopes of carbon differ in their number of neutrons, Carbon-12 has 6 neutrons, Carbon-13 has 7 neutrons, Carbon-14 has 8 neutrons. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.4 - Text Alternative Return to parent-slide containing images. Neon has 10 electrons, 2 in the inner shell and 8 in the outer shell. The 2 electrons in the inner shell are in a spherical 1s orbital. Two of the electrons in the outer shell are in a spherical 2s orbital. The remaining 6 electrons in the outer shell are in pairs in three dumbbell shaped p orbitals. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.6b - Text Alternative Return to parent-slide containing images. Carbon, oxygen, hydrogen, nitrogen, sodium, chlorine, calcium, phosphorous, potassium, sulfur, iron, magnesium Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.8 - Text Alternative Return to parent-slide containing images. When sodium gives an electron to chlorine, both have full outer shells. The positive charge on the sodium ion is attracted to the negative charge on the chloride ion, and this attraction forms a sodium chloride salt crystal. Return to parent-slide containing images. © McGraw Hill ‹#› Covalent bonds 2 - Text Alternative Return to parent-slide containing images. Hydrogen gas has two hydrogen atoms sharing a pair of electrons to form one covalent bond. Oxygen gas has two oxygen atoms sharing two pairs of electrons to form two covalent bonds. Nitrogen gas has two nitrogen atoms sharing three pairs of electrons to form three covalent bonds. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.9 - Text Alternative Return to parent-slide containing images. The oxygen has more pull on the electrons, giving it a partial negative charge and the hydrogen atoms a partial positive charge. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.10 - Text Alternative Return to parent-slide containing images. Two water molecules have a hydrogen bond between the positive charge on the hydrogen atom of one water molecule and the oxygen atom on the second water molecule. An organic molecule with an OH group can also form a hydrogen bond with water through its hydrogen atom and the oxygen atom of water. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.13 - Text Alternative Return to parent-slide containing images. A salt crystal dissolves in water when the Na+ ions interact with partial negative charges on water and Cl- ions interact with partial positive charges on water. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.14 - Text Alternative Return to parent-slide containing images. A hydrogen ion concentration of 10 -1 M has a pH of 1 (acidic) down to a concentration of 10 -14 M has a pH of 14 (basic). A list of the pH value of different solutions: hydrochloric acid pH 1, stomach acid pH 2, vinegar pH 3, tomatoes pH 4, coffee pH 5, urine pH 6, water pH 7, sea water pH 8, baking soda pH 9, great salt lake pH 10, ammonia pH 11, bleach pH 13, sodium hydroxide pH 14. Return to parent-slide containing images. © McGraw Hill ‹#› Figure 2.15 - Text Alternative Return to parent-slide containing images. Increasing base is on the x-axis and pH on the y-axis. As base is added the pH increases rapidly at first, then the curve flattens out. This is the buffering range as adding base does not affect pH much. At the end of the graph pH again increases rapidly as base is added. Return to parent-slide containing images. © McGraw Hill ‹#› Biological buffers - Text Alternative Return to parent-slide containing images. Carbonic acid breaks down to bicarbonate ion (HCO3-) and a hydrogen ion (H+). Return to parent-slide containing images. © McGraw Hill ‹#› 61262 22 CHO+6O 6HO+6CO reactants products ® – 2 OHH HO hydroxide ionhydrogen ion + ®+