SDO Navotas Sci9 Q2 Lumped (PDF) - Grade 9 Science

DIVISION OF NAVOTAS CITY 9 SCIENCE Quarter 2 S.Y. 2021-2022 NAVOTAS CITY PHILIPPINES Science – Grade 9 Alternative Delivery Mode Quarter 2 Second Edition, 2021 Republic Act 8293, section 176 states that: No copyright shall subsist in any work of the Governme...

DIVISION OF NAVOTAS CITY 9 SCIENCE Quarter 2 S.Y. 2021-2022 NAVOTAS CITY PHILIPPINES Science – Grade 9 Alternative Delivery Mode Quarter 2 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: Mark Anthony U. Andrade, Ronalyn A. Bautista, Jean May P. Buluran, Veronica A. Cane, Ronnie Mar Clemente, Aussie Claire A. Cruz, Ian Ezekiel C. Martin, Raquel G. Nachor, Claudine Alyssa T. Pangilinan, James Oliver K. Pingol, Lawrence Jay S. Santos, and June Kathleen A. Sayo Editors: Jean May P. Buluran and Aussie Claire A. Cruz Reviewers: Gemma C. Payabyab Illustrators: Mark Anthony U. Andrade, Jean May P. Buluran, Aussie Claire A. Cruz, Ian Ezekiel C. Martin, James Oliver K. Pingol, Lawrence Jay S. Santos, and June Kathleen A. Sayo Layout Artist: Jean May P. Buluran and Aussie Claire A. Cruz 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 Printed in the Philippines by ________________________ 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 What I Know................................................................................1 Module 1......................................................................................2 Module 2......................................................................................7 Module 3......................................................................................11 Module 4......................................................................................18 Module 5......................................................................................21 Module 6......................................................................................28 Module 7......................................................................................35 Assessment..................................................................................41 Answer Key..................................................................................42 References...................................................................................45 Multiple Choice: Choose the letter of the best answer. 1. Which of the following best describes the quantum mechanical model of the atom? A. It describes an electron's probability distribution that determines the probable location of an electron. B. It makes predictions based on Schrodinger’s mathematical equation. C. It is the currently accepted atomic model. D. All of the above are correct. 2. Chemical compounds can be classified as ionic or covalent compound, and these chemical compounds can be written in their chemical symbols. Which of the following is incorrectly paired? A. NaCl – ionic compound C. BrO – covalent compound B. F2 – ionic compound D. NH3 – covalent compound 3. How is the bond in Cl2 different from the bond in MgF2? A. The bond in Cl2 is metallic while the bond in MgF2 is covalent. B. The bond in Cl2 is ionic while the bond in MgF2 is covalent. C. The bond in Cl2 is covalent while the bond in MgF2 is ionic. D. There is no bond difference between the two. 4. Which of the following statements about organic compounds is/are true? I. Organic compounds contain calcium. II. Organic compounds contain carbon. III. Organic compounds can be produced by living organisms. IV. Organic compounds can be produced artificially. A. I, II, and III C. II and III B. I and III D. II, III, and IV 5. To which group does the molecule with the structure H―C≡C―H belong? A. alcohol B. alkane C. alkene D. alkyne 6. Acetone is used to clean the surface of the nail and to remove dirt on it. To which of the following functional groups does acetone belong? A. alcohol B. carboxylic C. ester D. ketone acid 7. How much mass is present in a mole of N2 gas if the atomic mass of N is 14 g? A. 14 g B. 28 g C. 6.02 x 1023 atoms D. 1.204 x 1024 atoms 8. Which of the following samples has the largest mass, in grams? A. 2 moles of CO B. 4 moles of H2O C. 3 moles of CO D. 5 moles of H2 9. Which of the following best explains the difference between empirical and molecular formulas? A. Empirical formula refers to the simplest whole number ratio of atoms; Molecular formula refers to the actual number of atoms. B. Empirical formula refers to the actual number of atoms; Molecular formula refers to the simplest whole number ratio of atoms. C. Empirical and molecular formulas are the same, they both refer to the simplest whole number ratio of atoms in a molecule. D. Empirical and molecular formulas are the same, they both refer to the actual number of atoms in a molecule. 10. What is the percentage of oxygen in carbon dioxide (CO2)? A. 27.3% B. 40.0% C. 70.3% D. 72.7% 1  MODULE 1 This module was designed and written with you in mind. It is here to help you master the Electronic Structure of Matter. 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 three lessons, namely: Lesson 1.1- Quantum Mechanical Model of an Atom Lesson 1.2- Electron Configuration Lesson 1.3- Quantum Numbers After going through this module, you are expected to: 1. predict the probable location of electron/s in an atom (electron cloud, Heisenberg's Uncertainty Principle); 2. describe and write the correct electron configuration of given elements; 3. describe the set of quantum numbers and complete the given set of quantum numbers for each given element; and 4. supply the following data from the electron configuration such as: period number, group number, number of paired and unpaired electron/s, number of valence electron/s, and number of core electrons. Lesson Quantum Mechanical Model 1.1 of the Atom The development of a better model of the atom was led by three physicists, namely Louie de Broglie, Erwin Schrodinger, and Werner Karl Heisenberg. De Broglie proposed that the electron, which has been known as a particle, could also be thought of as a wave. Schrodinger used mathematical equations to describe the probability of finding an electron in a certain position. On the other hand, Heisenberg discovered that for a very small particle like the electron, its location cannot be exactly known and how it is moving. This is called the uncertainty principle. These scientists believed that there is only a probability that electrons can be found in a three-dimensional space around the nucleus known as atomic orbitals. Therefore, the quantum mechanical model describes the probable location of electrons within the atom using atomic orbitals. These orbits Table 1.1.1. Energy Level, Sublevel and Orbitals are called levels and they are numbered 1, 2, 3, 4, and so forth, with the first level being the orbit closest to the nucleus. The levels can be broken down into sublevels. They are the following: s (sharp), p (principal), d (diffuse), and f (fundamental) sublevels. Level one has one sublevel – an s. Level 2 has 2 sublevels, s and p. Level 3 has 3 sublevels – s, p, and d, and Level 4 has 4 sublevels – s, p, d, and f. 2  The sublevels contain orbitals. These are spaces that have a high probability of containing an electron. Orbital can also be described as an area where electrons live. There can be a maximum of two electrons in one orbital. Since the s sublevel has just one orbital, it can only contain a maximum of 2 electrons. The p sublevel has 3 orbitals and can contain up to 6 electrons. The d sublevel has 5 orbitals and is intended for a maximum of 10 electrons. Finally, the f sublevel has 7 orbitals and can hold a maximum of 14 electrons. Activity 1.1: Let’s Count! Using the information on Table 1.1.1, answer the following questions. 1. How many sublevels are there on the 4th level? 2. How many electrons can be probably found on the 3rd level? 3. How many orbitals are there on the 2nd level? 4. How many sublevels are there in the 3rd energy level? 5. How many electrons can be found in a p sublevel? 6. How many orbitals in an f sublevel? 7. How many electrons are there in an f sublevel? 8. How many electrons can be found on the 1st level? 9. How many electrons can fit on the 2nd level? 10. How may electrons can be found in the fourth energy level? Lesson 1.2 Electron Configuration The arrangement of electrons outside the atom, or the distribution of electrons outside the nucleus, is referred to as electron configuration. This will help us to understand and predict more about the properties of elements. The electron configuration of the element Helium illustrates that there are two protons in the middle, which have a greater amount of positive charge to pull the electrons in its orbit Electron configuration is a method of distributing electrons in different sublevels or orbitals as well as in the main energy levels. A typical electron configuration consists of numbers, letters, and superscripts with the following format: 1. A number indicates the energy level 2. A letter indicates the type of sublevel or orbital (s, p, d, f) 3. A superscript indicates the number of electrons in the orbital Example 1: o 1s1 means there is 1 electron in the ‘s’ orbital of the first energy level of the Hydrogen atom. o The element is HYDROGEN because it has an atomic number of 1. o The atomic number is equal to the number of protons and electrons in a neutral atom. Example 2: 1s22s22p63s23p64s23d6 = no. of e-: 26 → atomic no.: 26 → Element: Iron Filling the Electron Subshells Configuration The order of electrons fills the lowest energy orbitals first. The subshells are given by “s”, “p”, “d”, and “f” and so on. The numbers at the bottom indicate the number of electrons that can fit into that type of sublevel. For example, at the “s” sublevel, it can accept only two electrons. 3 Filling of Atomic Orbitals Aufbau Principle This principle is named after the German word "Aufbeen", which means "build up". The Aufbau principle dictates that electrons will occupy orbitals with lower energies before occupying higher energy orbitals. The energy of an orbital is the sum of the principal and the AQN. According to this principle, electrons are filled in the following order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p… The order in which electrons are filled in atomic orbitals as per the Aufbau Principle is shown in the diagram on the right. Table 1.1.2. Maximum number of electrons and spins Pauli Exclusion Principle An orbital may be occupied by a maximum of two (2) electrons. Each electron will have a different spin. The maximum number of electrons in each of the sublevels and orbitals is shown in Table 2. Hund’s Rule When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron, then a second electron is added to each orbital. The table on the right illustrates how electrons occupy their energy orbitals using the Hund's Rule. Short electron configurations are often written in standard notation, particularly for elements with a high atomic number. Instead of using conventional notation, abbreviated or compressed notation may be used. As illustrated in the examples, the sequence of filled subshells that corresponds to a noble gas's electronic configuration is replaced with the noble gas's symbol in square brackets. Activity 1.2: The Electron Configuration and Orbital Diagram of Elements: Following the principles of writing the electron configuration of elements, fill in the table below with the information needed. Symbol Atomic Number Electron Configuration Ar - Argon K - Potassium Al - Aluminum O - Oxygen N - Nitrogen Si - Silicon Activity 1.3: The Table of Configurations: Complete the table below. Atomic Atomic Shorthand Element Expanded notation Symbol number notation Magnesium Mg 12 1s22s22p63s2 [Ne]3s2 4  Atomic Atomic Shorthand Element Expanded notation Symbol number notation Sulfur 37 Ga Cl 1s22s22p63s23p64s23d104p65s2 Silicon 55 Sodium [Kr]4d4 5s1 Lesson 1.3 Quantum Numbers Quantum numbers specify the layout of an electron’s possible position in an atom. An electron in an atom or ion has four quantum numbers to describe its state. Every electron in an atom has a specific and unique set of these four quantum numbers. 1. Principal Quantum Number (n) The principal quantum number, signified by (n), is the main energy level occupied by the electron. Energy levels are fixed distances from the nucleus of a given atom. They are described in whole number increments (e.g., 1, 2, 3, 4, 5, 6...). At location n=1, an electron would be closest to the nucleus, while at n=2 an electron would be farther, and n=3 farther yet. The principal quantum number also corresponds to the row number for an atom on the periodic table. 2. Angular Momentum Quantum Number (ℓ) The angular momentum quantum number, signified as (ℓ), specifies the shape of an orbital with a particular principal quantum number. The value of ℓ depends on the value of the principal quantum number, n. The angular momentum quantum number can have positive values of zero to (n−1). If n=2, ℓ could be either 0 or 1. 3. Magnetic Quantum Number (mℓ) The magnetic quantum number, signified as (mℓ), describes the number of orbitals within a sublevel. Electrons can be situated in one of three planes in three-dimensional space around a given nucleus (x, y, and z). For a given value of the angular momentum quantum number ℓ, there can be (2ℓ+1) values for mℓ. As an example: n=2 ℓ = 0 or 1 for ℓ = 0, mℓ = 0 for ℓ = 1, mℓ = −1, 0, +1 4. Spin Quantum Number (ms) The spin quantum number describes the spin of a given electron. An electron can have one of two associated spins; (+1/2) spin, or (−1/2) spin. An electron cannot have zero spin. We also represent spin with arrows ↑ or ↓. A single orbital can hold a maximum of two electrons and each must have the opposite spin. In summary, the four quantum numbers can be used to describe the outermost or valence electron of an atom. They are: n is the distance of the orbital from the nucleus (1,2,3,4...) ℓ is the shape of the orbital (0, 1, 2, 3, …, n−1) mℓ is the orientation of the orbital in space (−ℓ, …, -1, 0, +1, …, + ℓ) ms is the spin of the electrons (+1/2 or -1/2) 5  To determine the set of quantum numbers of a particular electron in an atom, take note of the following recommendations: Begin each element by writing out the electronic configuration in terms of s, p, d, and f orbitals. From this, you can directly get n and find ℓ based on its equivalent orbital. From ℓ you can figure out the range of mℓ, and then count one up in the sequence for each electron in the outermost subshell. ms is either −1/2 or +1/2. Activity 1.4: Quantum Numbers A. Complete the table by writing the following information: element symbol, electron configuration and the sets of quantum numbers of the outermost electron. Chemical Sets of Quantum Numbers Symbol Name of Magnetic and Electron Principal Azimuthal Magnetic Element Orbital Atomic Configuration (n) (l) Spin (ms) (ml) Number Chlorine Nitrogen Sodium Carbon Beryllium B. Write the set of quantum numbers for ALL the electrons of: 1. Oxygen 2. Magnesium My Element’s Profile Study the first 31 elements in the Periodic Table. Choose the element with the atomic number that corresponds to the day number of your birth. Create a profile of it by identifying and describing the following: 1. Element’s Name and Origin 2. Atomic Number and Mass Number 3. Electron Configuration 4. Quantum Numbers of the Outermost Electron 5. Significant use in your community. Note: You may use sources from the internet on the common uses of the element in your community. Be creative in the presentation of your output and be guided by the given rubrics. 6  MODULE 2 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 2 lessons, namely: Lesson 2.1 – Ionic and Covalent Compounds Lesson 2.2 – Physical Properties of Ionic and Covalent Compounds and their Effects on the Environment After going through this module, you are expected to: 1. identify and describe ionic and covalent compounds based on their chemical formula and chemical names; 2. enumerate and discuss the different physical properties of ionic and covalent compounds; 3. distinguish ionic from covalent compounds based on their physical properties; and 4. cite natural phenomena that use different physical properties of ionic and covalent compounds (ex. snowflakes, voltaic cells). Lesson 2.1 Ionic and Covalent Compounds Elements have the ability to combine with another element to form a chemical compound. These compounds are made up of strong chemical bonds that hold them together so that they appear as a single substance. These chemical compounds have two types depending on the chemical bonding that they have; ionic and covalent compounds. When a metal transfers an electron to a non-metal element, ionic compounds are formed. As a result, these elements become ions. When an atom gains or loses an electron, it forms an ion. A cation is a positively charged ion, whereas an anion is a negatively charged ion. When dealing with ionic compounds, you must first determine the element's ionic charge. The number of electrons that may be lost or gained by an element, in order for it to maintain a stable valence shell that satisfies the octet rule, is determined by its ionic charge. The potential gain and loss of electrons is described in the Periodic Table of Elements, as Figure 2.1.1. The ionic charge of elements in each specific group in the illustrated in Figure 1. Periodic Table of Elements. The sodium (Na) atom, for example, belongs to Group 1A, whereas the chlorine (Cl) atom belongs to Group 7A. The sodium atom contains 11 protons and 11 electrons in its outer shell, with an electron configuration of 1s22s22p63s1. It's simpler to lose this outside electron than it is to acquire seven more. The chlorine atom contains 17 protons and 17 electrons, and its electron configuration is 1s22s22p63s23p5, which implies it has 7 valence electrons and requires one more to be stable. As a result, when the outer electron of the 7 sodium atom transfers to the chlorine atom, both elements become stable, following the octet rule, with 8 valence electron of 8. In this case, an ionic compound called sodium chloride (NaCl) is produced, as illustrated in Figure 2. Covalent compounds are formed when Figure 2.1.2. Ionic compound of salt (NaCl) two non-metals combine. These non-metals have the tendency to "gain" electrons to fill their outer valence shell to follow the octet rule, but neither of them gains or loses an electron. As a result, they share their electron/s with one another so that both/all of them will become stable. A good example of a covalent compound can be seen when a fluorine atom combines with another fluorine atom, making them diatomic molecules by sharing their electrons. Each atom holds 6 outer electrons, and both achieve octet valence electrons by sharing the two electrons in the bond, as Figure 2.1.3. Fluorine atom presented in Figure 3. We can also determine the type of chemical bond that exists between compounds by getting its electronegativity difference. Electronegativity is the ability of the atom to attract electrons in a chemical bond. When the computed difference is between 0-1.9, then we can say that these atoms have covalent bonding and when the computed difference is greater than 1.9, then we can say that they have ionic bonding. As seen in Figure 4, the electronegativity value rises from left to right throughout the period, and from top to Figure 2.1.4. The electronegativity trend in Periodic table of bottom within the group. Photo credits: Elements https://sciencenotes.org/wpcontent/uploads/2019/09/ElectronegativityTrend2.png Activity 2.1: Do you know my bond? Use your Periodic Table to categorize the elements in each formula as metal or non-metal. Determine the type of chemical bond that exists between/among them. Formula Metal Non-Metal Type of Bond Hydrogen H2 O Covalent Oxygen 1. MgCl3 2. C4H10 3. NO2 4. Al2O3 5. C2H6 6. CBr4 7. BaF2 8. Na2S 9. Sr3N2 10. HF 8  The Physical Properties of Ionic Lesson and Covalent Compounds and 2.2 Their Effect on the Environment Chemical Compound A chemical compound is formed when two or more different atoms combine. These are chemical compounds that bond chemically and cannot be separated. The formation of chemical compounds can be classified into two types, namely: ionic compounds and covalent compounds. Ionic Compound The first type of chemical compound is the ionic compound. This is the bonding of compounds that gain and lose electrons, which are called ions. It happens when metal reacts with nonmetals. Let us identify the physical characteristics. Physical Properties 1. High boiling and melting points indicate a strong electrostatic force or attraction between oppositely charged ions. I And it takes more energy to move away from the ion attraction. This will cause an ionic substance to melt or boil. 2. Most ionic compounds are solids due to the organization and strong attractions of the ions, making them hard but brittle. Applied force may cause the atoms to be displaced and cause damage or breakage to the ionic compound. 3. Ionic compounds have high electrical and thermal Figure 2.2.1. Toothpaste conductivity. When ionic compounds are dissolved in water, Photo credits: https://www.pikrepo.com/fnihz/bl liquid ions can conduct electricity. ue-and-white-toothbrush-beside- 4. Ionic compounds are insulators. A solid ionic compound white-and-blue-toothpaste-soft-tube cannot conduct electricity. Examples: table salt, toothpaste, rust, crystal gems Covalent or Molecular Compound This type of bonding is the sharing of electrons with another atom. Like in ionic bonds, the effect of electron sharing as stated in the Octet Rule makes the atom stable. Usually, this bonding happens when a nonmetal reacts with a nonmetal, and sometimes between metals and nonmetals. Covalent Bond In this type of chemical bonding, an atom sharing an electron happens not individually but in a pair. The paired electrons are glued to the two atoms. Physical Properties 1. Low boiling point and melting point. Their bond can be separated easily, and a small amount of energy is needed to melt and boil the covalent compound. 2. With a weaker intermolecular attraction among the ions, they are soft and brittle. Figure 2.2.2 Liquid Ammonia 3. They are usually gases and liquids. Photo credits: https://commons.wikimedia.or 4. They are insoluble in water. g/wiki/File:Liquid_ammonia_bo Examples: CO2, water, and ammonia ttle.jpg Chemical compounds are formed naturally in our environment. Perhaps it is accurate to say that we are surrounded by unknown ionic and covalent compounds, and we have to deal with that on a daily basis. The formation of these chemical compounds in the air, water, or on land will cause a variety of phenomena which may or may not be beneficial to humans. 9 Ionic and covalent compounds are responsible for the several natural phenomena. What impact do these chemical compounds have on our everyday lives? 1. Acid Rain or Acid Precipitation During the water cycle, water accumulates and forms clouds, which eventually turn into rain. However, pollutants like sulfur dioxide (SO2) and nitrogen oxide (NO) emitted by automobiles, factories, and fossil fuel combustion may also accumulate. Acid rain or acid precipitation is produced when these chemical substances combine with water vapor. Weathering of structures, peeling paint on walls, and rusting of metals will occur as a consequence of this. It does, however, have a significant impact on living organisms. It will be harmful to trees' leaves and bark, which will inhibit their growth. It changes the fundamental makeup of soil and has an impact on plant development. It alters the pH level of water in the aquatic environment, which is essential for aquatic survival. 2. Ocean Acidification Carbon dioxide (CO2) levels in the atmosphere rise as a result of deforestation and vehicle emissions. The high CO2 concentration in the atmosphere may also be absorbed by the ocean, making it excessively acidic. Because of the large quantity of CO2, calcium carbonates (CaCO3) will be released. Aquatic water with a low pH may retain enough calcium carbonate (CaCO3) for marine creatures with an exterior protective shell, such as crustaceans. This is also critical for the tiny creatures' home, or coral reef. If carbon carbonates are depleted, there will be a reduction in the quantity of marine creatures available for sale on the market. To prevent these occurrences from escalating, let us offer our simple ways to help the environment. Let us stop dreaming about a safe and clean environment and start working to make it a reality. Activity 2.2: Comparing Physical Properties Identify the specific physical properties of ionic and covalent compounds by completing the table below. Choose the correct answer inside the parentheses. Properties Ionic Compound Covalent Compound Melting point (High or Low) Boiling point (High or Low) Polarity (High or Low) Hardness (Hard or Soft and Brittle) State as room temperature (Solid, Liquid, or Gas) Electrical conductivity (can conduct electricity or cannot conduct electricity) Thermal conductivity (can conduct heat or cannot conduct heat) Critical Thinking Questions: 1. Why do ionic compounds have a high melting point and boiling point compared to covalent compounds? 2. When can an ionic compound conduct electricity? 10 Activity 2.3: Fishy Bone Using the fish bone diagram, identify the different causes of acid rain and ocean acidification. Are they In or Out? Check your house for the following compounds and determine their practical uses. Can be found at home? Compounds Practical uses Yes No Sugar Water Ammonia Baking soda Sodium Fluoride (ingredient in toothpaste) MODULE 3 This module was designed and written to help you master the nature of chemical bonding. 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 3.1 – Charged Particles and Formation of Ions Using the Lewis Dot Structure Lesson 3.2 – Chemical Bonding and Writing Chemical Formula After going through this module, you are expected to: 1. determine how ions are formed; 2. differentiate cations from anions based on their tendency to lose or gain electrons; 3. show the formation of ions using Lewis Electron Dot Symbols (LEDS); 4. differentiate ionic from covalent bonds; and 5. write the chemical formulas of ionic compounds based on the charges of ions. 11  Charged Particles and Lesson Formation of Ions Using 3.1 the Lewis Dot Structure An atom is chemically neutral because the number of protons in its nucleus equals the number of its electrons. It also takes part in chemical reactions that result in the formation of molecules. An ion, on the other hand, is produced as a result of a chemical process in which the quantity of protons and electrons is unequally distributed. An ion is an atom that carries an electrical charge generated either by the removal Fig 3.1.1. Comparison of Oxygen Atom from Oxygen Ion or the acceptance of electrons from another neutral atom. Formation of Ions Metal atoms are referred to as "electron donors" because they have a higher probability of losing or donating electrons. On the other hand, non-metal atoms have a higher probability to accept or acquire electrons than metal atoms, so they are referred to as "electron acceptors." They gain electrons until the valence Fig 3.1.2. Formation of Ions shell acquires eight electrons, making it stable. When an atom receives or donates electrons, it transforms into an ion due to the fact that the number of protons does not match the number of electrons. For example, if an oxygen atom (a highly reactive non-metal atom) receives two electrons from another atom, the total number of electrons will be higher than the number of protons. Meanwhile, if a metal atom donates one electron, the number of electrons will be less than the number of protons. Take note: The valence electrons are the only particles that can be gained or lost since they are freely moving and located at the outermost layer of an electron shell. Protons, on the other hand, cannot be gained or lost because if the number of protons changes, the atom or element will change as well. According to the Octet Rule, to attain stability, the outermost electron shell must have 8 electrons. There are two types of ions: cations and anions. If an ion has fewer electrons than protons, then it is called a cation. These are ions that carry a positive charge. Meanwhile, if an ion carries a negatively charged particle, it is called an anion because it has more electrons than protons. Aside from the types of ions, there is another category of ions, namely monatomic ions and polyatomic ions. This category describes the presentation of certain Table 3.1.1. Some Monatomic and Polyatomic Ions 12 ions. Monatomic ions are ions that are composed of a single atom with a single possible charge. Polyatomic ions, on the other hand, are composed of a group of atoms that contain a single charge. Unconsciously, we are surrounded by atoms and ions. Our bodies are made up of atoms, and we require an ion, particularly electrolytes, to control osmotic pressure in our cells and keep muscle and nerve cells functioning. It allows electric impulses to flow freely throughout the body. A magnificent Aurora is a light display that occurs when electrically charged particles from the sun collide with particles from gases in the Earth's atmosphere, such as oxygen and nitrogen. Valence electrons are electrons found at the outermost energy level (orbital) of the atom. They play an essential role in forming bonds to produce compounds. It is important that you know how to identify the number of valence electrons so that Fig.3.1.3. Valence Electron you can illustrate how bonds are formed. of Lithium ion Period refers to the horizontal arrangement of the elements in the Periodic Table. It denotes how many valence shells or orbitals an element has. Family, on the other hand, refers to how the elements are arranged vertically. It represents the number of valence electrons in Figure 3.1.4. Lewis Dot Structure of Representative Family in the Periodic Table an element. Take note: Metals have low electronegativity and Non-metals have high electronegativity. Electronegativity is a measure of the tendency of an atom to attract a pair of electrons. The greater the number of electrons, the higher the tendency to attract electrons. Ionization energy is the energy required to gain or lose one or more electron/s from a neutral atom. The greater the ionization energy, the more difficult it is to lose valence electrons. Metal elements have low electronegativity and ionization energy. Therefore, metals tend to transfer or lose electrons. Non-metal elements have a greater tendency to attract (gain) electrons towards themselves because they have high electronegativity and ionization energy. Lewis Dot Structures (LEDS), also known as electron dot Figure 3.1.5. Lewis Dot Structure of Lithium atom structures, are graphical representations of the valence electrons of atoms within a molecule. This structure may be represented by the Lewis Dot Symbol, which consists of the chemical symbol of an element and the number of valence electrons are drawn around it as dots. The location of the dots begins at the top of the element symbol and moves clockwise to the right (3 o'clock). Then it continues down to the 6 o'clock position, then to the 9 o'clock position, and finally back to the top position until all dots match the amount of valence electrons of the atom. These Lewis symbols and Lewis structures aid in visualizing the valence electrons of atoms and molecules, whether they exist as lone pairs (valence electrons that are not shared with another atom) or within bonds. Activity 3.1. Lewis Dot Symbol Model Objectives: After performing this activity, you should be able to make a model of the Lewis- Dot Symbol of an element. Materials: Periodic Table, colored paper, bond paper, glue, markers, scissors or puncher Procedure: 1. Gather all the materials. 2. Locate and identify the element in period 2, family 4A. 3. Use the Periodic Table to answer the table below. 4. On a bond paper, create the diagram for this element using the data in the table. You may refer to the sample illustration in creating your model. 5. Cut/punch the colored paper into small dots. 13  6. Paste the dots in their proper position to complete your model. Element Element Symbol Family Valence Electron Lewis Dot Structure Example: S 6 6 Sulfur Sample Illustration: Element Illustration: Guide Questions: 1. What is Lewis Dot Structure? 2. What is the importance of Lewis dot structure? 3. Which part of the periodic table represents the number of valence electron? 4. What is valence electron? 5. Why are valence electron/s important? Lesson Chemical Bonding and 3.2 Writing Chemical Formula We know that all living things are made up of atoms. In most cases, atoms are not just floating around, but they are usually interacting with other atoms or groups of atoms. Atoms may be linked by strong bonds and formed into molecules, or they may form temporary, weak bonds with other atoms that they bump into. Both the strong bonds that hold molecules together and the weaker bonds that create temporary connections are important to the interaction of our bodies, and to the existence of life itself. Chemical bonds are the attraction forces that link atoms together. Bonds are formed when valence electrons interact. The type of interaction between the atoms depends on their relative electronegativity. Atoms with equal or similar electronegativity form covalent bonds in which the valence electron is shared between the two atoms. The electrons exist in between the atoms and are both attracted to the nuclei. This type of bond forms often between two non-metal elements. If the electronegativity difference between covalently bonded atoms is more than 0.4 but less than 1.9, the pair of atoms usually forms a polar covalent bond. The electrons are shared between the atoms, but the electrons are not equally attracted to both elements. The result of unequal sharing of electrons between the atoms of different elements developed a part in molecules that is slightly positive (δ+) and slightly negative (δ-) charged. Example: Water molecules (H2O) are composed of hydrogen and oxygen atoms. It is polar because the electrons are unequally shared between the two atoms. Figure 3.2.1. Polar covalent bonding of water molecule Table 3.2.1. Electronegative difference of water molecule Meanwhile, if the electronegativity difference is equal or less than 0.4, the pair of atoms usually forms a nonpolar covalent bond. It is formed between two atoms of the same element, or between atoms of different elements that share electrons more or less equally. 14  Example: Fluorine (F2) is a diatomic atom (contains 2 atoms). It is nonpolar because the electrons are equally shared between the two fluorine atoms. Figure 3.2.2. Non polar covalent Bonding of fluorine atom Table 3.2.2. Electronegative difference of fluorine atom Ionic bonding takes place if atoms have an electronegativity difference of greater than 1.9. This means that a metal atom can fully transfer the valence electron to a nonmetal atom. Once the electrons have been transferred to the non-metal atom, both the metal and the non- metal are considered to be ions. The two oppositely charged ions attract each other to form an ionic compound. Example: When sodium atom (metal) is attracted to chlorine atom (non-metal), it forms sodium chloride (ionic compound). It forms an ionic bond because the electrons are fully transferred from one atom to another. Figure 3.2.3. Ionic bonding between sodium and chlorine atom. Table 3.2.3. Electronegative difference between sodium and chlorine atom. Why do chemical bonds form? Atoms are trying to reach the most stable state that they can. Many atoms become stable when their valence shell is occupied by electrons or when they satisfy the Octet Rule. If atoms have an incomplete valence electron, they will strive to reach it by gaining, losing, or sharing electrons through bonds. The Octet Rule states that elements need to gain, lose, or share electrons to reach the electron configuration of the nearest noble gas. Why Noble Gas? Noble gas elements are the only group of elements in the periodic table that have complete valence electrons in their outermost shells, which makes them very stable. These are the elements found in Family 8A. Other elements are also striving for stability, which controls their reactivity and bonding to other atoms. Example: Sodium chloride, an ionic compound, is formed when a sodium atom (metal) completely transfers its valence electron to a chlorine atom (non-metal). The sodium atom originally had 11 electrons, whereas the chlorine atom had 17 electrons. Sodium loses one electron after ionic bonding. As a result, it has 10 electrons remaining. On the other hand, the chlorine atom acquires one, which results in a total of 18 electrons. When looking at the noble gas family in the Periodic Table, the sodium atom ended up having the same number of electrons as the neon ion (10), whereas the chlorine ion is the Figure 3.2.4. Lewis Dot of Ionic bonding between sodium and chlorine atom. same as the argon ion (18). After chemical bonding, elements become isoelectronic (having a similar electron configuration) to the nearest noble gas in the periodic table. Individual atoms become ions; metal atoms become cations, while non-metal atoms become anions. 15 Chemical Formula You've learned that some ionic compounds are composed entirely of a metal ion (cation) and a non-metal ion (anion). Certain ionic compounds, however, are composed of more complicated atoms, referred to as polyatomic ions. Many of the products we use on a daily basis contain these polyatomic ions. Ionic compounds can be represented using a chemical formula. A chemical formula is an expression of the atomic composition of Figure 3.2.5. The Anatomy of a Chemical Formula a compound, which consists of the number and type of atoms present in a substance or compound and the symbol/s of the element/s involved. Charge: Metals found in Group/ +1 Family 1A, 2A, and 3B have low +2 +3 ±4 -3 -2 -1 electronegativity and ionization energy, so they tend to transfer or lose electrons. On the other hand, non-metals, found in 5A, 6A, and 7A, have high electronegativity and ionization energy, which results in a greater tendency to attract electrons towards themselves. Thus, non- metals tend to gain electrons. Transition elements usually have more than one possible charge, like iron (Fe2+ and Fe3+). However, some have only one possible charge, like Figure 3.2.6. The Periodic Table of Elements silver (Ag1+), zinc (Zn2+) and cadmium (Retrieved from https://sciencenotes.org/periodic-table-of-elements-hd/) (Cd2+). Writing Chemical Formula Chemical formulas are written with the elements' chemical symbols followed by numeric subscripts indicating the relative ratios of the constituent atoms. Remember that in order to satisfy the Octet Rule, the charge of the constituent ions will be determined based on the number of valence electrons. When cations and anions are combined, an electrically neutral compound is formed. For instance, you are given calcium and chlorine. Calcium has a charge of +2 since it belongs to Group 2A. Chlorine, on the other hand, has a charge of -1 since it belongs to Group 7A. In other words, this compound contains Ca2+ (cation) and Cl– (anion). After losing and acquiring electrons, those ions will become stable, completely filling out their valence shells. Their ionic formula will be CaCl2 (calcium chloride), which indicates their neutral composition. Two chloride ions were required in the resulting compound due to the +2 charge of calcium. Two -1 chloride ions were required to balance out the +2 charge from calcium in order to make the neutral molecule CaCl2. How to write Chemical Formula? Rules Example No. 1 Example No. 2 1. Write the formulas for the cation Sodium ion + nitrite Aluminum + sulfite and anion (including the charges). Na1+ NO21- Al3+ SO3-2 2. Use the Criss-Cross method to balance the subscripts. 3. Balance the charges, if necessary, using the subscripts. Use Na1(NO2)1 Al2(SO3)3 parentheses if you need more than one of a polyatomic ion. 4. Simplify the formula. NaNO2 Al2(SO3)3 16 Activity 3.2. Bonding by Transfer of Electrons Objectives: After performing this activity, you should be able to illustrate how ionic bond is formed using Lewis-Dot Structure. Materials: Periodic Table Procedure and Example: Complete the table below and write the Lewis Symbol for the selected elements. Take note of the electronegativity value of both elements. Subtract the electronegativity value of the metallic element from the non- metallic element. Guide Questions: 1. After ionic bonding, what kind of an element forms a cation? 2. What about the element that turns into an anion? 3. How does an atom attain stability after ionic bonding? 4. Do all the attractions between metals and non-metals form ionic bonds? Why? 5. Why do ions form after ionic bonding? Activity 3.3. What’s The Formula? Objective: After performing this activity, you will be able to write the chemical formula of the compound formed between anions and cations. Directions: Write in the box the chemical formula of the compound formed by the cations and the anions. Guide Questions: 1. What do you call the ions on top of the column? 2. What do you call the ions on the left of the row? 3. What method did you use in getting the product between cation and anion? 4. What is the importance of writing the chemical formula of the compound? Supply the following Science journal- writing prompts with your own ideas. Write the completed journal in your notebook following the format in the box below. 17  MODULE 4 This module was designed and written to help you master the nature of chemical bonding. 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 4.1 – The Chemistry of Carbon Lesson 4.2 – Properties of Organic Compounds After going through this module, you are expected to: 1. discuss why carbon is a unique atom (valence electron, bond length, strength, multiple bond formation, etc.); 2. differentiate organic from inorganic compounds (from their chemical formula, uses, properties); 3. determine the different uses of organic compounds and cite examples of each organic compound; and 4. find the properties of common organic compounds through experimentation. Lesson 4.1 The Chemistry of Carbon Carbon is everywhere! It is considered as the building blocks of life on earth. It may appear in the form of diamonds or graphite, or as various kinds of carbon or organic compounds. It serves as the key ingredient for most life on earth; it gives us fuel for energy, it creates life, and it has an essential role in regulating our climate. Carbon atom has a unique ability to bond with Table 4.1.1. Different types of carbon chains variety of elements. It contains four valence electrons that form covalent chemical bonds, wherein electrons are shared to attain stability. This allows the formation of multiple stable bonds with other small atoms such as hydrogen, oxygen, nitrogen, and even carbon itself. With this, carbon atoms are capable of forming long and infinite straight or branched carbon- carbon chains, as well as closed or cyclic chain. Table 4.1.2. Different types of bonds These chains are composed of either single, double, or triple carbon-carbon bonds. As a result, carbon atoms form various enormous and complex molecules that can be found everywhere around us. Carbon compounds have the ability to form two different molecules with the same number of carbon atoms but with different geometrical arrangements. They are called isomers. Isomers are molecules or polyatomic ions with the same molecular formula but different shapes and properties. Carbon compounds can be in the form of gas, liquid or solid. In the non-living environment, we can find carbon compounds in the atmosphere, in carbonate rocks, and in fossil fuels such as coal, oil, and gasoline. In the living environment, carbon atoms form the structural molecular backbone of the significant molecules of life, such as proteins, carbohydrates, lipids, and nucleic acids. It is a versatile element; it can exist in very small 18 two-atom molecules such as carbon monoxide (CO) up to more complex molecules that contain thousands of atoms, such as proteins and DNA. Activity 4.1. The Anatomy of Carbon Atom Directions: Using the diagram below, answer the following questions. 1. What is the name of the element? 2. Where is element C located? Atomic Number ___ Period Number ___ Family Number ___ 3. How many valence electrons (outermost) are present in the diagram? 4. How many atom/s are allowed to bond in 1 element C atom? 5. Is it a metal or a non-metal? 6. If the atom of this element is bonded with a non-metal atom, will the valence electrons be shared or transferred? Lesson Properties of Organic Compound 4.2 Organic and Inorganic Compounds Organic and inorganic compounds are the two general categories of compounds in Chemistry. Generally, organic compounds are compounds that contain carbon atoms, mostly with carbon-hydrogen or C-H bonds. However, there are a few important exceptions to this rule, such as carbon dioxide (CO2) and other complex negative ions that contain carbon but aren’t organic. Compounds that are made up of two or more elements other than carbon, as well as certain carbon-containing compounds that lack carbon-carbon bonds, such as cyanides, carbides, and carbonates, are called inorganic compounds. Table 4.2.1. Difference Between Organic and Inorganic Compounds Points of Organic Compounds Inorganic Compounds Comparison Organic compounds are biological Inorganic compounds are simple Nature of and more complex in nature. They and mineral in nature. These existence are mainly found in most of the compounds are usually found in living things around us. non-living things. They exist as solids, liquids, or Most of them are solids. Very few State gases. exist as liquids and gases. Mostly, they are compounds with They are made from single carbon. They also contain hydrogen, elements. They do not contain oxygen, and their derivatives. carbon-carbon or carbon-hydrogen Examples (e.g., Nucleic acids, sucrose, bonds. (e.g., salt, metals, silver, enzymes, benzene, methane, fats sulfur, pure diamonds) and ethanol) Bond They form covalent bonds wherein They form ionic bonds, wherein Formation electrons are shared. electrons are transferred. They are soluble in organic solvents They are highly soluble in water, Solubility but are less soluble in water. but insoluble in organic solutions. However, some are soluble in water. Rate of Organic compounds have a slow rate Inorganic compounds have a high Reaction of reaction. rate of reaction. They are volatile compounds. They They are non-volatile compounds. Volatility evaporate easily at normal temperatures. 19  Points of Organic Compounds Inorganic Compounds Comparison They are highly inflammable. They They are not inflammable. Flammability can be easily set on fire. In most aqueous solutions, organic In most aqueous solutions, compounds are typically poor inorganic compounds are typically Conductivity conductors of electricity and heat. good conductors of electricity and heat. The intermolecular forces of organic Inorganic compounds, on the other compounds are weak; thus, their hand, tend to have strong Viscosity viscosity tends to be low. intermolecular forces, such as hydrogen bonding, thus they have a higher viscosity. They are used in industries like They are commonly used as foods, pharmaceuticals, fuels, etc. catalysts, pigments, coatings, They are also present in surfactants, and more. They are Uses photosynthesis and cellular used in the ceramic industry. In respiration. They cannot produce the electrical field, they are applied salt. to electric circuits. They can produce salt. Activity 4.2. My Chemical Formula Based on their chemical formulas, identify the following materials as organic or inorganic compound. 1. Benzene, C6H6 4. Sodium Chloride, NaCl 2. Hydrochloric acid, HCl 5. Carbon dioxide, CO2 3. Sucrose, C12H22O11 Activity 4.3. Organic or Inorganic? Below are some materials which are commonly used in everyday lives. Based on their properties, tell whether they are organic or inorganic. Organic or Materials Properties Inorganic? 1. Ethyl Ethyl alcohol, or ethanol, is a clear liquid that is highly Alcohol flammable. It easily mixes with water and with other organic compounds. 2. Salt Salt is a very common compound found in our homes. It appears as a solid, crystalline powder. It is soluble in water and not inflammable. 3. Water Water has the ability to dissolve many substances. It contains a polar covalent bond. 4. Gasoline Gasoline is a flammable liquid used primarily as a fuel. It is volatile and evaporates easily at normal temperatures. 5. LPG LPG, or Liquefied Petroleum Gas, is commonly used as a home heating and cooking fuel. It is colorless in both liquid and vapor states. It is highly flammable. Activity 4.4. Organic Compounds: Are they Useful? (This activity is lifted from the Department of Education Science 9 Learner’s Module (First Edition, with minor modifications to suit the distance learning modality.) Objective: In this activity, you will be able to recognize the uses of common organic compounds. Procedure: 1. Label the following products. Choose your answer from the word bank below. WORD BANK gasoline kerosene LPG acetone acetic acid ethanol 20  ________________________ ________________________ ________________________ ________________________ ________________________ ________________________ 2. Complete the table below. Using a check mark, indicate the uses of the given compounds. You may have more than one check mark per sample depending on its use/uses. Organic Compounds Uses Acetic Gasoline Ethanol Acetone LPG Kerosene Acid Beverage/ Drinks Food Antiseptic/ Disinfectant Fuel Cleaner Guide Questions: 1. What properties do the following organic compounds have in common? 2. Why do you think organic compounds like these are very important? Making Connections: Science at Home Organic compounds are literally everywhere around you. Look for three common organic compounds inside your home or within the premises of your neighborhood. Describe each of them by completing the table below. Name of the Material Properties Use/s at Home Example: strong, durable and moisture- Used as flooring Vinyl resistant (Vinyl tiles) 1. 2. 3. MODULE 5 This module was designed and written to help you master the concepts about Hydrocarbons. 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 three lessons, namely: Lesson 5.1- Classes of Hydrocarbons Lesson 5.2- Naming Hydrocarbons Lesson 5.3- Functional Groups 21 After going through this module, you are expected to recognize the general classes and uses of organic compounds (S9MT-IIh-18) Specifically, this learning resource material aims to: 1. differentiate molecular, empirical, and structural (expanded and condensed) formulas; 2. differentiate alkanes, alkenes and alkynes based on the presence of bonds and their physical properties; 3. name different structures of hydrocarbons (alkanes, alkenes, alkynes, alcohol, aldehydes, esters, carboxylic acid, ether, amines, amides, etc.; and 4. identify different functional groups and their uses in organic compounds. Lesson 5.1 Classes of Hydrocarbons Hydrocarbons are organic compounds that are made up of only two elements, hydrogen and carbon. These hydrocarbons have different types depending on their structure; the alkanes, the alkenes and the alkynes. Alkanes Hydrocarbons that contain single bonds are called alkanes. Alkanes are also known as paraffins. Each carbon atom is bonded to four other atoms, making it a saturated compound, and there are no double or triple bonds in the molecules. Alkanes’ general formula is expressed as CnH2n+2, where n is the number of carbon atoms. The suffix –ane is used in naming an alkane. Some examples of alkanes include methane (the simplest hydrocarbon), ethane and propane are shown on the right. In the IUPAC system, various prefixes may be used to indicate the presence of carbon Table 5.1.1. Prefixes corresponding to the number of atoms in an unbranched chain. carbon atoms Alkenes and alkynes Alkenes and alkynes are types of hydrocarbons that are unsaturated because they contain carbon atoms attached to less than four atoms. Alkenes, also known as olefins, contain a double bond between two carbon atoms and contain fewer hydrogen atoms when compared to alkanes. It has a general formula of CnH2n. The suffix –ene is used in naming an alkene. Ethene is an example of alkene. Another type of hydrocarbon is called alkynes. They are also called acetylenes. They contain a triple bond between two carbons and have the general formula of CnH2n-2. The suffix –yne is used in naming an alkyne. Some examples of alkynes are shown on the right. These hydrocarbons may be shown in a variety of ways. A molecular formula, an extended structural formula, a condensed structural formula, a carbon skeletal form, or a line form may all be used. In most cases, instead of the alternative representations, the expanded and condensed formulas are used in the study of Table 5.1.2. Comparison of the Different Representations of Butane hydrocarbons. 22 Activity 5.1. Let’s Compare! Complete the table to demonstrate the differences between alkane, alkene, and alkyne. Description Alkane Alkene Alkyne Number of Bonds General Formula Suffix used in Naming Examples (Give 2) Activity 5.2: Which Hydrocarbon Am I? Now that you're familiar with the various kinds of hydrocarbons and their characteristics, examine the provided structure and determine whether it's an alkane, alkene, or alkyne. Hydrocarbon Type Hydrocarbon Type 6. 1. 2. 7. 3. 8. 4. 9. 5. 10. Lesson 5.2 Naming Hydrocarbons The naming of organic compounds is known as organic nomenclature. There are a few rules for naming organic compounds that have been devised by the International Union of Pure and Applied Chemistry (IUPAC). Knowing these rules and given a structural formula, one should be able to write a unique name for every distinct compound. Likewise, given the IUPAC name, one should be able to write a structural formula. Given below are the established rules by the IUPAC for naming hydrocarbons: 1. To determine the suffix of the name, identify the functional group in the compound. a. Straight-chain hydrocarbons that contain only single-bond carbons are called alkanes. When naming these molecules, the final syllable in their names is "ane." b. Straight-chain hydrocarbons that contain at least one double bond carbon are called alkenes. In naming these molecules, the final syllable in their names is "ene." c. Straight-chain hydrocarbon molecules that contain at least one triple bond carbon are called alkynes. To name these molecules, the final syllable in their names is "yne." 23 2. Find the longest continuous carbon chain that contains the functional group and count the number of carbon atoms in this chain. Take note that it will not always be a straight chain. This number will determine the prefix of the compound's name. Prefixes are attached at the beginning of the word to indicate the number of carbons in the hydrocarbon chain. The IUPAC prefixes for the first 20 carbon chain lengths are introduced in Lesson 1. 3. Alkenes and alkynes with hydrocarbon chain lengths of four or greater require a numbering scheme for the chain to designate the location of the multiple bonds. The rules regarding this numbering scheme are: a. The numbering scheme starts at the molecule's end, closest to the position of the multiple bonds. b. The number in the name is determined by assigning a numerical value relative to the carbon in the chain where the multiple bonds first appear. c. Each multiple bond in the chain is represented by a number. Commas are used to separate these numbers. d. A dash separates the number from the name of the hydrocarbon. e. When there is more than one multiple bonds, prefixes are added to the "ene" or "yne" final syllable. 4. For ring hydrocarbons, the prefix "cyclo" is attached to the hydrocarbon name. 5. Combine the elements of the name into a single word in the following order: branched groups/halogen atoms in alphabetical order (ignoring prefixes); prefix of the main chain; and name that ends according to the functional group and its position on the longest carbon chain. Examples: Steps for Naming Example No. 1 Example No. 2 Hydrocarbons Examine the compound. The compound is a Since the molecule contains hydrocarbon with single Determine the functional a double carbon-carbon bonds between the carbon group. bond, it is an alkene and the atoms. It is an alkane and suffix is -ene. will have the suffix -ane. Find the longest continuous There are four carbons in the carbon chain and number the longest (straight) chain, so carbon atoms in that chain. There are six carbons in the the prefix is but-. longest chain, so the prefix is hex-. There is one methyl group attached to the main chain. No branched group is Look for any branched After the functional group, identified in the given groups. Name them and the branched group should hydrocarbon. The double identify their position in the have the lowest number bond occurs at the second carbon chain. possible. So, the methyl is carbon. attached to the third carbon atom (3-methyl). Combine the components of the hydrocarbon’s name into a single word following this The hydrocarbon’s name is The hydrocarbon is 2- order: prefix; suffix according 3-methylhexane. butene or but-2-ene. to the functional group and its position along the longest carbon chain. 24  Steps for Naming Hydrocarbons Example No. 3 Examine the compound. Determine the functional group. There is a triple carbon-carbon bond. This is an alkyne and will have the suffix -yne. Find the longest continuous carbon There are three carbon atoms in the longest chain, chain and number the carbon therefore the prefix will be prop-. The triple bond atoms in that chain. appears in the first carbon atom, so the numbering will begin from the left. This will make the suffix -1- yne. Look for any branched groups. There are no branched groups. Name them and identify their position in the carbon chain. Combine the components of the The hydrocarbon’s name is 1-propyne or prop-1- hydrocarbon’s name into a single yne. word following this order: prefix; suffix according to the functional group and its position along the longest carbon chain. Steps for Naming Hydrocarbons Example No. 4 Give the structural formula for 1-hexene. Then, find its molecular formula. Solution: Structural Formula: The prefix hex- tells us there are six carbon atoms in the chain. The suffix -1-ene tells that there is a Molecular Formula: C6H12 double bond between the first and (Note: Just count the number of each atom from the second carbon atoms. structural formula.) Activity 5.3. Classify Me! Analyze the given molecular formula. Determine if the compound is an alkane, alkene, or alkyne. 1. CH4 4. CH3CH2CCCH2CH3 2. CH3CH=CH2 5. CH2=CHCH2CH3 3. CHCH Activity 5.4: Identify Me! Fill in the table below with the compound's name or structural formula. IUPAC hept-3-ene octane ethyne Name Structural Formula Lesson 5.3 Functional Group The functional groups determine an organic compound's structural formula, properties, and even chemical reactivity. Regardless of the molecule in which they are found, 25 functional groups have their own unique characteristics. Alcohols, ethers, aldehydes, ketones, carboxylic acids, esters, amines, and amides are common examples of functional groups. 1. Alcohol Alcohols are transparent and colorless organic compounds that are commonly used as disinfectants, solvents, fuels and reagents. The present functional group in alcohols is the -OH group. The most common types of alcohol are methyl alcohol, ethyl alcohol, propyl alcohol, and isopropyl alcohol. Ethanol is one of the common alcohols that is produced through fermentation and can be found in alcoholic beverages. In the current situation, where we are battling with a deadly virus known as Severe Acute Respiratory Syndrome Coronavirus-2 (SARS-CoV-2), health experts recommended using isopropyl alcohol as an alternative if water and soap are unavailable. It acts as a disinfectant and can also kill viruses, and it is commonly used in hospitals. To name the alcohols, locate the longest chain of carbon that holds the –OH group. If a structure contains -OH, you just need to use the -ol suffix with the name of the parent alkane. For example, the name of the structure on the right is butanol, from the alkane butane, since the parent chain contains four carbons. 2. Ethers Ether contains an oxygen atom that is attached to two alkyl groups (R-O-R). This organic substance is extremely flammable. In addition, they are commonly used as solvents. Common examples of ether are dimethyl ether and diethyl ether. Dimethyl ether is used as an aerosol spray propellant while diethyl ether is used as an anesthetic for surgical procedures. When naming ether, it is necessary to specify the alkyls linked to oxygen. The name should include the alkyls that are ordered alphabetically before the term ether. For instance, the image on the right shows two alkyls, referred to as methyl, attached to oxygen. This structure is referred to as dimethyl ether. The prefix di- denotes two identical alkyl groups linked to oxygen. 3. Aldehydes The carbonyl group is the functional group present in aldehyde. The carbonyl group is formed by a double bond between a carbon atom and an oxygen atom. Aldehydes have the generic formula R-CHO. Aldehyde is derived from the phrase "alcohol dehydrogenated," and its most basic form is methanol or formaldehyde. In funeral homes and medical labs, formaldehyde is used to preserve the dead corpses of organisms. To rename aldehydes, use –al suffix with the name of the parent chain. For example, the structure on the right is composed of two carbon chains; therefore, the parent chain is ethane. Change the suffix to -al. The name of the structure is ethanal. 4. Ketones A ketone is a functional group that is characterized by the presence of a carbonyl group (O=C) connected to two other carbon atoms. Ketones and aldehydes are named similarly except for the suffix used and the number that denotes the carbonyl group's position. To name ketones, use the sffix -anone with the name of the parent chain. For instance, the longest carbon chain in the structure on the right is three. Therefore, the name of the parent chain is propane. Change the suffix from -ane to –anone, then indicate the location of the carbonyl group. The name of the structure is 2-propanone. 5. Carboxylic Acid Carboxylic acid is characterized by the presence of the carboxyl functional group (COOH). The most often encountered carboxylic acid is dicarboxylic acid, which contains two carboxyl groups. Oxalic acid, which is utilized as a cleansing agent, is a good example of this. Vinegar, which also contains a carboxyl group, is a common ingredient in a wide variety of recipes and is also used as a preservative. Acetic acid is the carboxylic acid found in vinegar. 26  To name a carboxylic acid, change the suffix of the parent chain to -oic acid. For instance, the longest continuous carbon chain on the right is four. Therefore, the name of the parent chain is butane. By adding the suffix –oic acid, the structure becomes butanoic acid. 6. Esters Esters are formed from the reaction between a carboxylic acid and alcohol. It has a general formula of RCOOR, wherein two alkyls are attached to O=C-O. Generally, esters are mildly polar and have pleasant odors. Some fruits contain esters, such as pineapples, bananas, raspberries, and oranges. In naming esters, use the alkyl as the first name while the second should be the stem name of the acid. The second name should end with the suffix –oate. Looking at the picture on the right, the first alkyl is methyl; the second alkyl attached to the functional group is butyl. The name of the structure is methyl butanoate. 7. Amines Amines are organic compounds containing the functional group of basic nitrogen atoms. They are derivatives of ammonia, NH3. The hydrogen of ammonia can be replaced by R or alkyl groups. Most of the amines are neurotransmitters in the brain, spinal cord, and other body parts. One of the common amines that act as neurotransmitters in the brain and spinal cord is histamine. Histamine is involved in inflammatory responses and its main role is to act as a mediator for itching. When naming the amines, add the suffix –amine to the alkane chain connected to it. For instance, the alkyl attached to the structure on the right is methyl, with a single carbon atom. Simply add the suffix - amine to the end of the name. So, the structure is known as methylamine. 8. Amides The amides are formed from the reaction of carboxylic acid and amine. Most of the amides are drugs, including penicillin. Penicillin is the most common antibiotic that treats bacterial infections in wounds. It was discovered by a famous scientist named Alexander Fleming. In naming the amides, add the suffix –amide to the stem of the acid's parent’s name, and omit the word acid. For example, the structure on the right is an amide derivative of acetic acid. To name it, just replace the suffix -ic with -amide and omit the term acid. Therefore, its name is acetamide. Activity 5.5 Know My Importance Activity 5.6 Know My Identity Complete the table below. Identify the following functional groups. Type of General Example Uses Compound Formula 1. _______________ 2. __________________ 3. _______________ 1. Alcohol 2. Ether 3. Aldehyde 4. _______________ 5. _______________ 6. _______________ 4. Ketone 5. Carboxylic Acid 6. Ester 7. _______________ 8. ________________ 9. ________________ 7. Amine 8. Amide 10. ______________ 27 My Hydrocarbon! Construct the structure of the hydrocarbon of your choice using materials available in your house. These materials will represent the hydrocarbon's hydrogen atoms, carbon atoms, and bonds. Examples: Marshmallows as hydrogen atoms Raisins as carbon atoms Toothpicks as bonds Write one paragraph that will describe the structure of your hydrocarbon. Be creative in the presentation of your output and be guided by the given rubrics. MODULE 6 This module was designed and written to help you master how pieces of matter are quantified by finding the mass or by counting or number of atoms, ions, or molecules. 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 four lessons, namely: Lesson 6.1 – Formula Mass vs. Molecular Mass Lesson 6.2 – What is in a Mole? Lesson 6.3 – Number of Particles Lesson 6.4 – Conversion of Units After going through this module, you are expected to: 1. differentiate molecular unit from formula unit and compute for the molecular mass; 2. define the term "mole" and calculate the number of moles in a given compound; 3. determine the number of particles based on the mass or number of moles; and 4. convert the number of moles, mass, and number of particles from one unknown to another. Lesson Formula Mass vs. 6.1 Molecular Mass You can easily calculate the number of particles in a given mass by first defining its molecular or formula mass. If a n element's atomic mass is the average mass of its atom, the 28 molecule's molecular mass is the average mass of its molecules. Numerically, the molecular mass of a substance (or formula mass in the case of an ionic compound) equals the total of the atomic masses of the elements/atoms included within the molecule or formula unit. Molecular Mass Formula Mass The sum of the atomic masses of all the The sum of the atomic masses of all the atoms in a molecule. atoms/ions present in one formula unit of any compound, whether molecular or ionic. SI Unit: atomic mass unit (amu) SI Unit: grams per mole (g/mol) Steps in Calculating Molecular Mass or Formula Mass Sample Problem 1.1: Calculate the molecular mass of carbon dioxide. 1. Write the chemical formula of the covalent compound. Carbon dioxide → CO2 2. Determine and write the atomic masses of each element present in the compound. C = 12 g O = 16 g 3. Count the number of atoms for each element. Then, multiply it to its respective atomic mass. Lastly, add the masses of each element. No. of Atomic atom mass CO2 → 2 x 16g = 32g 1 x 12g = 12g 44 amu Sample Problem 1.2: Calculate the formula mass of sodium chloride. 1. Write the chemical formula of the ionic compound. sodium chloride → NaCl 2. Determine and write the atomic masses of each element present in the compound. Na = 23g Cl = 35g 3. Count the number of atoms for each element. Then, multiply it to its respective atomic mass. Lastly, get the sum of the molar masses of each element. No. of Atomic atom mass NaCl → 1 x 35g = 35g/mol Interpretation: 1 x 23g = 23g/mol A mole of NaCl is equal 58 g/mol to 58g of NaCl. When a large number of molecules or formula units are involved in a visible chemical reaction, it is convenient to employ a unit of measurement called the mole (mol). One mole of a substance has the same mass as its relative molecular mass given in grams. Activity 6.1. Decode the Message Calculate the molecular or formula mass of each compound and write the letter in their respective places to decode the message. Atomic Masses of the Elements: C – 12g H – 1g O – 16g F – 19g Ba – 137g Na – 23g Ca – 40g 1. Carbon dioxide, CO2 O. 40g/mol 2. Water, H2O O. 18 amu 3. Calcium fluoride, CaF O. 59g/mol 4. Methane, CH4 D. 16 amu 5. Barium oxide, BaO G. 44amu 6. Sodium hydroxide, NaOH J. 153g/mol 7. Hydrogen peroxide, H2O2 B. 34 amu ______ ______ ______ ______ ______ ______ ______ ! 1 2 3 4 5 6 7 29  Lesson 6.2 What is in a Mole? Chemists determine the number of atoms and molecules by relating them to their masses. They developed a unit of measurement called the mole. The number of atoms represented by an element's atomic mass is expressed in grams (g). In the year 1896, renowned German chemist Wilhelm Ostwald introduced the term "mole," which is derived from the Latin word "mole," that means "great amount" or "pile." One mole is the number of atoms contained in precisely 12 grams of the isotope 12C. A mole is denoted by the SI symbol mol. A mole is just a number. It is similar to how a dozen equals twelve and a pair equals two, one mole of a material equals 6.02 x 1023 particles. One mole of an element or a compound can be calculated or interpreted using atomic mass or molar mass. Take note that the atomic mass or molar mass can be expressed in grams per mole (g/mol). ELEMENT COMPOUND a. Atomic mass - expressed in grams of an 1 mol C = 12 g 1 mol H2O = 18 g element 1 mol Na = 23 g 1 mol NaCl = 58 g b. Molar mass - expressed in grams of a compound using the atomic masses of 1 mol Hg = 200 g 1 mol CO2 = 44 g Table 6.2.1. Example Interpretation of the Equivalent Value of each element 1 Mole in Atomic Mass and Molar Mass One mole is also used to talk about mass. One mole of a substance is equal to the atomic mass, or molar mass, in grams. For instance, the relative atomic mass of carbon is 12 g. It means that 1 mole of C is equal to 12 g. The diagram on the right serves as your guide on computing values relating to mole. Conversion of Element from Mass to the Number of Moles Sample Problem 2.1. How many numbers of moles are there in 15 grams of lithium? Given: mLi = 12.3g Li Required to Find: number of moles = ? mol Li Equation 𝟏 𝒎𝒐𝒍 𝒆𝒍𝒆𝒎𝒆𝒏𝒕 𝒏𝒖𝒎𝒃𝒆𝒓 𝒐𝒇 𝒎𝒐𝒍𝒆𝒔 = 𝒎𝒂𝒔𝒔 𝒐𝒇 𝒆𝒍𝒆𝒎𝒆𝒏𝒕 (𝒎) 𝒙 𝒂𝒕𝒐𝒎𝒊𝒄 𝒎𝒂𝒔𝒔 𝒆𝒍𝒆𝒎𝒆𝒏𝒕 Solution number of moles =15 g Li 𝒙 𝟏 𝒎𝒐𝒍 𝑳𝒊 7 g Li 𝟏𝟓 𝒎𝒐𝒍 𝑳𝒊 = 7 number of moles = 2.143 mol Li or 2.14 mol Li Answer There are 2.14 mol present in 15 g of Li. Conversion of Element from the Number of Moles to Mass Sample Problem 2.1. How much mass is contained in 2.14 mol of lithium? Given: number of moles = 2.14 mol Li Required to Find: mLi = ? g Li Equation mLi = 𝒎𝒂𝒔𝒔 𝒐𝒇 𝒆𝒍𝒆𝒎𝒆𝒏𝒕 (𝒎) 𝒙 𝟏 𝒎𝒐𝒍 𝒆𝒍𝒆𝒎𝒆𝒏𝒕 𝒂𝒕𝒐𝒎𝒊𝒄 𝒎𝒂𝒔𝒔 𝒆𝒍𝒆𝒎𝒆𝒏𝒕 Solution mLi = 2.14 mol Li 𝒙 𝟕𝒈 𝑳𝒊 1 mol Li 𝟏𝟒.𝟗𝟖 𝒈 𝑳𝒊 = 1 mLi = 14.98g Li or 15g Li Answer There are 14.98 g present in a 2.14 mol of Li. 30 Conversion of Compound from the Number of Moles to Mass Sample Problem 2.3. How many numbers of moles are there in 98.3 g of aluminum hydroxide, Al(OH)3? Given: m = 98.3g Al(OH)3 Required to Find: number of moles = ? mol Al(OH)3 Equation 𝟏 𝒎𝒐𝒍 𝒄𝒐𝒎𝒑𝒐𝒖𝒏𝒅 𝒏𝒖𝒎𝒃𝒆𝒓 𝒐𝒇 𝒎𝒐𝒍𝒆𝒔 = 𝒎𝒂𝒔𝒔 𝒐𝒇 𝒄𝒐𝒎𝒑𝒐𝒖𝒏𝒅 (𝒎) 𝒙 𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔 𝒄𝒐𝒎𝒑𝒐𝒖𝒏𝒅 Solution 1. Get the molar mass of the compound, Al(OH)3. Note: Atomic masses of each element: Al - 27g, O – 16g, H – 1g 2. Solve for the number of moles. 𝟏 𝒎𝒐𝒍 Al(OH)3 𝒏𝒖𝒎𝒃𝒆𝒓 𝒐𝒇 𝒎𝒐𝒍𝒆𝒔 = 98.3 g Al(OH)3 𝒙 78 g Al(OH)3 𝟗𝟖.𝟑 𝒎𝒐𝒍 Al(OH) = 3 78 Number of moles = 1.260 mol Al(OH)3 or 1.26 mol Al(OH)3 Answer There are 1.26 mol present in a 98.3 g of Al(OH)3. Activity 6.2.1. Mass Counted Activity 6.2.2. The Mole of Fortune State the masses of one mole of each Answer the following by calculating the substance. You may refer to the periodic number of moles or atomic mass in each of table and round it off to the nearest whole the given questions. number. Example: Asked: Mole Given:15 grams of lithium (Li) Answer: 2.143 mol Li or 2.14 mol Li ASKED GIVEN Mole 1. 22 g of gold (Au) Mole 2. 2.3 g of water (H2O) Mass 3. 9.8 mol of calcium (Ca) Mass 4. 3.3 mol of potassium sulfide (K2S) Lesson 6.3 Number of Particles The mole concept is a useful way of expressing the amount of a substance. The distinctiveness of a substance is defined by the quantity of each type of atoms or ions it contains. In dealing with p