Module 1 in Science 3b Chemistry PDF

***MODULE 1*** *In Science 3b* *Chemistry* Prof. Rita D. Gordo Subject Teacher **Module 1 in Science 3b ( Grade 9 ) : Chapter 8- Chemical Formulas and Equations** **PROF. RITA D. GORDO** **Asso. Professor III** **UEP- Laboratory High School** **College of Education** **Overview** In this...

***MODULE 1*** *In Science 3b* *Chemistry* Prof. Rita D. Gordo Subject Teacher **Module 1 in Science 3b ( Grade 9 ) : Chapter 8- Chemical Formulas and Equations** **PROF. RITA D. GORDO** **Asso. Professor III** **UEP- Laboratory High School** **College of Education** **Overview** In this module, you will learn how chemists measure number of atoms, molecules, or ions. The formula of a compound tells us the number of atoms in a molecule and the amount of element present. With this knowledge, scientist develop new drugs, cosmetics, fabrics, and dyes. You will also study the way atoms and molecules react and change. Chemical reactions keep on going around us all the time. The burning of fuel in the kitchen stove or gasoline in the car engine are examples of chemical reactions happen. It is important that you have some knowledge of chemical reactions, so that you have a better understanding of the processes occurring in your body and in your immediate environment. **Learning Outcomes** - Name and write formulas of binary compounds using the periodic table of elements. - Name and write the formulas of compounds given some common polyatomic ions and active metals - Explain the Law of Conservation of Mass given a balanced chemical equation - Describe the four kinds of evidences of chemical reaction - Distinguish between exothermic and endothermic chemical reaction - Write correctly balanced equation from a given word equation - Explain the steps involved in writing a balanced equation - Identify chemical reactions that occur in everyday life - Classify and balance chemical reactions as combination, decomposition , displacement, or double-displacement reactions **Pre-Test ( Please see attached sheet )** **Lesson 8.1 -- Elements, Compounds, and the Periodic Table of Elements** In your previous lesson, you have learned that elements are arranged according to groups (families) and periods (series). The elements of group 1A have an oxidation number of + 1.which means that these elements lose one electron to other elements to become chemically stable. On the other hand, elements of group 7A have an oxidation number of -1, which means that each element acquires an electron to become chemically stable. Can you tell the oxidation numbers of the elements in groups 2A, 3A, 5A. and 6A ? If you know the oxidation numbers of common elements based on their positions in the periodic table. You will find it easy to write chemical formulas of binary compounds. **Table8. 1** **OXIDATION NUMBERS OF COMMON ELEMENTS** 1A 2A 3A 4A 5A 6A 7A ------ ------ ------ ------ ------ ------ ------ (+1) (+2) (+3) (+4) (-3) (-2) (-1) (-4) Li Be B C N O F Na Mg Al Si P S Cl K Ca Se Br Sr **Lesson 8.2 -- Binary compounds** A **binary compounds** is a compound consist of two elements. **A Metal and a Nonmetal** The binary compounds formed by metals and nonmetals are usually ionic in nature. To write the correct formula of ionic compounds, you must know the ionic charges of the cations and anions. You can easily remember these if you know the positions of the elements in the periodic table. In writing a chemical formula, the sum of the total positive charges of the compound must be equal to the sum of the negative charges. The compound must have a net zero charge. When the charge on the positive ion is not equal to the charge on the negative ion, a **subscript** which is a number written below the chemical symbol must be used to balance the negative and positive charges. In most cases, the positive ion is first written followed by the negative ion. Consider the following examples to illustrate formula writing: 1. K^+^ and Cl^-^ ( K^+^ ) ( Cl^-^ ) or KCl (+ 1) + (-1) = 0 2. Ca^+2^ and Br^-1^ (Ca^+2^ ) ( Br^-1^ ) or CaBr~2~ (+2 ) + 2 ( -1) = 0 **Note** that in the examples given, the net ionic charge of the positive ions and the negative ions is equal to zero. There is a short-cut method to get the correct formula of a binary compound. Simply write each ion with its charge (starting with the positive ion) then switch the charges ( but not the plus or minus sign), and write them as subscripts. You can perfect this process through practice. Subscripts are always reduced to lowest terms. Remember that the "crisscross" method is based on the transfer of electrons and the conservation of charges. The total positive charges must equal the total negative charges. The net charge on the compound is always zero. In naming binary compounds, the metal is named first, followed by the nonmetal. Change the ending of the nonmetal into --ide. Example Calcium Bromide, Sodium Chloride, Potassium Iodide, Magnesium Nitride, Sodium Oxide, Aluminum Sulfide. The example above are binary compounds contain metals with fixed oxidation numbers. How about metals with variable oxidation numbers (transition metals). The same procedure is followed in writing the formula. However, in naming, the oxidation number must be specified. There are two methods of specifying the oxidation number- one uses is the stock system and the other make use the Latin names of the metals. In the **stock system** the oxidation number of the metal is indicated by a Roman numeral enclosed in a parenthesis. However, when using the Latin name of the metal suffixes-**ous** and -**ic** are used to denote the lower and higher states respectively. Consider the following examples: Cu^+1^ O^-2^ → Cu~2~O → Cuprous Oxide or Copper (1) oxide Fe^+3^ O^-2^ → Fe~2~O~3~ → Ferric Oxide or Iron (III) Oxide Table 8.2 names of common cations using the two system **Table 8-2** **Common Cations With Variable Oxidation Numbers** Pb^+2^ Plumbous or Lead (II) Fe^+3^ Ferric or Iron (III) --------- ----------------------- -------- --------------------------- Pb^+4^ Plumbic or Lead (IV) Hg^+1^ Mercurous or Mercury (I) Cu ^+1^ Cuprous or Copper(I) Hg^+2^ Mercurous or Mercury (II) Cu^+2^ Cupric or Copper (II) Sn^+2^ Stannous or Tin(II) Fe^+2^ Ferrous or Iron(II) Sn^+4^ Stannic or Tin(IV) **A Nonmetal and a Nonmetal** The type of bond between two nonmetals is generally covalent in nature. Let us recall our lesson in grade 8 about Lewis Electron Dot Formula. How do you determine the formula of covalent compounds using the Lewis Electron dot Structure? Naming the covalent compounds having two nonmetallic elements is similar to naming ionic compounds, except that prefixes are used. Refer to table 8-3. **Table 8-3** **GREEK PREFIXES USED IN NAMING COVALENT COMPOUNDS** **Greek Prefix** **Number** **Greek Prefix** **Number** ------------------ ------------ ------------------ ------------ Mono 1 hexa- 6 di- 2 hepta- 7 tri- 3 octa- 8 tetra- 4 nona- 9 penta- 5 deca- 10 An example when carbon with oxygen to form an oxide, carbon is named first since it is a less electronegative element (Group 4) than oxygen (Group 6). The prefix mono- is used only in the second element. The suffix-ide is then added to the stem of the name of the more electronegative atom. The two compounds carbon form with oxygen are: CO → Carbon monoxide CO~2~ →Carbon dioxide **Hydrogen** **and a Nonmetal** A binary compound which is composed of hydrogen and a more electronegative element is named like any other binary compound of nonmetals. For example, HCl is called hydrogen chloride. The word hydrogen comes first; then the second word is made affixing the suffix --ide to the root of the nonmetal. Hydrogen may also be combined with fluorine, bromine, and iodine to form HF, HBr, and HI. The names are hydrogen fluoride, hydrogen bromide, and hydrogen iodide, respectively. When these substances are dissolved in water, they become aqueous acids. A different method of naming the substance is applied. The prefix **hydro-** is attached to the root word of the nonmetal and the suffix-**ic** is added. The word acid becomes the last term. Example H~2~S~(aq\ )~ → hydrosulfuric acid HCl~(aq)~ → hydrochloric acid **Lesson 8.3 -- Ternary Compounds** **Metals with Polyatomic Ions** Polyatomic compounds are formed in the same way as binary compounds. A **polyatomic ion** is a stable group of atoms that carries an overall electrical charges. The atoms in a polyatomic ion are bonded together by covalent bonds. One or more atoms are group that carry a positive or negative ionic charge. Therefore, the group as a whole has an ionic charge. Table 8.4 list of formulas and charges of common polyatomic ions. These polyatomic ions form ionic bonds in the same manner as single ions do. Parenthesis are placed around the polyatomic ion and the subscript is written outside the parenthesis whenever a multiple of the polyatomic ion is necessary. For example to determine the formula of magnesium hydroxide. The following steps must be followed: 1. Mg^+2^ and OH^-1^ 2. Mg^+2^ ( OH) 1. 2 3. Mg(OH)~2~ The total positive and negative charges must be equal. Placing the hydroxide ion in the parenthesis , the correct formula is Mg(OH)~2~. The parenthesis are not used when a single polyatomic ion is present, as in potassium chlorate, calcium carbonate and sodium hydrogen carbonate. KClO~3~ CaCO~3~ NaHCO~3~ **Table 8.4** **Formulas and Charges of Common Polyatomic Ions** **Name** **Formula** **Charge** **Name** **Formula** **Charge** -------------- -------------- ------------ ------------- ------------- ------------ Ammonium NH~4~ -1 Chlorate ClO~3~ -1 Nitrite NO~2~ -1 Chromite CrO~2~ -1 Nitrate NO~3~ -1 Cyanide CN -1 Hypochlorite ClO -1 Perchlorate ClO~4~ -1 Chlorite ClO~3~ -1 Hydroxide OH -1 Bicarbonate HCO~3~ -1 Sulfite SO~3~ -2 Acetate C~2~H~3~O~2~ -1 Sulfate SO~4~ -2 **Note:** Others are found in your periodic table of elements Example : MgNO~3~ Magnesium Nitrate Al(OH)~3~ Aluminum Hydroxide **Oxyacids and Oxyanions** You have learned how binary acids are formed. Now you will learn that there are also acids containing three or more elements. Oxygen is always present in this kind of acids, which are referred to as binary acids or oxyacids. Table 8-5 gives the partial list of common oxyacids and their corresponding anions. Try your best to memorize the formula of common oxy acids. **Table 8-5** **COMMON NAMES OF OXYACIDS AND THEIR ANION** **Name of Acid** **Formula of Acid** **Name of Anion** **Formula of Anion** ------------------ --------------------- ------------------- ---------------------- Boric Acid H~2~BO~3~ Borate BO~3~^-3^ Carbonic Acid H~2~CO~3~ Carbonate CO~3~^-2^ Nitric Acid HNO~3~ Nitrate NO~3~^-1^ Nitrous Acid HNO~2~ Nitrite NO~2~^-1^ Phosphoric Acid H~3~PO~4~ Phosphate PO~4~^-3^ Phosphorous Acid H~3~PO~3~ Phosphite PO~3~^-3^ Sulfuric Acid H~2~SO~4~ Sulfate SO~4~^-2^ **Lesson 8.4- Chemical Reactions** Chemical reactions are important in many professions. From the time we get up in the morning to the time we sleep at night. Chemical changes are taking place, within us and outside of us. Plants grow through photosynthesis, foods that we eat are digested by the body, metals corrode, raw are being converted to useful products, new medicines are being developed, more versatile materials are being made. **Evidences of Chemical Reactions** Various chemical changes that occur around us have significant effects to our environment and consequently to our health. Chemical changes occurring in industries results to products that are useful to us, The wastes we throw continue to undergo chemical changes and has an impact on our well-being as well. The irresponsible use of fertilizers, herbicides and pesticides have negatively affected plants and aquatic life. We continue to pollute the atmosphere with vehicle and industrial gas emissions. When you drop an Alka-Seltzer tablet in a glass of water, bubbles are formed. The release of gas indicates that a chemical reaction is taking place. Similarly, when a piece of zinc metal is added to hydrochloric acid, hydrogen gas bubbles are released. Many chemical reactions give off some form of energy. Light energy was given off when magnesium ribbon was burned. Heat energy is also released. In burning newspaper, What form of energy is released? Can you give similar examples? **Energy and Chemical Reactions** Some chemical reactions release energy, others absorb energy.. The law of conservation of mass and energy applies to chemical reactions. Chemical reactions observe energy conservation. This means that every time a chemical reaction takes place, energy is neither created nor destroyed. Any energy released in a reaction was present in the chemical bonds of the original substances. Any absorbed in a reaction becomes part of the chemical bonds of the new substances. There are several chemical reactions that release energy. These types of reactions are called **exothermic reactions**. Whenever you cook your food or use your car energy is released. All exothermic reactions fit this chemical pattern. Exothermic reactions most often produce energy in the form of heat, light or electricity. Chemical reactions that absorb energy are called **endothermic reactions.** These types of reaction are less common. For example, when electric energy is added to water, the water molecules absorb the electricity. The water molecules undergo a chemical change, breaking apart into hydrogen and oxygen. The chemical description for an endothermic is shown in this pattern. Chemical reactions that takes place daily in our environment and in the human body can be classified as: 1. Composition or direct combination reaction 2. Decomposition or breaking-apart reaction 3. Single-displacement reaction 4. Double-displacement reaction Details of these will be discussed in lesson 8.5. **Lesson 8.5 -- Chemical Equations** A chemical equation is a shorthand of a quantitative description of a chemical reaction. With your knowledge of formulas of compounds, you are now ready to write chemical equations. Basic terms and symbols are necessary in writing equations. The reactants are always written on the left side and products on the right. An arrow , →, double arrow , or an equal sign = usually separates. Table 8-6 gives you the different symbols used in chemical equations: **Table 8-6** **SYMBOLS USED IN EQUATIONS** Let us take the reaction of hydrogen with oxygen for example. Burning of hydrogen gas in the presence of oxygen results in a big explosion and the production of water. To write the chemical equation for this reaction, 1. Identify the reactants and the product. A **reactant** is a substance used up during a chemical reaction. A **product** is a substance formed after a chemical reaction. In the example given, hydrogen burns in the presence of oxygen. The reactants are hydrogen and oxygen. This results in the production of water, which is the product. 2. Write the **word equation**; the **reactants** to the left and the **product** to the right. Hydrogen + oxygen → water 3. Replace the names of the substances with their chemical symbols or chemical formula. H + O → H~2~O Unfortunately hydrogen and oxygen atoms are very rare in nature. They are found as **diatomic** molecules instead. **Diatomic** molecules are molecules that are made of only two atoms. ( See table 8-7 for other diatomic molecules.) Thus, hydrogen gas and oxygen are written as H~2~ and O~2~ respectively. Correcting the equation we get: 2H~2~ + O~2~ → 2H~2~O **Table 8-7** **Diatomic Molecules** **Name of Elements** **Formula** ---------------------- ------------- Hydrogen H~2~ Nitrogen N~2~ Oxygen O~2~ Fluorine F~2~ Chlorine Cl~2~ Bromine Br~2~ Iodine I~2~ Sample Exercises: 1. Gaseous hydrogen combines with gaseous nitrogen to produce gaseous ammonia. N~2~ + 3 H~2(g)~ → 2 NH~3(g)~ 2. Solid aluminum oxide forms instantly when fresh aluminum metal is exposed to oxygen in the air. 4 Al~(s)~ + 3O~2(g)~ → 2Al~2~O~3(s)~ **Types of Chemical Reaction** A. **COMBINATION REACTION** A combination reaction takes place when two or more substances combine to form a more complex molecules. The general equation for combination reaction is A + B→AB, where A and B are elements or compounds. When iron fillings are burned with yellow to form black iron (II) sulfide, a combination reaction occurs. Fe + S~2~ → Fe S~2(g)~ Depending on the reactants involved, a combination reaction may follow different patterns. 1. Metal + nonmetal → Salt Fe ~(s)~ + S ~(g)~ → FeS~(S)~ 2. Nonmetal oxide + water → oxyacid SO~2(g)~ + H~2~O~(l)~ → H~2~SO~4(aq)~ 3. Metal + oxygen → metal oxide 2Mg~(s)~ + O~2~ → 2 MgO~(S)~ 4. Metal oxide + water → metal hydroxide 5. Nonmetal + oxygen → nonmetal oxide 6. Metal oxide + nonmetal oxide → Salt MgO + SO~3~ → MgSO~4~ B. **DECOMPOSITION REACTION** In a decomposition reaction, one substance, is usually a compound, breaks down to form two or more substances which may be elements or compounds. A **catalyst** , which may be in the form of heat, electricity, radiation, or even enzyme can aid or bring about decomposition reaction. The general equation for decomposition reaction is AB → A + B, where A and B are elements or compounds. It is important that you know the different types of decomposition reactions in order to predict the possible products that will be formed. Consider the following reactions: 1. When an oxide is heated, oxygen generally is given off as one of the products. 2HgO~(s)~ → 2 Hg~(l)~ + O~2(g)~ 2. Some carbonates( like limestone ), when heated, decomposition to yield carbon dioxide. Caco~3(s)~ → CaO~(S)~ + CO~2(g)~ 3. Some compounds such as baking soda ( sodium hydrogen carbonate or sodium bicarbonate) when heated decompose yield carbon dioxide, water, and carbonate salt. Baking soda can be used to put out a flame because of the production of carbon dioxide. 2 NaHCO~3(s)~ → Na~2~CO~3(s)~ + H~2~ O~(l)~ + CO~2(g)~ 4. **Hydrates** are substances that contain one or more molecules of water for each formula unit. These compounds, upon application of heat, readily decompose. The water is driven off, leaving the anhydrous(without water) salt. Consider the following examples below. CuSO~4~.5H~2~O → CuSO~4(s)~ + 5 H~2~O~(g)~ C. **SINGLE-DISPLACEMENT REACTION** In a **single-displacement** ( also known as substitution or replacement ) reaction, one element reacts by replacing another element in a compound. This reaction can be represented by the general equation: A + BC → AC + B To determine which element will be displaced or substituted , you need to know the electromotive or activity series. The **electromotive or activity series** is an arrangement of elements, which shows what metal can displace another metal from an aqueous solution of its salt or acid. You should be able to interpret and use this series so that you can complete And predict chemical reactions involving replacement reactions. **Table 8-8** **The Activity Series** **Table 8-9** **Activity Series of Halogens** As you can see in the activity series , zinc is higher than nickel, and therefore it can displace nickel. Sodium can replace one hydrogen ion in water and magnesium can replace one hydrogen ion from sulfuric acid. Note that hydrogen is a diatomic molecule. Zn~(s)~ + NiSO~4(aq)~ → ZnSO~4(aq)~ + Ni~(s)~ Na~(s)~ + H~2~O~(l)~ → Mg(s) + H~2~SO~4(aq)~ → Another displacement reaction involves a nonmetal displacing a nonmetal ion from its salt or acid. Consider the following reaction. F~2(g)~ + 2 NaCl~(aq)~ → 2Na F~(aq)~ + Cl~2(aq)~ Note that halogens are diatomic molecules in the free state. Cl~2~ + Na I → Zn + AgNO~3~ → **D. DOUBLE-DISPLACEMENT OR IONIC REACTION** Two compounds are involved in **double-displacement or ionic reaction.** The positive ion of one compound exchanges with the positive cation of another compound. This reaction may also be called **metathesis** or double replacement reaction. Note that in this reaction, there are four separate species, namely, A, B, C, and D, whereas in single --displacement reactions there are only three. The general equation is AB + CD → AD + CB Where : A and C = cations B and D = anions Consider the following reactions: 1. A salt and a base Ca(NO~3~)~2(aq)~ + 2 NaOH~(Aaq)~ → 2 NaNO~3(aq)~ + Ca(OH)~2(aq)~ 2. Two salts 3. A salt and an acid Ba(NO~3~)~2(aq)~ + H~2~SO~4(aq)~ → BaSO~4(s)~ + 2 HNO~3(aq)~ 4. Metal Carbonate and an acid 5. An acid and a base 2HCl~(aq)~ + Mg(OH)~2(aq)~ → MgCl~2(aq)~ + H~2~O~(l)~ Note that this reaction occurs in the stomach when milk of magnesia, Mg(OH)~2~, is used as an antacid. This is known as a **neutralization** reaction. **Lesson 8.6 -- Balancing Chemical Equations** Chemist determine the identity of the reactants and products of a reaction by **experimental observation.** For example, when you heat a small amount of mercury(II) oxide in a test tube, the amorphous orange compound turns brown then black and finally disappears. Moisture is also observed near the mouth of the test tube. The grayish substance left clinging at the middle portion of the test tube is mercury and the smoke seen is oxygen. In balancing a chemical equation, the following steps may be used: 1. Read the description of the chemical reaction, What are the reactants, the products, and their states? Write the correct formula for each substance. 2. Write unbalanced equation that summarizes the information from step 1. 3. Balance the equation using inspection method. Add coefficients before the formula of each substance. Do not change the identities ( formulas ) of any of the reactants or products. 4. Check to see that the coefficients used give the same number of each of atom on both sides of the equation. Be sure that the coefficients used are the smallest integers that give the balanced equation. To balance the chemical equation, H~2(g)~ + O~2(g)~ → H~2~O~(g)~ Step 1. Count the number of each element in the reactant side and in the product side. H~2(g)~ + O~2(g)~ → H~2~O~(l)~ ( unbalanced ) Reactants Products ---------------- ----------- ---------- Hydrogen ( H ) 2 2 Oxygen ( O ) 2 1 Notice that the number of oxygen atoms in the reactant side is greater than the product side. Step 2. Adjust the coefficients of the formula to balance the equation. In balancing an equation, the smallest whole number coefficients are generally used. What coefficient should you place before H~2~O to balance the equation? H~2(g)~ + O~2(g)~ → 2 H~2~O~(l)~ Placing a " 2 " in front of H~2~O increases the number of hydrogen atoms and oxygen atoms in the product. There are now 4 hydrogen atoms and 2 oxygen atoms in the product. **Atom** **Reactant** **Product** ---------- -------------- ------------- H 2 4 O 2 2 This time, there are less hydrogen atoms in the reactant that there are in the product. What number should you place before H~2~ to balance the equation? Try placing a " 2 ". 2 H~2(g)~ + O~2(g)~ → 2 H~2~O~(l)~ **Atom** **Reactant** **Product** ---------- -------------- ------------- H 2(2) 4 O 2 Step 3. Check if the equation is completely balanced. **Atom** **Reactant** **Product** ---------- -------------- ------------- H 4 4 O 2 2 Remember that choosing the right coefficient is done by trial and error. If you carefully check the numbers of each atom every time a coefficient is charged, you reduce the trial and eliminate the error. How would you balance the following chemical equations? Practice Exercises 1. Na + S~8~ → Na~2~S 2. AlCl~3~ + LiH → AlH~3~ + LiCl 3. NaHCO~3~ + H~2~SO~4~ → Na~2~SO~4~ + CO~2~ + H~2~O **Feedback** How did you go on so far with this module? Were the lesson helped you attain the focus points intended for the lesson. Well, these are simple concepts in formula writing, chemical equation and balancing equation that are explained clearly that leads to a more meaningful learning and much better retention of what has been learned. **Summary** - Chemical changes occur constantly in everyday life as energy changes, color changes, formation of precipitate and evolution of gas. - Chemical formulas represent compounds. They indicate the number of atoms of each element present in a compound. - In formula writing, the total positive charges must be equal to the total negative charges. The net charge on the compound is zero. - A chemical equation is a shorthand method of expressing a chemical chang using symbols and formulas. - Chemical equations are balanced based on the Law of Conservation of Matter (mass) Which states that the matter(mass) is neither created nor destroyed. The number of atoms or moles of atoms of each element must be the same on both sides of the equation. - The types of chemical reactions are combination, decomposition, single-displacement and double-displacement reaction. - In simple binary compounds, the more positive element is named before the more negative element that is suffixed by --**ide**. - Names of polyatomic ions is often end in --ate, except when there are several anions containing the same elements, the name of the one with less oxygen ends in --**ite.** - When two nonmetals form several different covalent compounds, prefixes such as mono^-^ , di^-^ , tri^-^ , etc. re used to indicate numbers of atoms of a particular element. - The Law of conservation of mass states that the total mass of reactants is equal to the total mass of products when chemical reaction takes place. **Suggested Readings** - Ultraviolet Radiation - Chlorofluorocarbons - Antacids and Laxative **References** Mendoza, Estrella E., Religioso, Teresita F. You and The Natural World, Chemistry Phoenix Publishing House, Inc., Quezon City, 2008. Madriag, Estrella A., et al. Science Links- Worktext for Scientific and Technological Literacy., Rex Book Store, Manila, 2017. Mapa, Amelia P. PhD., Fidelino, Trinidad B., Rabago, Lilia M. , PhD. Chemistry Textbook, Science and Technology, Third Year 2001.