Physical Science SHS 1st/2nd Sem 2021-2022 PDF

DIVISION OF NAVOTAS CITY Physical Science 1st or 2nd Semester S.Y. 2021-2022 NAVOTAS CITY PHILIPPINES Physical Science for Senior High School Alternative Delivery Mode 1st or 2nd Semester Second Edition, 2021 Republic Act 8293, section 176 states that: No copyright shall subsist in any work of the Government of the Philippines. However, prior approval of the government agency or office wherein the work is created shall be necessary for exploitation of such work for profit. Such agency or office may, among other things, impose as a condition the payment of royalties. Borrowed materials (i.e., songs, stories, poems, pictures, photos, brand names, trademarks, etc.) included in this module are owned by their respective copyright holders. Every effort has been exerted to locate and seek permission to use these materials from their respective copyright owners. The publisher and authors do not represent nor claim ownership over them. Published by the Department of Education Secretary: Leonor Magtolis Briones Undersecretary: Diosdado M. San Antonio Development Team of the Module Writers: Christian James T. Malapitan, Heydiliza A. Santos, Don King O. Evangelista, Lea Z. Malaca, Jasmin B. Tiongson, Editor: Hydiliza A. Santos Reviewer: Russell P. Samson Illustrator: Rodel R. Rimando, SDO-La Union, Region I Layout Artist: Russell P. Samson Management Team: Alejandro G. Ibañez, OIC- Schools Division Superintendent Isabelle S. Sibayan, OIC- Asst. Schools Division Superintendent Loida O. Balasa, Chief, Curriculum Implementation Division Russell P. Samson, EPS in Science Grace R. Nieves, EPS In Charge of LRMS Lorena J. Mutas, ADM Coordinator Vergel Junior C. Eusebio, PDO II LRMS Inilimbag sa Pilipinas ng ________________________ Department of Education – Navotas City Office Address: BES Compound M. Naval St. Sipac-Almacen Navotas City ____________________________________________ Telefax: 02-8332-77-64 ____________________________________________ E-mail Address: ____________________________________________ [email protected] Table of Contents Quarter 1 or Quarter 3 What I Know................................................................................1 Module 1......................................................................................2 Module 2......................................................................................11 Module 3......................................................................................20 Module 4......................................................................................30 Module 5......................................................................................36 Module 6......................................................................................43 Module 7......................................................................................47 Module 8......................................................................................50 Assessment..................................................................................52 Quarter 2 or Quarter 4 What I Know................................................................................54 Module 9......................................................................................55 Module 10....................................................................................62 Module 11....................................................................................66 Module 12....................................................................................70 Module 13....................................................................................73 Module 14....................................................................................75 Module 15....................................................................................81 Module 16....................................................................................84 Assessment..................................................................................89 Answer Key..................................................................................91 References...................................................................................96 Directions: Choose the letter of the best answer. Write the chosen letter on a separate sheet of paper. 1. What causes the formation of heavier atomic nuclei during big bang? A. the increase in temperature C. the unchanging temperature B. the cooling down of the universe D. none of the above 2. Sodium chloride is made up of sodium (Na) ion and chlorine (Cl) ion. Which of the following types of bonds is describe the compound? A. covalent bond C. nonpolar covalent bond B Ionic bond D. polar covalent bond 3. Why are dispersion forces high in molecules with great number of electrons? A. The electron distribution of big molecules is easily polarized. B. The nucleus in the molecules has greater effective shielding effect. C. The electrons move freely around the nucleus resulting to greater energy. D. The electrons in the molecules can easily jump from one orbital to another. 4. Cellulose is the most abundant organic chemical on Earth. Cellulose consists of long chains of glucose. What type of carbohydrate is cellulose? A. Monosaccharide C. Oligosaccharide B. Disaccharide D. Polysaccharide 5. Which of the following would NOT increase the rate of reaction? A. adding catalyst B. raising the temperature C. increasing the volume of the container D. increasing the concentration of the reaction 6. Why the dust particles suspended in the air inside unheated grain elevators can sometimes react explosively? A. because of high kinetic energy. B. because of high activation energy. C. because it has the catalytic effect on the reaction. D. because it has a large surface area for the reaction. 7. Which of the following statements is FALSE for the chemical equation given below in which nitrogen gas reacts with hydrogen gas to form ammonia gas assuming the reaction goes to completion? N2 + 3H2 2NH3 A. One mole of N2 will produce two moles of NH3 B. The reaction of 14 g of nitrogen produces 17 g of ammonia C. The reaction of one mole of H2 will produce 2/3 moles of NH3 D. The reaction of three moles of hydrogen gas will produce 17 g of ammonia. 1 8. What is the ratio of the actual yield to the theoretical yield, multiplied by 100%? A. mole ratio C. percent yield B. excess yield D. Avogrado yield 9. Hydroelectric power is a renewable source of electricity. Which of the following energy phase that the hydroelectric power plant is a source of electricity? A. electricity is extracted from water B. water is converted into steam to produce electricity C. potential energy possessed by stored water is converted into electricity D. kinetic energy possessed by stored water is converted into potential energy 10. What is the best way to dispose of hazardous chemicals at home? A. bury them on a yard B. carefully pour them down the drain C. read the labels to see how to dispose on each D. put them in a leak-proof container in the trash MODULE 1 This module was designed and written with you in mind. It is here to help you master the formation of the elements during the Big Bang and during stellar evolution as well as the distribution of the chemical elements and the isotopes in the universe. The scope of this module permits it to be used in many different learning situations. The language used recognizes the diverse vocabulary level of students. The lessons are arranged to follow the standard sequence of the course. But the order in which you read them can be changed to correspond with the textbook you are now using. The module includes lessons, namely: Lesson 1.1: Formation of Light Elements during Big Bang Lesson 1.2: Stellar Nucleosynthesis and the Formation of Heavy Elements Lesson 1.3: From Outer Space to Laboratory After going through this module, you are expected to: 1. Give evidence for and describe the formation of heavier elements during star formation and evolution 2. Explain how the concept of atomic number led to the synthesis of new elements in the laboratory Lesson Formation of Light Elements 1.1 during Big Bang 2 In the beginning….. The sudden dropped of temperature to billion degrees caused the combination of protons and neutrons. They stayed together to form nuclei of the simplest elements like hydrogen, helium and lithium. However, in normal circumstances, protons and neutrons repel each other because a positively charged nuclei (proton) do not interact with an uncharged nucleus (neutron). Nuclear fusion (meaning: fusion = union) brought together these neutrons and protons utilizing the very high temperature. The formation of the nuclei of light elements a few minutes after the Big Bang is called Big Bang nucleosynthesis or primordial nucleosynthesis (nucleo refers to nuclear and synthesis means combination or parts to form a whole). The formation of nuclei of ordinary hydrogen, its two isotopes, deuterium and tritium, and the common isotope of helium is shown in the illustration below. Nucleosynthesis Note: As the Universe cooled down to about a trillion degrees, protons and neutrons fused to form heavier atomic nuclei proton p p deuterium p p p n n n n n n neutron tritium helium n It is clearly shown above that the combination of these particles is capable of forming helium, (and as hydrogen nuclei are simply protons anyway, hydrogen does not need to form by combinations of anything) and the unstable helium-3, tritium and deuterium simply break down. But why does the helium not also simply break down again under pressure, and why do heavier elements not form? In fact, heavier elements do form, but like the helium-3, tritium and deuterium they are formed as unstable isotopes. Ordinary helium, (helium-4) is the most stable isotope formed in this mixture of fundamental particles during the early stage of Big Bang. Because of very high pressures and temperatures, many helium nuclei did break down again, but it must be remembered that at the time, the Universe was rapidly expanding and cooling. Remember that almost all nuclei heavier than hydrogen formed would be unstable and break down (see the illustration). The only other elements formed to even a small degree was deuterium, and lithium-7, which though stable requires the collision of several particles simultaneously, and it therefore very unlikely. The exact ratios of hydrogen, helium, lithium and deuterium created during the period of Big Bang nucleosynthesis would have been very sensitive to the density of the mixture of fundamental particles from which they were made. Because the mixture was not entirely uniform in density, different amounts of each were created in different areas of the compacted early universe. As an average though, 3 approximately twelve protons (hydrogen nuclei) were left for each helium nucleus produced. proto p n n p p p Lithium-5 n p n p Unstable n neutro p n p n p n p p Helium-5 n n Unstable n p n p n p n p n p p Beryllium-8 n n p n p n Unstable So, after four minutes, the Universe had made the first basic nuclei, and this marks the end of primordial nucleosynthesis. Currently, the universe was still too hot to allow these nuclei to attract electrons and form atoms. After 100 millenia it had created the first atoms. Activity 1: Formation of elements Directions: Arrange the correct sequence of stages of formation of light elements in the below using the Graphic organizer below. Use the choices below the graphic organizer. Sequential Order 1. 2 3 4. 5. Protons, Electrons combine Stars ignite by neutrons emerge with H, He and burning H and from the cooling other nuclei to He. The cosmos quark soup form neutral atoms lights up. Simple atomic Gravity causes the nuclei are formed diffuse H2/He gas (H, He) to form clouds which collapse into stars 4 Lesson Stellar Nucleosynthesis and the 1.2 Formation of Heavy Elements ….and there was light After Big Bang, the universe continued to expand. The clouds of gas floating are made of clouds of hydrogen and helium, which are left over from supernovas. Matter started to form lumps, and some of which clumped together and formed galaxies. These clumped matters are called protogalaxies (also called primeval galaxy, is a cloud of gas which is forming into a galaxy). Inside these protogalaxies, the internal lumps collapsed due to gravitational forces. The compression brought by gravitational forces heats the interior of these clumps which resulted to an increase in internal temperature. An increased internal temperature initiated the nuclear fusion, involving hydrogen and producing helium along with a tremendous amount of energy. This caused the shining of the star. The production of nuclei heavier than hydrogen in stars is called stellar nucleosynthesis. Stellar nucleosynthesis is the process by which elements are created within stars by combining the protons and neutrons together from the nuclei of lighter elements. All of the atoms in the universe began as hydrogen. Fusion inside stars transforms hydrogen into helium, heat, and radiation. Heavier elements are created in different types of stars as they die or explode. When hydrogen in the core of a star is depleted, the core of the star shrinks, and the core temperature rises. When the temperature is high enough, helium fusion becomes possible, and the heavier carbon nucleus is formed. Depending in the mass of the star this series of fusion reactions in the core followed by shrinking and temperature rise that trigger more fusion takes place successively forming elements. The burning of helium to produce heavier elements then continues for about 1 million years. Largely, it is fused into carbon via the triple- alpha process in which three helium-4 nuclei (alpha particles) are transformed. The alpha process then combines helium with carbon to produce heavier elements, but only those with an even number of protons. The combinations go in this order: 5 1. Carbon plus helium produces oxygen. 2. Oxygen plus helium produces neon. 3. Neon plus helium produces magnesium. 4. Magnesium plus helium produces silicon. 5. Silicon plus helium produces sulfur. 6. Sulfur plus helium produces argon. 7. Argon plus helium produces calcium. 8. Calcium plus helium produces titanium. 9. Titanium plus helium produces chromium. 10. Chromium plus helium produces iron. In stars more massive that about ten solar masses, the successive burning to produce heavier nuclei eventually produces a nonburning core of iron. Iron-56 and other nuclei with mass number of 56 are stable nuclei and their fusion with a helium nucleus would require energy instead of producing it. As a result, stellar nuclear burning stops at iron. With no further production of energy at the core the gravitational forces take over and the core collapses, quickly and catastrophically. The massive star that ends up as supernovas are manufacturers of elements heavier than iron. These heavier elements like gold, silver, lead and mercury require the very special conditions of pressure and heat that exist inside a supernova during those few seconds of the collapse. Supernovae are also distributors of the elements in the universe because during the explosion, all the elements that have been formed are flung into space. Our own Sun and the planets in the solar system were formed from the debris of earlier stars that have run through their life cycles. Since then, the Sun shines! Activity 2: Formation of elements Directions: Answer the following, identify the number of proton and neutron, together with the name of the product yield in fusion. (refer to the example below). Write your answer on the space provided for. Blue Circle (Proton) Orange Circle (Neutron) Deuterium Tritium Helium – 4 1 neutron 1 proton 1 proton 2 protons 1 neutron 2 neutrons 2 neutrons 6 Using your periodic table of elements, identify the name of the element, if the number of protons is 2, its atomic number is the same, therefore the name of the element is HELIUM, to identify the atomic mass, follow the formula (atomic mass = number of proton + Number of Neutron). Thus Helium 4 is the product of two deuterium, the number for added to the name of the product refers to its atomic mass. Helium - 4 Tritium ______________ ___________ ______ proton _______ protons ______ protons ______ neutrons _______ neutrons ______ neutrons Lesson From Outer Space to Laboratory 1.3 Timeline of Element Discovery Note sometimes there are two dates for an element, when discovery and isolated were separated. For more recently discovered elements, discovery dates and credit may be hotly disputed even today. Ancient Times: Before 1 A.D. Gold, Silver, Copper, Iron, Lead, Tin, Mercury, Sulfur, Carbon Time of the Alchemists: 1 A.D. to 1735 Arsenic (Magnus ~1250), Antimony (17th century or earlier), Phosphorus (Brand 1669), Zinc (13th Century India), 1735 to 1745 Cobalt (Brandt ~1735), Platinum (Ulloa 1735) 1745 to 1755 Nickel (Cronstedt 1751), Bismuth (Geoffroy 1753), 1755 to 1765 No new elements discovered in this date range. 1765 to 1775 Hydrogen (Henry Cavendish 1766), Nitrogen (Rutherford 1772), Oxygen (Priestley; Scheele 1774), Chlorine (Scheele 1774), Manganese (Gahn, Scheele, & Bergman 1774) 1775 to 1785 Molybdenum (Scheele 1778), Tungsten (J. and F. d’Elhuyar 1783), Tellurium (von Reichenstein 1782) 1785 to 1795 7 Uranium (Peligot 1841), Strontium (Davey 1808), Titanium (Gregor 1791), Yttrium (Gadolin 1794) 1795 to 1805 Vanadium (del Rio 1801), Chromium (Vauquelin 1797), Beryllium (Discovery: Louis Nicloas Vauquelin 1798; Isolated: Friedrich Wöhler & Antoine Bussy 1828), Niobium (Hatchett 1801), Tantalum (Ekeberg 1802), Cerium (Berzelius & Hisinger; Klaproth 1803), Palladium (Wollaston 1803), Rhodium (Wollaston 1803-1804), Osmium (Tennant 1803), Iridium (Tennant 1803) 1805 to 1815 Sodium (Davy 1807), Potassium (Davy 1807), Barium (Davy 1808), Calcium (Davy 1808), Magnesium (Black 1775; Davy 1808), Boron (Discovery: Gay- Lussac & Thenard June 1808; Isolated: Humphry Davy July 1808), Iodine (Courtois 1811) 1815 to 1825 Lithium (Discovery: Johan August Arfvedson 1817; Isolated: William Thomas Brande 1821), Cadmium (Stromeyer 1817), Selenium (Berzelius 1817), Silicon (Berzelius 1824), Zirconium (Klaproth 1789; Berzelius 1824) 1825 to 1835 Aluminum (Wohler 1827), Bromine (Balard 1826), Thorium (Berzelius 1828) 1835 to 1845 Lanthanum (Mosander 1839), Terbium (Mosander 1843), Erbium (Mosander 1842 or 1843), Ruthenium (Klaus 1844) 1845 to 1855 No new elements discovered in this date range. 1855 to 1865 Cesium (Bunsen & Kirchoff 1860), Rubidium (Bunsen & Kirchoff 1861), Thallium (Crookes 1861), Indium (Riech & Richter 1863) 1865 to 1875 Fluorine (Moissan 1866) 1875 to 1885 Gallium (Boisbaudran 1875), Ytterbium (Marignac 1878), Samarium (Boisbaudran 1879), Scandium (Nilson 1878), Holmium (Delafontaine 1878), Thulium (Cleve 1879) 1885 to 1895 Praseodymium (von Weisbach 1885), Neodymium (von Weisbach 1885), Gadolinium (Marignac 1880), Dysprosium (Boisbaudran 1886), Germanium (Winkler 1886), Argon (Rayleigh & Ramsay 1894) 1895 to 1905 Helium (Discovery: Pierre Janssen and Norman Lockyer 1868; Isolated: William Ramsay, Per Teodor Cleve, Abraham Langlet 1895), Europium (Boisbaudran 1890; Demarcay 1901), Krypton (Ramsay & Travers 1898), Neon (Ramsay & Travers 1898), Xenon (Ramsay & Travers 1898), Polonium (Curie 1898), Radium (P. & M. Curie 1898), Actinium (Debierne 1899), Radon (Dorn 1900) 1905 to 1915 Lutetium (Urbain 1907) 1915 to 1925 Hafnium (Coster & von Hevesy 1923), Protactinium (Fajans & Gohring 1913; Hahn & Meitner 1917) 8 1925 to 1935 Rhenium (Noddack, Berg, & Tacke 1925) 1935 to 1945 Technetium (Perrier & Segre 1937 ), Francium (Perey 1939), Astatine (Corson et al 1940), Neptunium (McMillan & Abelson 1940), Plutonium (Seaborg et al. 1940), Curium (Seaborg et al. 1944) 1945 to 1955 Mendelevium (Ghiorso, Harvey, Choppin, Thompson, and Seaborg 1955), Fermium (Ghiorso et al. 1952), Einsteinium (Ghiorso et al. 1952), Americium (Seaborg et al. 1944), Promethium (Marinsky et al. 1945), Berkelium (Seaborg et al. 1949), Californium (Thompson, Street, Ghioirso, and Seaborg: 1950) 1955 to 1965 Nobelium (Ghiorso, Sikkeland, Walton, and Seaborg 1958), Lawrencium (Ghiorso et al. 1961), Rutherfordium (L Berkeley Lab, USA – Dubna Lab, Russia 1964) 1965 to 1975 Dubnium (L Berkeley Lab, USA – Dubna Lab, Russia 1967), Seaborgium (L Berkeley Lab, USA – Dubna Lab, Russia 1974) 1975 to 1985 Bohrium (Dubna Russia 1975), Meitnerium (Armbruster, Munzenber et al. 1982), Hassium (Armbruster, Munzenber et al. 1984) 1985 to 1995 Darmstadtium (Hofmann, Ninov, et al. GSI-Germany 1994), Roentgenium (Hofmann, Ninov et al. GSI-Germany 1994) 1995 to 2005 Nihonium – Nh – Atomic Number 113 (Hofmann, Ninov et al. GSI-Germany 1996), Flerovium – Fl – Atomic Number 114 (Joint Institute for Nuclear Research and Lawrence Livermore National Laboratory 1999), Livermorium – Lv – Atomic Number 116 (Joint Institute for Nuclear Research and Lawrence Livermore National Laboratory 2000), Oganesson – Og – Atomic Number 118 (Joint Institute for Nuclear Research and Lawrence Livermore National Laboratory 2002), Moscovium – Mc – Atomic Number 115 (Joint Institute for Nuclear Research and Lawrence Livermore National Laboratory 2003) 2005 to Present Tennessine – Ts – Atomic Number 117 (Joint Institute for Nuclear Research, Lawrence Livermore National Laboratory, Vanderbilt University and Oak Ridge National Laboratory 2009) Will There Be More Elements? The discovery of 118 elements completes the first seven periods of the periodic table, but scientists are working to synthesize new elements. When another discovery is verified, another row (period) will be added to the table. However, before scientists or chemist discovered these elements, they studied first the basic – the concept of atomic number and the isotopes. 9 These are example of isotopes. What have you noticed? Normally all of them are Hydrogen element. However, you will notice that they have different atomic weight (number on the upper part of the element). What thus that implies to us? What does this mean? Atomic Mass Atomic Number 1. The atomic mass is equal to the sum of the number of proton and neutron of an element. 2. The Atomic number is equal to the number of protons of an element Here is a simple description of the relationship of atomic number and atomic mass to the subatomic particles of an element Atomic number = Number of Proton Atomic Mass = Number of Proton + Number of Neutron Number of Neutron = Atomic Mass – Atomic Number Number of Electron = Number of Proton Following the formula above, we can simply identify the differences in terms of the number of neutrons of each isotope. Number of Proton 1 1 1 Number of Neutron 0 1 2 Atomic Mass 1 2 3 Number of Electron 1 1 1 Activity 3. Isotopes Directions: Indicate the number of protons, neutrons, atomic mass, and electron of the following isotopes of carbon element. Indicate your answer in the table below, show your solution on the space provided for/separate worksheet Number of Proton Number of Neutron Atomic Mass Number of Electron 10 Activity: Infographic Making Directions: In this activity, you will create an infographic that demonstrates your understanding of the Big Bang. The infographic needs to include attention to the following: Topic: The topic of the infographic is specific in nature and is intended to inform or convince the viewer. Type: The type of infographic chosen (for example, timeline or informational) highly supports the content being presented. Objects: The objects included in the infographic are relevant and support the topic of the infographic. Data visualizations: The data visualizations present accurate data and are easy to understand. Style: Fonts, colors, and organization are aesthetically pleasing, appropriate to the content, and enhance the viewer’s understanding of the information in the infographic. Citations: Full bibliographic citations for all sources used are included. MODULE 2 This module was designed and written with you in mind. It is here to help you master Physical Science. The scope of this module permits it to be used in many different learning situations. The language used recognizes the diverse vocabulary level of students. The lessons are arranged to follow the standard sequence of the course. But the order in which you read them can be changed to correspond with the textbook you are now using. The module is divided into two lessons, namely: Lesson 2.1 – Electronegativity and Polarities of Molecules Lesson 2.2 – Molecular Structures and Polarities Lesson 2.3 – Polarities and Properties After going through this module, you are expected to: 1. Recall the definition of electronegativity. 2. Compute for electronegativity differences between two bonded elements. 3. Infer the polarity of a molecule given its molecular geometry. 4. Differentiate polar from non-polar compounds based on its properties. 11 Lesson Electronegativity and Polarity 2.1 Recall from your Junior High School years. Elements combine to form compounds. When they form compounds, they either give up, take, or share their valence electrons (outermost electrons). When they give up, take, or share their valence electrons, just like tug-of-war, there are differences in their pull of electrons. This is determined by their electronegativity. Electronegativity (EN) is a measure of the relative tendency of an atom to attract electron to itself when chemically combined with another atom. The higher the value of electronegativity, the more it tends to attract toward itself. For the representative elements, electronegativity usually increases from left to right across periods and decreases from top to bottom within group the most reactive metals then. The elements in the lower left-hand corner of the periodic table have the lowest electronegativity values. These are consistent with the trend in ionization energy. Figure 1 General trend in electronegativity of representative elements Although electronegativity is an atom’s ability to attract the electrons, it also involved in bonding. Because no two elements have the same electronegativities, in a covalent bond between different elements, one of the atoms attracts the shared pair more strongly than does the other. The differences of the electronegativities are determined by the EN or electronegativity difference. The result of their EN will determine the polarity of the bond, whether it is polar covalent, non-polar covalent, or ionic. The value of electronegativity is first determined by Linus Pauling. Electronegativity Difference Type of Bond (∆EN) Ionic ≥1.7 Polar Covalent 0.5 to 1.6 Nonpolar Covalent ≤ 0.4 12 The formula for getting the EN of a pair of elements is determined by this equation: EN = | ENelement 1 – ENelement2 | Let us try for the following element pair: HCl EN of H =2.1 EN of Cl =3.0 EN = 0.9 Therefore, the bonds that exist between hydrogen and chlorine in hydrogen chloride is polar covalent because it is between 0.5 to 1.6. Look at the electron density diagram or the diagram that shows the distribution of valence electrons of HCl. You can see that Cl is a lot bigger than H. Going back to our tug-of-war illustration, we can infer that Cl will pull the H atom towards itself. Therefore, it is polar covalent. H is now a partially positive (ẟ+) element and Cl now is a partially negative (ẟ-) element. The pull now is towards chlorine. The arrow above is called a electric dipole. Let us try another pair of elements: Cl2 EN of Cl =3.0 EN = 0 The electron density diagram of Cl 2 shows two chlorine atoms. Since they are the same element of the same size, if we apply the tug-of-war illustration, there will be an equal pull among the poles. Thus, we cannot see a electric dipole or a dipole moment. Moreover, the electronegativity difference is zero. Therefore, we consider it as a nonpolar covalent compound. The presence of a polar bond in a molecule often makes the entire molecule polar. In a polar molecule one end of the molecule is slightly negative, and one end is slightly positive, in the hydrogen chloride molecule the partial charges on hydrogen and chlorine atoms are electrically charged region, or poles. A molecule that that has two poles is called a dipolar molecule, or dipole, hydrogen and chloride molecule is a dipole. When polar molecules are placed in an electric field, the negative ends of the molecules orient toward the positively charged plate and the positive ends of the molecules orient toward the negatively charged plate. The positive and negative charges in a nonpolar molecule experience force in opposite direction as a result of their opposite polarities. This force causes the electron cloud of a nonpolar molecule to be displaced in the direction of the attraction. 13 Let us have a final example: ENNa = 0.9 ENCl = 3.0 (∆EN) = 0.9 – 3.0 = │-2.1│= 2.1 ionic EN differences in an ionic bond is so strong that there is no sharing of electrons but rather a giving and taking of electrons. In your Grade 9 Science, you have learned that metals are always in the business of giving up electrons and non-metals are always taking electrons to form their complete octet. As per Dr. Magno of UP Diliman, there are limitations to the concept of electronegativity when applied to some molecules, such as BF 3 (EN = 1.94) and SiF4 (EN = 2.08). These molecules have a large electronegativity difference that we would think that they are ionic, but both compounds are covalent. They exist as gases in room temperature, a property not held by ionic compounds. How can we explain this covalent nature? The electronegativity of an atom in a molecule is not precisely constant. It may vary slightly to either side of the stated value depending on its environment. For example, fluorine is less electronegative when it bonds with metalloids like B and Si, or nonmetals such as C that it bonds with metals such as Na or Mg 1. Use the illustration below as a summary to this lesson: (Source: Raymond Chang’s Chemistry 9th Edition) Activity 1. Give and Take or Sharing Directions: A. Identify the type(s) of Bond(s) found on the following molecules: (Ionic – metal and non-metal, Covalent – both non-metal) Compounds Type(s) of Bonds 1. CCl4 2 Li2O 3 NF3 4 CAs 5 SO2 6 Mg (OH)2 7 BeCl2 1 For more information, you can refer to Marcelita Magno’s Foundational Course in College Chemistry, pp 279 – 280 for more information. 14 Directions: B. Determine if the bond between atoms in each sample below is nonpolar covalent, polar colavent or ionic. Show the electronegativity differences in each. 8 H2 9 PCl 10 F2 11 NaBr 12 NF 13 MgO 14 CH 15 HCl Lesson The Role of Molecular Geometry 2.2 in the Polarity of a Molecule You just have learned how to predict the type of bond polarity simply by calculating the electronegativity difference of atoms (specifically two atoms). How about for those molecules consisting of more than two atoms like H2O, CCl4, NH3 and CO2? For polyatomic molecules, both the bond polarity and molecular shape determine the overall molecular polarity. In terms of molecular geometry, the valence shell electron pair repulsion (VSEPR) theory would help us to determine the spatial arrangement of atoms in a polyatomic molecule. The molecular geometry is important to determine if a molecule is polar or not. Valence shell electron pair repulsion theory, or VSEPR theory, is a model used in chemistry to predict the geometry of individual molecules from the number of electron pair surrounding their central atoms. It helps predict the spatial arrangement of atoms in a polyatomic molecule. The shapes are designed to minimize the repulsion within a molecule. According to VSEPR, there are seven basic shapes of molecules stipulated below. You can predict the shape of the molecules if you know the number of total valence electrons, number of bonding pairs, non-bonding pairs, and lone pairs in the central atom. For deeper discussion regarding drawing plausible Lewis structures and molecular geometry, it will be discussed in Chemistry 1 (STEM/GAS). For Physical Science, we will deal more with symmetry and polarity. 15 Here is a simple flowchart that you can follow to know if a molecule is polar or non-polar given its molecular structure. NO Is the shape YES symmetrical in 3D? NO Are all atoms The molecule is bonded to the POLAR central atoms the same? YES The molecule is NONPOLAR Figure 4. Flowchart to determine if a molecule is polar or nonpolar Remember that as you look at the structure of the element, you need to see it at a 3-D perspective. A flat thinking will not suffice. You need to think that some of the elements is at the BACK of this page or at the FRONT of this page. Let us have some examples so that you can visualize. Examples: CH4 (methane) has four hydrogen atoms attached to the central Carbon atom. A. Is the shape symmetrical in 3D? Yes B. Are all atoms bonded to the central atoms the same? Yes Therefore, CH4 is nonpolar. (Note: Broken lines in a Lewis Diagram means that the bonded atom is at the BACK of the page.) 16 H2O (water) has two hydrogen atoms attached to the central oxygen atom with two lone pairs at the central atom. A. Is the shape symmetrical in 3D? Yes B. Are all atoms bonded to the central atoms the same? No. There is a presence of lone pairs. Therefore, water is a polar molecule. Polar molecules have dipole moments, which is the sum of the direction of its polarity. For water, the dipole moment is upward towards oxygen which is more electronegative. CO2 is a combination of carbon and oxygen atoms double bonded to the central carbon atom. Is the shape symmetrical in 3D? Yes A. Are all atoms bonded to the central atoms the same? Yes Therefore, CO2 is a nonpolar molecule. CH3Br is an organic molecule with the halogen bromine. A. Is the shape symmetrical in 3D? No Therefore, CH3Br is a polar molecule. The existence of lone pairs in the central atom has a big factor in making a molecule polar. For example, H2O (Bent) - polar due to two lone pairs and NH3 (Trigonal pyramidal) - polar due to one lone pair To summarize: Nonpolar molecules are symmetric with no unshared electrons. Polar molecules are asymmetric, either containing lone pairs of electrons on a central atom or having atoms with different electronegativities bonded. Activity 2 – Geometries and Polarities Directions: Given the following molecules and their respective molecular geometries, give the polarity of the molecule if its polar or non-polar. The first one is done for you. Molecule ∆ EN Bond Polarity Molecular Polarity of Molecule (refers to pairs) Geometry (considering the geometry) P – Cl Polar Non-polar 1. PCl5 |2.2 – 3.0| 0.8 17 2. BeCl2 3. CH2Cl2 4. OF2 5. SF6 Lesson Polarities and Properties 2.3 As you have learned on Lesson 1, there is a certain attraction to magnetic and electric fields for polar molecules and there is no reaction for non-polar molecules. This also has a relationship with their physical properties. Remember the rule: what happens microscopically (or atomically) has an effect macroscopically (physically). The table below shows the general properties of polar and nonpolar molecules. Polar molecules Nonpolar molecules Usually exist as solids or Usually exist as gases at liquids at room temperature room temperature High boiling point Low boiling point High melting point Low melting point High surface tension Low surface tension 18 Low vapor pressure High vapor pressure Low volatility High volatility Soluble in water Insoluble in water Examples: Water Examples: Waxes, Oils, Carbon Dioxide Solubility of Polar and Nonpolar Compounds Solubility is defined as the ability of a solid substance to be dissolved in a given amount of solvent while miscibility is the ability of the two liquids to combine or mix in all proportions, creating a homogenous mixture. The general rule to remember about the solubility and miscibility of molecular compounds can be summarized in a phrase, “like dissolves like” or “like mixes with like”. This means that polar substances will only be dissolved or mixed with polar substances while nonpolar substances will be soluble or miscible with another nonpolar substance. If we will go back to your simple experiment above, does the rule like dissolves like hold? Water is a polar compound, and it mixes and dissolves with polar compounds such as alcohol. However, it doesn’t dissolve well with oil, which is a nonpolar compound. Surface Tension Because water is a very special compound, having been polar, it exhibits a lot of special properties like surface tension. In the chemistry sense, surface tension is the energy required to increase the surface area by a unit amount. The stronger the forces are between the particles in a liquid, the greater its surface tension. Polar molecules have a higher surface tension than nonpolar molecules, due to the extra lone pairs that it has that connects it to other molecules of the same compound, also known as intermolecular forces, to be discussed in the next module. Look at the table of surface tension of different compounds at 20oC. Substance Surface Tension (J/m2) Diethyl ether (slightly polar) 1.7 x 10-2 Water (polar) 7.3 x 10-2 Water has the highest surface tension of all the liquids. This property is vital for surface aquatic life because it keeps plant debris resting on a lake surface which provides shelter and nutrients for fishes, microorganisms, and insects. Activity 3 – Like Dissolves Like Directions: Generally, like dissolves like. Polar molecules dissolve polar molecules and ionic compounds. Nonpolar molecules dissolve other nonpolar molecules. Alcohols, depending on what alcohol it is, have the characteristics of both, tend to dissolve both types of solvents, but do not dissolve ionic solvents. In the table below are solutes and solvents. Check the appropriate column as to whether the solute is soluble in a polar or a nonpolar solvent. 19 Solvents Carbon Solutes Water tetrachloride (non- Alcohol polar) NaCl (ionic) I2 (non-polar) Ethanol (alcohol) C6H6/benzene (non-polar) Toluene (non- polar) Activity: Directions: Compute the electronegativity differences of each pairs of elements and arrange them afterwards from the LEAST POLAR to the MOST POLAR. Element Pair Electronegativity Arrangement Difference C – O C – N C – Cl C – F C – H MODULE 3 This module will provide you the knowledge, skills and understanding of intermolecular forces of attraction. The module is divided into two lessons: Lesson 3.1 – Intermolecular Forces of Attraction Lesson 3.2 – Effects of Intermolecular Forces of Attraction At the end of this lesson, you are expected to: 1. Differentiate intramolecular forces from intermolecular forces of attraction. 2. Identify the intermolecular forces of attraction that exists between given molecules. 3. Relate intermolecular forces of attraction to different properties of matter such as viscosity, surface tension, and capillarity. 20 Lesson Intermolecular Forces of 3.1 Attraction Intramolecular vs Intermolecular Forces of Attraction From your English class, you know that prefixes have meaning. Inter- means between while Intra- means inside. Therefore, intramolecular forces of attraction are forces that exists between the elements in a compound while intermolecular forces are forces that exists between two molecules of a compound. Intramolecular forces are stronger than intermolecular forces of attraction. Intramolecular forces are also known as bonding forces. They are strong because they involve larger charges, and they are closer together. Intramolecular forces usually determine the chemical property of the compound. Intramolecular forces are more permanent than intermolecular forces of attraction. Intermolecular forces, on the other hand, are forces that form between molecules, atoms, or ions. They are relatively weak because they involve smaller charges that are far apart. Intermolecular forces usually determine the physical property of the compound. Intermolecular forces are also known as van der Waals forces. Let us use the following comparison to differentiate the strengths of intermolecular or intramolecular forces of attraction. H2O (g) → 2 H2 (g) + O2 (g) absorbs 927 kJ/mol (energy needed to break the bonds between hydrogen and oxygen atoms in water molecule) H2O (l) → H2O (g) absorbs 40.7 kJ/mol (energy needed to break the forces between water molecules to transform the liquid water to water vapor) We can see from the illustration above that it takes less energy to turn liquid to gas (breaking intermolecular forces) rather than separating hydrogen and oxygen in a water molecule (breaking bonding forces). The table below shows the comparison between bonding and nonbonding (intermolecular) forces of attraction. A. Bonding Forces (Intramolecular Forces) 21 B. Nonbonding forces (Intermolecular Forces) Source: Martin Silberberg’s Chemistry: The Molecular Nature of Matter and Change, pg 437 Types of Intermolecular Forces A. Ion – Dipole Forces This happens when an ion and a nearby polar molecule attract each other. The ion comes from an ionic compound that dissociated or separated. Ions, by definition, are charged particles. Remember the basic rule of physics, opposites attract. So, if you have a polar molecule with partially negative and partially positive ends, the partially negative end with attract to an ion that is positive; and the partially positive end will attract the ion that is negative. The most basic example of ion-dipole forces takes place when an ionic compound dissolves in water. The partially negative oxygen atom in water is attracted to the Na+ ion in NaCl while the terminal partially positive H from water is attracted to Cl- ions. Source: https://lh3.googleusercontent.com/proxy/Z6ZJEuQ1UStY73Wp1xvvYzBfs5FizSggCIm2B7wVa0lPWYE0hPpDz5sPYym- 9X2Dx6lKudiWQYMt1f7DIYHTl88Ib1eVYRn3-7NOX7VBfY8 B. Dipole-Dipole Forces When polar molecules lie near one another, their partial charges act as tiny electric fields that orient them and give rise to dipole-dipole forces, where the positive pole of one molecule attracts to the negative pole of the other molecule. You also need to take look at the molecular structure or geometry of the compound for you to take a look if the dipoles are available for intermolecular forces. Dipole – dipole forces exist between polar covalent molecules. 22 B.1. Hydrogen Bond – A special type of Dipole-Dipole Force A special type of dipole-dipole force arises between molecules that have a terminal H atom bonded to a small, highly electronegative atom with lone electron pairs, specifically fluorine (F), oxygen (O). and nitrogen (N). The F, O, and N must be terminal or exposed so that it can bond with other H from another molecule. This type of dipole-dipole force is a relatively strong form of molecular attraction. The oxygen atom should be exposed so that the partially positive H can interact with it. C. Dispersion (London) Forces Consider a neutral atom such as Helium or a nonpolar molecule such as H2. The electrons in each of the case are in constant motion relative to the nucleus so that in a given atom or molecule, the centers of positive and negative charges no longer coincide, and a temporary dipole is produced. Its direction could change with the movement of electrons. The electric field generated by a temporary dipole can induce a dipole in phase with itself in neighboring atoms and weak attractions set in between nearby induced dipoles. These temporary cohesive forces are known as London dispersion forces. These are the weakest intermolecular forces, and they are the most temporary. These forces are present between all particles, whether atoms, ions, molecules. Dispersion forces are the only force existing between nonpolar particles. However, because it exists in all particles, dispersion forces contribute to the overall energy of attraction of all substances. Therefore, all molecules exhibit London Dispersion Forces. D. Induced Dipole Forces A nearby electric field can distort the electron cloud of a different molecule, specifically a nonpolar molecule. In effect, the field induces a distortion in the electric cloud. For a nonpolar molecule, this distortion creates a temporary, induced dipole moment. For a polar molecule, it enhances its dipole moment that is already present. The ease with which the electron cloud of a particle can be distorted is called polarizability. There are two types of induced dipole forces: ion that induces a 23 nonpolar molecule to become a dipole, and a dipole molecule that induces a nonpolar molecule to become a dipole. D.1. Ion-Induced Dipole Forces Ion-induced dipole forces is a charged-induced dipole forces when an ion induces a nonpolar molecule so that it creates a temporary dipole in the nonpolar molecule. This is a temporary interaction only and after the ion leaves the system, the nonpolar molecule returns to being a nonpolar molecule. One of the best examples of ion-induced dipole is the interaction of Fe2+ ion in hemoglobin and the oxygen molecule (O2) in our blood. Oxygen molecule is a nonpolar compound. Because Fe2+ is an ion, it induces a dipole in the oxygen molecule, thus binds itself in the oxygen. Its role is for the delivery of oxygen to the bloodstream. D.2 Dipole-Induced Dipole Interaction Dipole-induced dipole interaction are weaker than the ion-induced dipole. These arise when a polar molecule distorts the electron cloud of a nearby nonpolar molecule. The best example of this is the solubility of oxygen in water. Paint thinners and grease solvents also function through dipole-induced dipole forces. Identifying Intermolecular Forces Between Molecules For you to correctly identify the intermolecular forces between the molecules, you need to correctly identify as well the polarity of that molecule. With that, you can determine the intermolecular forces that exists between the molecules. The schematic diagram below can help you to identify the intermolecular forces in a sample. Interacting Particles Ions Present Ions not Present Polar + Nonpolar Nonpolar Molecules Ion + Polar Molecule Polar Molecules Only Molecules (Dipole- Only (Dispersion forces (Ion-Dipole Forces) (Dipole - Dipole Forces) Induced Dipole Forces) ONLY) Ion + Nonpolar H bonded to N, O, F Molecule (Ion - Induced (Hydrogen Bonding) Dipole Forces) DISPERSION FORCES ARE PRESENT IN ALL TYPES OF MOLECULES 24 Strength of Intermolecular Forces We can now list down all the intermolecular forces and the strength that it exhibits. The stronger the intermolecular force, the higher energy it needs for us to break those forces. The weaker it is, the less energy needed. The chart below tells us of the relative strengths of intermolecular forces. This can be referred to when trying to relate the IMF to the properties of substances that will be discussed in the following lesson. Relative Strengths of Intermolecular Forces Ion-dipole Strongest H-bonding Dipole-dipole Dipole-induced dipole London dispersion forces Weakest Activity 1: Directions: Check the type of intermolecular forces that exists between each of the following molecules. Check ALL that applies. The first one is done for you. Molecules Ion – Dipole Dispersion Hydrogen Bonding Induced Induced Dipole – London Dipole- Forces Dipole Dipole Dipole Ion – 1. CO2 / 2. Na+ and 3. HBr 4. I2 5. HF 6. NH3 7. CH3OH 8. CH4 9. C8H10 10. O2 25 Lesson Effects of Intermolecular 3.2 Forces of Attraction Since intermolecular forces happens in all types of compounds, it affects the physical characteristics of the substance. Listed below are some of the properties of substances that are greatly affected by the presence of intermolecular forces. a. Boiling Point In its strictest sense, the boiling point of a substance depends on the equilibrium vapor pressure exerted by a liquid or solid above the liquid or solid. This means that a substance reaches its boiling point if its rate of condensation is equal to its rate of vaporization in a closed container. For example, at 100 oC, the vapor pressure of water is equal to the atmospheric pressure of 1.00 atm. Since the vapor pressure of water is equal to the vapor pressure of the environment, thus, boiling occurs. Vapor pressure of a substance, however, is dependent on the strength of the intermolecular forces present in that substance. When the intermolecular forces are strong, the vapor pressure is low. Consequently, boiling will happen in a higher temperature since it needs more energy to break the intermolecular bonds for the substance to change it into vapor. In essence, strong intermolecular forces produce lower rates of evaporation and a lower vapor pressure. Therefore, we need to raise the temperature for the vapor pressure to increase. The reverse is also true, weak intermolecular forces produces a higher rate of evaporation and a higher vapor pressure. Therefore, we need lower amounts of temperature for it to boil. Example: Select whether which of the following pairs of substances has the higher boiling point. 1. MgCl2 or PCl3? Answer: MgCl2 is an ionic compound, but PCl3 is a polar covalent compound. Since PCl3 is composed of polar compounds, the IMF present in it is dipole-dipole while in MgCl2, it is ionic bonding forces. The forces in MgCl2 is stronger – thus MgCl2 has a higher boiling point. 2. CH3NH2 or CH3F? Answer: Let us look at their structures. Checking their molecular masses, they have almost the same molecular mass. Since CH3NH2 has the N—H bonds, it means that it can form hydrogen bonds. CH3F, although having the F atom, is not connected to H, so dipole-dipole forces occurs but not hydrogen bonds. Therefore, CH3NH2 has the higher boiling point. 26 Boiling Point and Structures For both nonpolar compounds that are considered isomers or with the same molecular formula, the strength of the dispersion forces (which ultimately affects their boiling points) is influenced by their molecular shape. Take into example n- pentane and 2,2-dimethylpropane which has the same molecular formula C 5H12. The figure on the left is n-pentane, and on the right is 2,2-dimethylpropane. We can see that n-pentane has more contact points (shaped like a cylinder) rather than 2,2-dimethylpropane (compact shape). Thus, n-pentane has more dispersion forces, and a higher boiling point. Laboratory data shows that the boiling point of n- pentane is 36.1oC while 2,2-dimethylpropane has a boiling point of 9.5oC. Boiling Points, Size, and Polarizability The strength of dispersion forces also depends on the size of the substance or the number of electrons in the substances. The ease with which the electron distribution is distorted explains the amount of dispersion forces that a substance exhibits. The distortion of the electron distribution is known as polarizability. The greater the polarizability of the electron distribution the greater are the dispersion forces. When the dispersion forces are high, the boiling and melting points are also high. Br2 and F2 are both diatomic gases. They are also both nonpolar, but Br 2 is a bigger molecule than F2. The polarizability of Br2 is greater than F2 so it has greater dispersion forces. This explains why Br2 has a higher boiling point than F2. Greater amount of energy is needed to overcome the big dispersion forces in Br 2 than in F2. b. Melting Point The condition given above for boiling point is also true for melting point. When bond breaks easily, it affects the melting point of the substance. The greater the intermolecular forces, the greater the temperature it needed for it to melt. c. Volatility Volatility is the tendency of substances to evaporate at normal temperatures. For example, rubbing alcohol is volatile because it evaporates rapidly in our hands while cooking oil does not. The more volatile a substance, the weaker is its intermolecular force because it means that it just need less energy for the substance to turn into vapor. d. Viscosity Viscosity is the measure of a liquid’s resistance to flow. When a liquid flows, the molecules slide around and past each other. When a material is viscous, it means that it has intermolecular forces that impedes its motion. It means that liquids that have stronger intermolecular forces is more viscous than other. e. Capillarity and Surface Tension Surface tension is the amount of energy required to stretch the surface area of a liquid. Liquids with high intermolecular forces have higher surface tensions. It means that it has a stronger network that holds the liquid together. For example, you can put a lot of drops of water in a one-peso coin and still hold the shape of the 27 liquid before it breaks. However, if the water is soapy, it means that there is an interaction of nonpolar compounds such as soaps that breaks that surface tension. Surface tension is also the reason why paperclip can float in water. It is also the reason why mosquitos can lay their eggs on water without sinking. Because surface tension happens, capillarity happens as well. It is the rising of a liquid through a narrow space against the pull of gravity. Two forces, cohesion and adhesion push surface tension on walls. Cohesion is the intermolecular attraction between like molecules while adhesion is the attraction between unlike molecules. To illustrate, the attraction between the water molecules is cohesion, while the attraction between the water molecules and the sides of the glass tube is adhesion. Source: https://ib.bioninja.com.au/_Media/water-cohesion-and-adhesion_med.jpeg If cohesion is greater than adhesion, there will be depression or lowering. If there is lowering, there is a lower height of the liquid in the capillary tube. The meniscus or the curve at the top of the liquid inside the glass tubing is curving downwards or convex meniscus. If adhesion is greater than cohesion, there will be rising of the liquid in the tubing, and the meniscus will be curving upwards or called as concave meniscus. Look at the illustrative example below: Source: https://www.quirkyscience.com/wp-content/uploads/2018/02/Figure-1-1.png Glass is mostly silicon dioxide (SiO2), so the water molecules form hydrogen bonds with the oxygen of the tube’s inner walls. Because the adhesive forces between the water molecules (hydrogen bonding) is greater than the cohesive forces between the water and the glass tubing, water creeps up to the wall of the tubing. The cohesive forces now give rise to surface tension and then pulls the liquid surface upward. These forces (cohesive and adhesive) combine to raise the water level and thus produces the concave meniscus. This explains the action why water rises from the roots of the plants to the leaves through the tubes inside the plant called xylem and phloem. In the case of mercury, since mercury (Hg) has stronger cohesive forces (due to metallic bonding). We know that metallic bonding is stronger than the cohesive forces that will happen to Hg and the glass. Therefore, the liquid Hg tends to pull 28 away from the walls. At the same time, the surface atoms are being pulled towards the interior of the Hg by its high surface tension, so the level drops. These combined forces produce a convex meniscus. This is the principle behind the barometer. The stronger the intermolecular forces possessed by a molecule, the higher is the surface tension of the substance. Activity 2: Directions: Fill in the blanks with the words GREATER or LOWER. 1. The stronger the forces between the particles, the _____________ the melting point. 2. The stronger the forces between the particles, the _____________ the vapor pressure. 3. The stronger the forces between the particles, the ______________ the viscosity. 4. CH4 has a _______________ surface tension than H2O. 5. SiO2 has a _______________ boiling point than SO2. 6. HCl has a _______________ boiling point than I 2. 7. C2H6 has a _______________ vapor pressure than C 4H10. 8. CH3CH2OH has a _______________ boiling point than CH3CH2CH3. 9. CH3CH2OH has a _________________ viscosity than CH3COCH3 10. H2O has a ____________ intermolecular force than CO2. Activity: World of Water Directions: Water is made from two gases that are flammable but together they make a substance that extinguishes flame. These two elements, hydrogen, and oxygen, when bonded together allowed life on earth to flourish. The ability of water to form hydrogen bonds presents many interesting properties which are useful to life. Create an output, whether a (1) poster, (2) video presentation, (3) essay, or (4) collage, or anything that you can propose to your teacher showing what you have researched on the following topic options about the amazing world of water. Choose one topic and relate intermolecular forces of attraction. I. Water and the Human Body II. The triple point of water III. Water and its role in agriculture IV. Water and the “universal solvent” claim V. Water and religion and myths VI. The shape of water and the beauty of snow VII. Water and its high specific heat capacity 29 MODULE 4 This module will provide you the knowledge, skills and understanding of intermolecular forces of attraction. Explain how the structures of biological macromolecules such as carbohydrates, lipids, nucleic acid, and proteins determine their properties and functions. At the end of this lesson, you are expected to: 1. Recall the different types of biomolecules from previous Science lessons. 2. Relate the structure of biomolecules to the intermolecular forces present in the biological macromolecule. 3. Cite the importance of these intermolecular forces in the properties and functions of the biomolecules. Lesson Biological Macromolecules 4 Carbohydrates Carbohydrates can be represented by the stoichiometric formula (CH 2O)n where n is the number of carbons in the molecule. Therefore, the ratio of carbon to hydrogen to oxygen in the compound is 1:2:1. Carbohydrates can be classified into three: monosaccharides, disaccharides, and polysaccharides. Carbohydrates act as energy storage or food reserves in plants and animals. This property of carbohydrates is due to many carbon-hydrogen bonds that exists in the molecule. Breaking down carbon-hydrogen bonds easily releases energy faster. Because of the abundance of polar -OH groups in the molecule, they are highly polar. This makes carbohydrates soluble in many body fluids, such as blood, which is 70% water. This makes carbohydrates easily available to the body. 30 From: https://alevelbiology.co.uk/wp-content/uploads/2019/11/Structure-and-Function-of- Carbohydrates_1.png Carbohydrates can be represented either via Fischer or Haworth projections. The illustration above is an example of a Fischer projection of a carbohydrate. We can classify carbohydrates as to the functional group that it possesses, if it is an aldose or ketose. It is an aldose if it contains an aldehyde. It is a ketose if it contains a ketone. It can be easily seen in the structure. An aldehyde is a terminal functional group (R-C=O -H) while a ketone is not a terminal functional group (R-C=O-R). When the linear structure (Fischer projection) closes, it produces a Haworth projection of the molecule, in this case, glucose molecule. This is the cyclic structure of a carbohydrate which you usually see in books. Carbohydrates can also be classified by the number of molecules (saccharides) it has either monosaccharide (one), disaccharide (two), or polysaccharide (many). Monosaccharide (one saccharide) Glucose used in dextrose, blood sugar; the form utilized by the human body Galactose found in milk and milk products Fructose found in fruits and honey Disaccharides (two saccharides) Maltose glucose + glucose found in malt Sucrose glucose + fructose found in regular table sugar, sugarcane, and sugar beet Lactose glucose + galactose found in milk and milk products Polysaccharides (many saccharides) Amylose storage form of glucose in plants (starch) Amylopectin storage form of glucose in plants (starch) Glycogen storage form of glucose animal; stored in the liver and muscles 31 Cellulose structural material in plants--cell wall in wood, wood fiber cannot be digested by humans You may notice that glucose, fructose and galactose have the same chemical formula, C6H12O6. However, they have different functions. It is because they may have the similar chemical formula but different chemical structure. This is what we call isomers. Glucose and galactose are really stereoisomers – same order, same aldose, but different arrangement in space, while fructose is a structural isomer of glucose and galactose. Their difference is that glucose and galactose are aldoses, while fructose is a ketose. Disaccharides form by dehydration reaction which eliminates water molecule in the combination of two monosaccharides. The bond that is formed between two monosaccharides is called a glycosidic bond/linkage. Long chains of monosaccharides result in polysaccharides. Below are the structures of amylose and amylopectin. The structures of amylose and amylopectin affects food chemistry. For example, rice with high amounts of amylopectin will be very sticky once it is cooked. If the rice grains is high in amylose, it will be fully separated once cooked. If we have excess starch in our body, our body stores it in forms of glycogen so that if we lack energy, there is an immediate source of energy stored. Look at the illustrations below. Can you think of the reason why glycogen is branched and how will it help the fast release of energy? What do you think is the reason why cellulose is very strong? Lipids A fat molecule consists of two main components: glycerol and fatty acids. Glycerol is an alcohol with three carbons, five hydrogens, and three hydroxyl (OH) groups. Fatty acids have a long chain of hydrocarbons with a carboxyl group attached and may have 4-36 carbons; however, most of them have 12-18. In a fat molecule, the fatty acids are attached to each of the three carbons of the glycerol molecule with an ester bond through the oxygen atom. The bond formed is called ester bonds and the process is esterification (condensation reaction). There are different classifications of lipids: triglyceride, phospholipid, wax, and steroid. The lipid family is one of the most varied in terms of structure, but they share the common property of being insoluble in water. Fat and oil are the most common examples of lipids. They are under triglycerides because they are composed of glycerol and three fatty acids. Fat refers to solid triglyceride usually from animal sources such as meat, milk, butter, margarine, eggs, and cheese. Oil refers to liquid triglycerides from plant sources. Examples are olive oil, corn oil, sunflower oil, and soybean oil. Animal fats contain high percentages of saturated fatty acids while plant oils are mostly unsaturated fatty acids. 1. Saturated fats have two carbons attached to each carbon (except the one at the end). Saturated fats are unhealthy fats like butter. They are usually solid at room temperature. In blood chemistry, they are called LDL or low-density lipoprotein. 2. Unsaturated fats contains a double bond. They are missing at least one hydrogen and are curl in shape. The unsaturated fats are healthy and include oils. In blood chemistry, they are under HDL or high-density lipoprotein. 32 They are insoluble in water because of their lack of many polar and hydrogen bonding functional groups. When placed in water, lipid molecules cling together, exposing their polar groups to their surrounding molecules while their nonpolar groups stay within the interior of their lipid cluster. Because of this property as well, lipids are at best an effective cell membrane component. Proteins Proteins are composed of four elements, namely: carbon, hydrogen, oxygen, and nitrogen. Sulfur and other metals are sometimes also found in proteins. If carbohydrates are made up of saccharides, proteins are made up of amino acids. There are twenty amino acids that exists in living organisms. Nine of these amino acids are essential. These are histidine, isoleucine, leucine, lysine, methionine, phenylalanine, threonine, tryptophan, and valine. It means that it cannot be made by our body and must come from food. Amino acids bond to other amino acids via peptide bonds or linkages. The 20 amino acids that make up proteins have side groups with varying properties. Hence, the number and sequence of amino acids affect the properties and functions of a particular protein. For example, hemoglobin, the protein found in red blood cells as is used to carry oxygen has 574 unique amino acid sequences. If one of the proteins is wrong in the sequence, the function malfunctions. Insulin, on the other hand, has 51 amino acid sequences. Protein structures are very complex, and researchers have only very recently been able to determine the structure of complete proteins down easily and quickly to the atomic level. (The techniques used date back to the 1950s, but until recently they were very slow and laborious to use, so complete protein structures were very slow to be solved.) Early structural biochemists conceptually divided protein structures into four levels to make it easier to get a handle on the complexity of the overall structures. To determine how the protein gets its final shape or conformation, we need to understand these four levels of protein structure: primary, secondary, tertiary, and quaternary. Because of different intermolecular attractions that happen in the tertiary level of protein structures, the functions of the proteins also change. The picture below shows the different intermolecular forces that may happen in the tertiary protein structure. 33 Primary structure is the amino acid sequence. Secondary structure is local interactions between stretches of a polypeptide chain and includes α-helix and β-pleated sheet structures. Tertiary structure is the overall the three-dimension folding driven largely by interactions between R groups. Quaternary structures is the orientation and arrangement of subunits in a multi-subunit protein. Because of its structure, proteins have various functions. Below is a table to summarize it and some of its examples. Type Examples Functions Digestive Amylase, lipase, pepsin, Help in digestion of food by Enzymes trypsin catabolizing nutrients into monomeric units Transport Hemoglobin, albumin Carry substances in the blood or lymph throughout the body Structural Actin, tubulin, keratin Construct different structures, like the cytoskeleton Hormones Insulin, thyroxine Coordinate the activity of different body systems Defense Immunoglobulins Protect the body from foreign pathogens Contractile Actin, myosin Effect muscle contraction Storage Legume storage proteins, Provide nourishment in early egg white (albumin) development of the embryo and the seedling Two special and common types of proteins are enzymes and hormones. Enzymes, which are produced by living cells, are catalysts in biochemical reactions (like digestion) and are usually complex or conjugated proteins. Each enzyme is specific for the substrate (a reactant that binds to an enzyme) it acts on. The enzyme may help in breakdown, rearrangement, or synthesis reactions. Enzymes that break down their substrates are called catabolic enzymes, enzymes that build more complex molecules from their substrates are called anabolic enzymes, and enzymes that affect the rate of reaction are called catalytic enzymes. It should be noted that all enzymes increase the rate of reaction and, therefore, are considered to be organic catalysts. An example of an enzyme is salivary amylase, which hydrolyzes its substrate amylose, a component of starch. Hormones are chemical-signaling molecules, usually small proteins, or steroids, secreted by endocrine cells that act to control or regulate specific physiological processes, including growth, development, metabolism, and reproduction. For example, insulin is a protein hormone that helps to regulate the blood glucose level. Proteins have different shapes and molecular weights; some proteins are globular in shape whereas others are fibrous in nature. For example, hemoglobin is a globular protein, but collagen, found in our skin, is a fibrous protein. Protein shape is critical to its function, and this shape is maintained by many different types of chemical bonds. Changes in temperature, pH, and exposure to chemicals may lead to permanent changes in the shape of the protein, leading to loss of function, known as denaturation. Nucleic Acids 34 Nucleic acids play an essential role in the storage, transfer, and expression of genetic information. Nucleic acid was discovered by a 24-year-old Swiss physician named Friedrich Miescher in 1868. He was puzzled that an unknown substance in white blood cells did not resemble carbohydrates, proteins, or lipids. He was able to isolate the substance from the nucleus and initially called it nuclein. He eventually was able to break down nuclein into protein and nucleic acids. He found out that nucleic acids contain carbon, hydrogen, oxygen, nitrogen, and phosphorus The most common examples of nucleic acids are DNA (deoxyribonucleic acid) and RNA (ribonucleic acid). DNA is a nucleic acid that carries the genetic code of organisms. It is fondly termed as the blueprint of life. RNA, on another hand, carries the information from the DNA to the cellular factories for the synthesis of proteins. If carbohydrates are composed of saccharide units, proteins of amino acids, and lipids of fatty acids, nucleic acids are composed of nucleotides. Nucleic acids are also known as polynucleotides. Nucleotides are made up of a nitrogeneous base, a five-carbon sugar, and a phosphate group. The nitrogeneous base can be a pyrimidine or purine. Purines include adenine and guanine, while pyrimidines are cytosine, thymine, and uracil. This plays an important part in the zipping of the DNA since a purine base pair with a specific pyrimidine base through hydrogen bonding. Adenine bonds with thymine with two hydrogen bonds, and guanine with cytosine with three hydrogen bonds. It is important that they are strong but NOT permanent. The double helix structure of the DNA protects the nonpolar nitrogenous bases in the molecules by orienting them in the middle. The polar phosphate groups are exposed so that the DNA will be soluble in the aqueous polar environment. This protects the information stored in the DNA ensuring that the DNA sequence stays intact by keeping the DNA structure stable. Furthermore, the helix is held together by hydrogen bonds that forms the two strands of the DNA. This allows to form a stable double helix and thus be able to protect the important genetic information that makes up the body. The nucleotides are joined together using phosphodiester bonds that makes the backbone of the DNA. Source: https://www.researchgate.net/profile/A_Turut/publication/283467276/figure/fig1/AS:613983497232405@1523396481947/Double-helix- structure-of-DNA.png Activity 1: Macromolecule Comparison Table Directions: Complete the table with correct information based on your lesson above. 35 Biological Function/s Monomer Bonds Present Examples Macromolecule (subunit) Carbohydrates Lipids Proteins Nucleic Acids Activity – The Biological Macromolecules in My Food Modified from: https://lhsblogs.typepad.com/files/macromolecules-in-your-food.pdf The Nutrition Facts Label tells you what nutrients (components of food your body needs to grow and stay healthy) and how much of those nutrients are in found in one serving. The Nutrition Facts label can help you make choices about the food you eat. The Nutrition Facts label is on the outside of most food packages but isn't on most fresh foods (like fruits and vegetables). Below is an example of a Nutrition Facts label and explanations of the information found on the label. Find the food label of something you have eaten (or would consider eating) and complete the following information. 1. What are the main ingredients of this food? 2. Identify the total grams of: a. Carbohydrates: _______________ b. Fats: ____________ c. Proteins: _______________ 3. What surprised you as you read the nutrition facts for this food item? Do the steps to at least three food labels. MODULE 5 This module was designed and written with you in mind. It is here to help you to look deeper into chemical reactions. The scope of this module permits it to be used in many different learning situations. The language used recognizes the diverse vocabulary level of students. The lessons are arranged to follow the standard sequence of the course. But the order in which you read them can be changed to correspond with the textbook you are now using. The module is devoted in a single lesson, namely: Lesson 5 – Rates of Chemical Reactions After going through this module, you are expected to: 1. describe how chemical reactions occur; 2. explain the Collision Theory and it’s premises; 3. illustrate and interpret energy diagrams; 4. explain the different factors that affect the rate of reaction. 36 Lesson Rates of Reaction 5 Chemical Reactions Occur One evidence that a chemical reaction has occurred is the production of other substances as shown by their properties which are different from those of the original substances. Chemical reactions happen because molecules interact with each other and collides with each other effectively to form a new product. This is called the collision theory. Collision Theory The collision theory explains how a chemical reaction takes place. According to this theory, two conditions must be satisfied for a chemical reaction to occur: 1) Particles of reactants must collide with one another in correct orientation; and 2) Colliding particles must have sufficient energy. No reaction can take place between two particles if they are far apart. They must come in contact so that they may be able to break bonds, exchange atoms, and form new bonds. Remember that chemical reactions involve bond breaking and bond forming in the intramolecular level, that is inside the molecule. Therefore, right amounts of energy and right orientation is needed. Let us take this illustration as an example: Source: https://mrtremblaycambridge.weebly.com/uploads/9/7/8/8/9788395/___5001988_orig.jpg Nitrous oxide (NO) and ozone (O3) reacts with each other. If the orientation of the molecules is incorrect, the collision will be ineffective. If there is an ineffective collision, the reaction will not proceed. However, if the orientation of the molecules is correct (meaning, they are in the right positions), the collision will be effective, and there will be new products, NO2 and O2. Let us have another illustration to show you the importance of the sufficient amount of energy for the reaction to proceed. 37 Source: https://saintschemistry10.weebly.com/uploads/5/1/9/3/51932861/gw500h283_orig.jpg In the first situation, the two different molecules are in the correct orientation. However, the reactants are moving too slow. When molecules move too slow, it means that their energies are not enough. Related to this is their kinetic energy. If there is less kinetic energy, molecules move slower. If there is not enough energy, the reactants will just bounce, and no reaction will happen. The energy needed for a reaction to proceed is known as its activation energy. Energy Diagrams Most reactions involving neutral molecules cannot take place at all until they have acquired the energy needed to stretch, bend, or otherwise distort one or more bonds. This critical energy is known as the activation energy of the reaction. Activation energy diagrams of the kind shown below plot the total energy input to a reaction system as it proceeds from reactants to products. The x-axis shows the direction of the reaction or the time it takes for the reaction to take place, and the y- axis shows the energy needed for the reaction to proceed. Source: https://upload.wikimedia.org/wikipedia/commons/6/6a/212_Enzymes-01.jpg Let us analyze the first energy diagram. All molecules have inherent energy. We call that energy potential energy. That is the reason why the reactants do not start with zero energy. Remember, energy is neither created nor destroyed. So, when we want to combine two reactants, we need to activate the reactants for them to form new bonds and to break their old bonds. We call that energy activation energy. The difference between the energy of the reactants versus the energy of the product is called enthalpy. In the strictest chemistry sense, enthalpy is the sum of the internal energy and the product of pressure and volume. We can see here that the energy of the reactants is higher than the energy of the products. Therefore, is energy released or is energy absorbed by the product? Energy is released. We call these reactions exothermic reactions. The prefix exo- means release. 38 Source: https://upload.wikimedia.org/wikipedia/commons/3/3a/Activation_Energy-c3bc.png The diagram above now shows the reaction of H2 and Cl2 to form HCl. Again, it requires correct orientation and the enough energy for the reaction to proceed. Is the energy of the reactants less or more than the energy of the products? The energy of the products is more than the energy of the reactants. Therefore, energy is absorbed by the products. We now call these reactions endothermic reactions. Shown below is a comparison chart between endothermic and exothermic reactions. Comparison Chart BASIS FOR ENDOTHERMIC REACTIONS EXOTHERMIC REACTIONS COMPARISON Meaning Chemical reactions Chemical reactions where involving the use of energy the energy is released or at the time of dissociation evolved in the form of to form a new chemical heat is known as the bond is known as the exothermic reaction. endothermic reaction. Energy The endothermic process The exothermic process requires energy in the form evolves or releases in the of heat. form of heat. Enthalpy ΔH is positive, as heat is ΔH is negative, as heat is (ΔH) absorbed. evolved. Examples 1. Conversion of ice into 1. Formation of ice from water vapor through water. boiling, melting or 2. Burning of coal evaporation. (combustion). 2. Breaking of the gas 3. The reaction between molecules. water and the strong 3. Production of anhydrous acid. salt from hydrate. 39 Factors Affecting Rate of Reaction A. Concentration According to the collision theory, there must be collision between reacting particles so that a reaction can occur. An increase in the concentration of the reactant means that there will be more particles colliding with each other in each time interval, thereby increasing the chances that a chemical reaction occurs. Therefore, an increase in the concentration of reactants causes an increase in the rate of the reaction. In the illustration above, if there are more particles of the reactants inside a container, there is more chance that they will collide with each other successfully. Concentration of each reactant can be expressed through percent by volume, percent by mass, molarity, molality, and other units. It can also be related to the reactants mass or volume. B. Pressure Increasing the pressure inside a container without changing the amount of substance inside decreases the volume of the substance (by virtue of the gas laws). The molecules will have less space to move and are more likely to collide successfully. Pressure is expressed in pascals, atm, or mmHg. The effect of pressure in reaction rates is very evident in gases as it follows laws such as the Dalton’s Law of Partial Pressures. C. Temperature Temperature is directly proportional to kinetic energy. An increase in temperature results in an increase in the kinetic energy of the reacting particles. The increased speed of the particles causes a greater proportion of these collisions to takes place in a given time. So, for example, with the reactants at 0 oC and at 100oC, what temperature will the reaction proceed faster? At 100 oC as it supplies activation energy quickly for the reaction to proceed. As a rule of thumb, rate of reactions roughly doubles or triples for every 10 Co rise in temperature. Source: https://cnx.org/resources/99129ce8ae5f998f706b70f2ed583002b106868e/CG12C3_004.png 40 D. Particle Size/Surface Area Particle size has something to do with the rate of reaction. Sugar has two forms, there is powdered sugar and there are sugar cubes. When you do coffee, which do you think will have the faster dissolution? The powdered sugar or the sugar cubes? The powdered sugar will dissolve faster. Making the particle size of the sugar smaller increases the surface area exposed which consequently results in an increase in contact area. Similarly, a reaction between two immiscible liquids can only take place at the interface between them. This interface can be increased by violent stirring which will create more areas for interaction and, therefore, lead to an appreciable increase in the rate of reaction. Another example, kindling wood burns faster than a whole log of wood. This is because there is a larger surface area in contact with oxygen in kindling wood rather than in a log of wood. Kerosene or gaas that is exposed over a wider area will burn faster than the same amount of kerosene that is just confined in a drum. The particle size and the surface area have a large effect on the rate of reaction. Source: https://www.pathwayz.org/Node/Image/url/aHR0cHM6Ly9pLmltZ3VyLmNvbS9aejZnUmE5LnBuZz8x E. Presence of a Catalyst A catalyst is usually a substance that, when added to a reaction mixture, increases the rate of the reaction but is itself unchanged after the reaction is completed. However, there also substance that slow down reactions. They are often referred to as inhibitors. Let us analyze the following energy diagram for a reaction with and without a catalyst. Source: https://upload.wikimedia.org/wikipedia/commons/f/fe/Carbonic_anhydrase_reaction_in_tissue.svg What happens to the activation energy when a catalyst is present in the reaction? The activation energy is lower. Therefore, the presence of a catalyst 41 speeds up a chemical reaction by lowering its activation energy. In this example, the catalyst is the enzyme carbonic anhydrase. Enzymes in our body acts as catalysts to speed up chemical reactions. Activity 1: Reaction Rates Directions: The following statements describe the factors that affect the rate of chemical reactions. Use the word bank below to complete the statements. WORDBANK catalyst heat concentration energy collision surface area temperature rate of reaction

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