Physical and Chemical Changes

Transcription

1 Physical and Chemical Changes Gezahegn Chaka, Ph.D. Objectives To observe physical and chemical changes. To identify physical and chemical changes. To write balanced chemical equations to represent these physical and chemical changes. Introduction In a physical change the appearance of a substance changes but its composition and identity remains unaltered. Examples of physical changes include freezing of water to form ice, dissolving of sugar in a hot cup of tea, and the filing of metal to produce powder or filings. There are a number of evidences that suggest whether a physical change has taken place or not. Some of the evidences are changes in state or dissolution of a solid. In a chemical reaction a change in the composition and identity of a substance occurs. The observation of chemical changes (chemical reactions) and the proper descriptions of these changes are important concerns of chemistry. There are many clues an observer might note that would suggest a chemical change is taking place, such as the formation of a gas or a precipitate, a color change, the disappearance of a solid, or the warming or cooling of a reaction mixture caused by the evolution or absorption of heat. Another method of describing a physical or a chemical change is with a chemical equation. For physical changes, symbols written in parenthesis can be used to indicate the changes in the physical state of substances. In such equations, (aq) means in aqueous solution, (s) is for solid, (l) is for liquid, and (g) is for gas. For chemical changes, the chemical equation is a statement that represents the reacting species (the reactants) and those produced (the products) with formulas and symbols. A proper chemical equation is based on chemical evidence. Simply writing an equation does not mean the reaction takes place.

2 Also to be in agreement with the law of conservation of mass, the equation must be balanced. Only balanced equations can be used in chemical calculations. Predicting the Products of Chemical Reactions Obviously, the most difficult step is to predict the products of a chemical reaction. Rather than trying to memorize each and every chemical reaction, it is helpful to try and understand a few trends. A large number of inorganic chemical reactions may be classified into one of the four general categories: single displacement, double displacement, combination, and decomposition reactions. Single displacement reactions involve the reaction of an element with a compound such that the element replaces one of the elements in the compound (setting it free) while combining with the other element. In double displacement reactions, the positive and the negative ions simply exchange partners. Such reactions occur in solution if one of the three criteria is satisfied: (a) formation of a precipitate, (b) formation of a gas, or (c) formation of a weak electrolyte. For example, cations (positive ions) can only be paired with anions (negative ions). Two cations will not pair up; neither will two anions. In double replacement reactions, therefore, the cations and anions just switch with each other. Combination reactions consist of the direct union of two simple substances to form a third substance that is more complex. On the other hand, a compound often decomposes only after the addition of energy (heat, light, electricity, etc.); otherwise, it would not be possible to obtain the product. The products of a decomposition reaction are often small, stable molecules such as carbon dioxide, water, and oxygen gas which are familiar to you. In any case, you must account for each type of atom since no atom will be created or destroyed in a chemical reaction. The seven elements which exist as diatomic molecules (H2, N2, O2, F2, Cl2, Br2, and I2) need to be represented as diatomic in a chemical equation.

3 A splint test is used to identify gases produced in a chemical reaction. A wood splint is ignited and then extinguished. While the splint is still glowing, it is inserted into the test tube where the reaction has occurred. If the gas produced by the reaction is oxygen the splint will momentarily glow more brightly and may even ignite. If hydrogen gas is present, there will be an audible pop as the glowing splint is inserted into the test tube. If the splint is completely extinguished, then the gas produced must be a nonflammable gas such as CO2, N2, or Cl2. Chemicals copper(ii) sulfate silver nitrate iron magnesium zinc potassium chlorate hydrochloric acid manganese(iv) oxide sodium chloride calcium carbonate iodine Considerations before you begin the experiment: Observe what changes occur in each of the following procedures and classify the change as a physical change or a chemical change. Represent the changes in correctly balanced chemical equations. What evidence suggests that a physical or a chemical reaction took place in each procedure?

4 Procedures GOGGLES MUST BE WORN WHILE WORKING IN THE LABORATORY 1. Place about 10 ml of the blue CuSO4 solution in a small beaker. Take a piece of steel wool, which is composed mainly of elemental iron, and dip it into the blue solution; hold it with tongs in the solution for about 30 seconds and take it out. Examine the coating on the steel wool. Record your observation. Now put the steel wool in the solution and let it remain for at least 15 minutes while you perform the other procedures. Record your observation. 2. Working in the fume-hood, place a small piece of zinc in a test tube. Add enough hydrochloric acid to just cover it. Perform a splint test on any gases that are produced. Record your observation. 3. Working in the fume-hood, place on a watch glass a few drops of the liquid from the test tube in Procedure 2. Place the watch glass on wire gauze and gently heat until the water evaporates. What is the residue (compound) on the watch glass? Record your observation. 4. Put 10 drops of NaCl solution in a test tube. Add 5 drops of AgNO3 solution. Record your observation. 5. Working in the fume-hood, add several drops of dilute hydrochloric acid (HCl), drop by drop, to a test tube containing a tiny amount of calcium carbonate (CaCO3). Perform a splint test on any gases that are produced. Record your observation. 6. Take a piece of clean copper wire (use steel wool to clean wire). Hold one end with tongs, heat the other end to red hot, and let the wire cool. Record your observation. 7. Working in the fume-hood, ignite a small strip of magnesium ribbon by holding it with tongs in the flame of the burner. Do not look directly at the burning magnesium flame. Record your observation. 8. Mix a small quantity of solid KClO3 with solid MnO2 in a test tube and attach the test tube to a stand with a clamp at a 45º angle. Working in the fume-hood, heat the content of the test tube over the Bunsen burner flame. Perform a splint test on any gases that are produced. Record your observation.

5 9. Place a small quantity of CaCO3 in a test tube and attach the test tube to a stand with a clamp at a 45º angle. Working in the fume-hood, heat the content of the test tube over the Bunsen burner flame. Perform a splint test on any gases that are produced. Record your observation. 10. Working in the fume-hood, support a 150 ml clean and dry beaker with a wire gauze on a ring stand. Place 0.5 g of iodine in the beaker. Place an evaporating dish on top of the beaker. Place a small amount of ice in evaporating dish. Heat the iodine in the 150 ml beaker very gently with low Bunsen burner flame. Observe the process inside the beaker and at the bottom of the evaporating dish. Record your observation. References 1. T. G. Greco; L. H. Rickard; G. S. Weiss; Eds.; Experiments in general chemistry: Principles and Modern Applications. Ninth edition, Pearson Prentice Hall, C. H. Henrickson; L. C. Byrd; N. W. Hunter; A Laboratory for General, Organic, and Biochemistry, Fifth Edition, McGraw Hill, Southeastern Louisiana University; Department of Chemistry and Physics, Chemical Reactions and Chemical Equations, Chemistry 103, spring 2006.

6 LABORATORY REPORT SHEET Physical and Chemical Changes Name Date Procedure Observations Physical Change Chemical Change Gas Produced (if any) Conclusions