Module-2023-Chem-for-Engineering-v3-MIDTERMS.pdf

.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 1 UNIT 1 INTRODUCTION TO CHEMISTRY: NATURE OF SCIENCE, SCIENTIFIC METHOD AND MEASUREMENTS THE NATURE OF SCIENCE Chemistry has been called the “central science” because it is important to so many other fields of scientific study. This text and the course are designed to help connect pieces of information already picked up, increase the understanding of chemical concepts, and give a more coherent and systematic picture of chemistry. The ultimate goal is to help the students understand the natural world. This type of perspective of the world is what enables chemists and engineers to devise strategies for refining metals from their ores, as well as to approach the many other applied problems that will be explored. Chemistry is the branch of science concerned with the properties, composition, and structure of substances and the changes they undergo when they combine or react under specified conditions. The following are some examples of how chemistry is important in our daily lives: 1. It provides men the basic necessities in life like shelter, food and clothing. 2. It provides men luxuries in life like convenient transportation, advance and fast means of communication, use of computers, use of cosmetics, perfumes, etc… 3. Researchers in chemistry help improve the synthesis of chemicals needed to combat disease such as antibiotics, anesthetics, antiseptics, hormones and others. 4. It explains the composition of major classes of foods and their nutritional values 5. It helps in the advancement of scientific and technological studies like telecommunication systems and computer studies. 6. It enhances the awareness on how the body works and on the chemical changes that occur within the body system. 7. It leads to the discovery of organophosphorus pesticides which along with other pesticides, reduce crop losses. Branches of Chemistry 1. Inorganic Chemistry- the study of all the elements and their compounds with the exception of carbon and its compounds investigates the characteristics of substances that are not organic, such as nonliving matter and minerals found in the earth's crust. 2. Organic Chemistry- Branch of chemistry dealing with compounds of carbon. 3. Analytical Chemistry- This kind of chemistry deals mostly with the composition of substances. Collection of techniques that allows exact laboratory examination of a given sample of material. Chemists perform qualitative analysis or substances in a sample & quantitative analysis for the amount of each substance. Qualitative Chemistry- the atoms and molecules present are identified, with particular attention to trace elements........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 1 Quantitative Chemistry- the exact weight of each constituent is obtained as well 4. Biochemistry- encompasses the study of the chemical nature of living material and of the chemical transformations that occur within it. A science that is concerned with the composition and changes in the formation of living species. This type of chemistry utilizes the concepts of organic and physical chemistry to make the world of living organisms seem much clearer. 5. Physical Chemistry- is concerned with the physical properties of materials, such as their electrical and magnetic behavior and their interaction with electromagnetic fields. This chemistry is defined as dealing with the relations between the physical properties of substances and their chemical formations along with their changes. SCIENTIFIC METHOD The number of steps can vary from one description to another (which mainly happens when data and analysis are separated into separate steps), however, this is a fairly standard list of the six scientific method steps that you are expected to know for any science class: 1. Purpose/Question Ask a question. 2. Research Conduct background research. Write down sources to cite references. 3. Hypothesis Propose a hypothesis. This is a sort of educated guess about what is expected. It is a statement used to predict the outcome of an experiment. Usually, a hypothesis is written in terms of cause and effect. Alternatively, it may describe the relationship between two phenomena. 4. Experiment Design and perform an experiment to test the hypothesis. An experiment has an independent and dependent variable. 5. Data/Analysis Record observations and analyze the meaning of the data. 6. Conclusion Conclude whether to accept or reject hypothesis. There is no right or wrong outcome to an experiment, so either result is fine. Whether the hypothesis is accepted or rejected, it may be revised to form a new one for a future experiment. The History of Chemistry Timeline of Development I. Prehistoric era. The history of chemistry started when people stated to use of fire, cook food and baked pottery, production of wine and use of cosmetics II. Greek Civilization.  Thales assumed that all matter was derived from water  Democritus said the ATOM is the simplest unit of matter  Empedocles said that all matter was composed of four elements: fire, air, water, and earth.  Aristotle described the four qualities were found in nature: heat, cold, moisture, and dryness........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 2 III. Beginning of Christian Era - End of 17th Century (Alchemy “al chemia”)  Alchemists attempted to transmute cheap metals to gold. This is the precursor of Chemistry.  Paracelsus – Auroleus Phillipus Theostratus Bombastus von Hohenheim. Searched for medicine to cure sickness. He was credited with the introduction of opium and mercury into the arsenal of medicine. His works also indicate an advanced knowledge of the science and principles of magnetism. The precursor of chemical pharmacology and therapeutics and the most original medical thinker of the sixteenth century."  Galileo – introduced accurate measurements  Robert Boyle - disproved Aristotle’s four elements theory “Skeptical Chemist”. He conceptualized the Boyle’s Law. IV. End of 17th Century -Mid 19th Century (Traditional Chemistry)  George Ernst Stahl – Phlogiston theory  Joseph Priestley - Isolated oxygen by heating mercuric oxide  Jan Baptista van Helmont - kinds of air-like materials “gas” Carbon Dioxide  Antoine Lavoisier - He disproved the phlogiston theory. He is the Father of modern chemistry.  Mikhail Vasilyevich Lomonosov - Lomonosov rejected the phlogiston theory, and anticipated the kinetic theory of gases. Lomonosov was the first person to record the freezing of mercury, and to hypothesize the existence of an atmosphere on Venus. He demonstrated the organic origin of soil, peat, coal, petroleum, and amber. In 1745 he published a catalogue of over 3,000 minerals, and in 1760 he explained the formation of icebergs.  John Dalton - developed the Atomic Theory  Rudjer Joseph Boscovich - developed the modern atomic theory V. Mid 19th Century – Present (Modern or 20th Century Chemistry)  Heinrich Geissler – developed the first vacuum tube  William Crookes – discovered the cathode rays  Eugene Goldstein - discovered proton  Michael Faraday- Invented the electric motor  Wilhelm Conrad Röntgen – discovered x – rays  Henri Becquerel – discovered spontaneous radioactivity  Joseph John Thomson – discovered the electron and its properties  Robert Andrews Millikan – determined the mass of an electron  James Chadwick – discovered the neutron.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 3  Ernest Rutherford – determined the three types of radioactivity  Marie and Pierre Curie – determined the radioactive properties emitted by uranium, thorium, radium & polonium  Niels Bohr - Proposed that electrons could only reside in certain energy levels or quanta  Enrico Fermi - neutron bombardment & nuclear fission  Dmitri Ivanovich Mendeleev – developed the Periodic Law and the properties of the chemical elements. The Father of the Periodic Table.  Henry Moseley - Determined the atomic numbers of all the known elements. Arranged the periodic table according to increasing atomic numbers. SYSTEMS OF MEASUREMENTS I. Systems of Measurement A. The English System Not easy to deal with since each unit has a corresponding value that is used when converting from one unit to another. B. The Metric System It is used in all scientific studies. It is in multiples of tens and is easier to use. It is established and modified when necessary by international agreement. C. The International System of Units (Le Système International d’Unitès, SI) It is a more improved version of the metric system and is founded on seven base units and two supplementary units. Measurement Unit Symbol Base Units Length meter m Mass kilogram kg Time second s electric current ampere A Temperature Kelvin K amount of substance mole mol luminous intensity candela cd Supplementary Units plane angle radian rad solid angle steradian sr Multiples or fractions of base units are indicated by the use of prefixes Prefix Abbreviation Factor Scientific Notation tera- T- 1,000,000,000,000 1012 giga- G- 1,000,000,000 109 mega- M- 1,000,000 106.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 4 kilo- k- 1,000 103 hecto- h- 100 102 deka- da- 10 10 deci- d- 0.1 10-1 centi- c- 0.01 10-2 milli- m- 0.001 10-3 micro- µ- 0.000 001 10-6 nano- n- 0.000 000 001 10-9 pico- p- 0.000 000 000 001 10-12 femto- f- 0.000 000 000 000 001 10-15 atto- a- 0.000 000 000 000 000 001 10-18 - other SI units, derived units, are obtained from the base units by algebraic combination Volume – refers to the amount of space occupied by an object 1. Regular solids 2. Cylinder 3. Irregular solids Mass – refers to the amount of matter present in an object. Mass is used interchangeably with weight in Chemistry. Although, the two terms differ technically when used in Physics. Density – it is an intrinsic property of matter that is defined as the mass of matter present per unit volume, expressed in the following formula: M D= V Where M = mass and V = volume SCIENTIFIC NOTATION - very large and very small numbers are frequently encountered in scientific studies  scientific notation is used to simplify the handling of these cumbersome values - when employed, the value is expressed in the form a x 10n where a = the decimal part, a number with one digit to the left of the decimal point and all others to the right n = the exponent of 10, positive or negative integer or zero.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 5  number can be converted into this form by moving the decimal point until there is only one nonzero digit to the left of it  for each place the decimal point is moved to the left, n is increased by one  for each place the decimal point is moved to the right, n is decreased by one for example, 0.082057 = 8.2057 x 10-2 29,979,000,000 = 2.9979 x 1010 96484.6 = 9.64846 x 104 1. Multiplication - decimal parts are multiplied and the exponents of 10 are added algebraically ex. (3.0 x 105)(2.0 x 102) = (3.0 x 2.0) x 105+2 = 6.0 x 107 2. Division - decimal parts are divided, exponent of 10 in the denominator is algebraically subtracted from exponent of 10 in the numerator ex. 6.89 x 10-7 = 6.89 x 10(-7)-(+3) 3.36 x 103 3.36 = 2.05 x 10-10 3. Addition and Subtraction - numbers must be expressed with the same power of 10 - the answer, which has the same power of 10, is found by adding or subtracting the decimal parts ex. (6.25 x 103) + (3.0 x 102) = (6.25 x 103) + (0.30 x 103) = 6.55 x 103 4. Taking a Root - write the number in such a way that the exponent of 10 is divisible by 2, then take square root of the decimal part and divide the power of 10 by 2 - when a cube root is taken, write the number in such a way that the exponent of 10 is divisible by 3, then take cube root of the decimal part and divide the power of 10 by 3 5. Raising to a Power - when a number is squared, decimal part is squared and the exponent of 10 is multiplied by 2 - when a number is cubed, decimal part is cubed and the exponent of 10 is multiplied by 3.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 6 - in general, (a x 10n)p = ap x 10p(n) SIGNIFICANT FIGURES - every measurement is uncertain to some extent - exactness, or precision, of a measurement depends upon the limitations of the measuring device and the skill with which it is used  indicated by the number of figures used to record it - digits in a properly recorded measurement are significant figures  include all the figures that are known with certainty plus one more, which is an estimate Rules: 1. Zeros used only to locate the decimal point are not significant ex. 0.003 = 1 significant figure Zeros that arise as a part of a measurement are significant. ex. 0.0005030 = 4 SFs Sometimes, it is difficult to determine the number of significant figures in a value that contains zeros, like 700.  problem can be avoided by using scientific notation 700 = 7.00 x 102  3 SFs = 7.0 x 102  2 SFs = 7 x 102  1 SF Another convention, using the decimal point  if decimal point is indicated, ex. 700. all figures preceding the decimal point are significant. If decimal point is not used, then the zeros are not significant.  system not universally employed 2. Certain values, such as those that arise from the definition of terms, are exact. For example, by definition, 1 L = exactly 1000 mL  may be considered to have an infinite number of SFs (zeros) after the decimal point. Values obtained by counting may also be exact ex. H2 molecule contains exactly 2 atoms population of the world is not exact, only estimated 3. At times, the answer to a calculation contains more figures than are significant.  answer must then be rounded off to the correct number of digits.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 7 a. If the figure following the last number to be retained is less than 5, all the unwanted figures are discarded and the last number is left unchanged: 3.6247 is 3.62 to 3 SFs b. If the figure following the last number to be retained is greater than 5, or is 5 with other digits following it, the last figure is increased by 1 and the unwanted figures are discarded: 7.5647 is 7.565 to four SFs c. If the figure following the last number to be retained is 5 and there are no figures or only zeros following the 5, the 5 is discarded and the last figure increased by 1 if it is an odd number or left unchanged if it is an even number  last figure of the rounded-off value is always an even number  zero is considered to be an even number 3.250 is 3.2 to 2 SFs 7.635 is 7.64 to 3 SFs 8.105 is 8.10 to 3 SFs 4. The result of an addition or should be reported to the same number of decimal places as that of the term with the least number of decimal places. 161.032 5.6 32.4524  199.1  4 SFs 199.0844 5. The answer to a multiplication or division is rounded off to the same number of significant figures as is possessed by the least precise term used in the calculation. 152.06 x 0.24 = 36.4944  36  2 SFs (least precise tern is 0.24).......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 8 Report: Familiarization of Name: Activity Common Laboratory Section: #1 Apparatus Date: Instructor: Draw the following apparatus in the table below and give the uses of each: Test tube Test tube brush Stirring rod Erlenmeyer flask Tripod Glass funnel Beaker Iron ring Mortar and pestle Florence flask Iron stand Iron clamp Volumetric flask Thermometer Clamp holder Distilling flask Test tube holder Evaporating dish Graduated cylinder Nichrome wire Triple beam balance Alcohol lamp Tirril burner Water trough Pipet Aspirator bulb Dropper Test tube rack Watch glass Stainless/porcelain spatula Crucible with cover Crucible tong Hard glass test tube Drawing of Apparatus Name of Apparatus Use/s.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 9 Drawing of Apparatus Name of Apparatus Use/s.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 10 Drawing of Apparatus Name of Apparatus Use/s.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 11 Drawing of Apparatus Name of Apparatus Use/s Questions for research: 1. Which among the given apparatus are used for measuring accurate volumes of liquids? 2. Which among the given apparatus are used as reaction vessels? 3. Which of the given apparatus are used for high-temperature ignition? What material makes them highly resistant to breakage at very high temperatures?.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 12 Activity # 1b Measurements OBJECTIVES At the end of the activity, the students shall be able to: 1. perform the proper techniques in doing measurements 2. convert units of measurements MATERIALS/APPARATUS ruler, calculator, graduated cylinder, triple beam balance, blocks of wood, beaker PROCEDURE A. Measuring the Volume of Regular Solids 1. Using a ruler, measure the length, width and height of a regular object (wooden block, book, eraser, etc.) in millimeter. 2. Record measurements in the data sheet. Convert the measurements into centimeters. 3. Compute for the volume of the regular using the formula V=LxWxH in both mm3 and cm3. B. Measuring the Volume of cylinder 1. Measure the diameter of a cylinder in centimeter and compute for the radius by taking half of the diameter. 2. Measure height of the cylinder in centimeter 3. Compute for the volume of the cylinder using the formula V=1/2πr2h C. Measuring the Volume of Irregular solids (Water displacement method) 1. Fill up a 50 mL graduated cylinder with 25 mL of water. Record this as the initial volume, Vi. 2. Drop the piece of irregular solid (key, stone, bracelet, etc) into the cylinder with water. 3. Read the volume and record this as Vf. 4. Compute for the volume of the object, Vo, using the formula Vo=Vf-Vi. D. Measuring the Volume of liquids 1. Fill up a beaker with water up to approximately the 25 mL mark. 2. Transfer the water from the beaker into the graduated cylinder and record the actual reading from the graduated cylinder. 3. Repeat the same procedure using say sauce as sample........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 13 E. Measuring the Mass 1. Get a triple beam balance and examine the different parts and their uses. 2. Proper technique in using the triple beam balance will be demonstrated by your instructor. 3. Determine the masses (in grams) of the different objects that you used in procedures A to D. For liquids, use weighing by difference technique by subtracting the weight of the container from the weight of the liquid and container. 4. Convert the masses of the objects in g to kg and mg. F. Computing the Density of an Object 1. Using your data for volumes of the different objects of wooden block, water and irregular objects and the data for masses from procedure E, compute for the densities of the different objects using the formula, D=M/V. Units for density may vary between, g/cm3 or g/mL........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 14 Report: Names of members: Activity MEASUREMENTS Section: Date: # 1b Group: Instructor: DATA AND RESULTS: A. Volume of Regular Solid Regular solid Millimeter (mm) Centimeter (cm) Length Width Height Volume B. Volume of cylinder Cylinder Millimeter (mm) Centimeter (cm) Diameter Radius Height Volume C. Volume of irregular solid Initial volume of water (Vi) Final volume of water Volume of object (Vo) in in mL (Vf)in mL mL.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 15 D. Volume of liquids Liquid samples Volume of liquid using Volume of liquid using beaker (mL) graduated cylinder (mL) Water Soy sauce E. Mass of objects Samples g mg kg Wooden block Water Key F. Density Samples Mass (g) Volume (mL or Density cm3) Wooden block Water Key G. Questions for research 1. Which between the two unit systems used in the ruler is more convenient to use? Why? 2. Is there any discrepancy observed in the volume of water measured using beaker and the graduated cylinder? How does it affect the accuracy of your measurement? 3. In what unit should the volume of an irregular solid be expressed? Why?.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 16 4. Ice and liquid water have the same composition, However, their densities differ at standard atmospheric pressure and at 0oC. Ice has a density of 0.9167–0.9168 g/cm3, while water has a density of 0.9998–0.999863 g/cm3. How do you account for the icebergs on the sea or floating ice cubes on water? Explain briefly........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 17 Homework #1 Name: Date: Score: Course, Yr & Sec: Measurements in Chemistry 1. Determine the number of significant figures in each of the following measured values. (6 pts) a. 23,009 b. 0.00231 c. 0.3330 2. Round off each of the following numbers to the number of significant figures indicated in parentheses. (4 pts) a. 3883 (two) b. 0.0003011 (two) c. 4.4050 (three) d. 2.1000 (three) 3. Write out the names of the metric system units that have the following abbreviations. (4 pts) a. mg b. pg c. Mm d. dL 4. Express the following numbers in scientific notation. (6 pts) a. 37.06 b. 0.00571 c. 437.0 d. 4370 e. 0.20340 Problem solving: Show the solution and circle your final answer. 5. A typical loss of water through sweating per day for a human is 450 mL. What is the volume, in liters, of sweat produced per day? (2pts) 6. The smallest bone in the human body, which is in the ear, has a mass of 0.0030 g. What is the mass of this bone in pounds? (2 pts) 1 kg=2.2 lbs., 1kg = 103g.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 18 8. What volume of gasoline, in milliliters, would be required to fill a 17.0-gal gasoline tank? (2 pts) 1gallon=3.785L, 1mL= 10-3L 9. Air has a density of 1.29 g/L at room temperature. State whether each of the following will rise or sink in air: a.) Helium gas (density =0.18 g/L) b.) Argon gas (density = 1.78 g/L) (2 pts).......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 19 UNIT 2 PROPERTIES OF MATTER Matter Pure Mixture Substance Compoun Heterogen Homogen Elements ds ous ous non- inert Suspensio metals metalloids acids bases salts oxides Colloid Solutions metals gases n Matter is anything that occupies space and has mass. Matter can exist in three physical states: 1. gas or vapor 2. liquid 3. solid Gas This state of matter has no fixed volume or shape. It conforms to the volume and shape of its container. Gases can be compressed or expanded to occupy different volumes. Liquid A liquid has a distinct volume, independent of its container, but it has no specific shape. It assumes the shape of the container it is in. Liquids cannot be appreciably compressed. Solid A solid has a definite shape and volume; it is rigid. Solids cannot be appreciably compressed........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 20 Substances A pure substance has a fixed composition and distinct properties. Most matter we come in contact with in our daily lives is not a pure substance, but a mixture of substances. Physical and Chemical Properties Every pure substance has a unique set of properties - characteristics which allow us to distinguish it from other substances. These properties fall into two general categories: physical and chemical. Physical properties - properties we can measure without changing the basic identity of the substance. Chemical properties - describe the way a substance may change or "react" to form other substances. Physical and Chemical Changes Substances can undergo various changes in properties, these changes may be classified as either physical or chemical. Physical changes - a substance changes its physical appearance but not its basic identity. All changes of state (e.g. solid to liquid to gas) are physical changes. Chemical changes - also known as chemical reactions, a substance is transformed into a chemically different substance. Indices of chemical change in matter: 1. Formation of a new substance. 2. Release of heat and light energy. 3. Formation of precipitate of insoluble substance. 4. Changes in color, odor and taste 5. Release of gas Mixtures Mixtures refer to combinations of two or more substances in which each substance retains its own chemical identity and hence its own properties. Heterogenous mixtures are not uniform throughout the sample, and have regions of different appearance and properties Homogenous mixtures are uniform throughout the sample, however, the individual substances retain their individual chemical and physical nature. Homogenous mixtures.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 21 are also called solutions, however, the most common type of solution is described by a solid (the solute) dissolved in a liquid (the solvent). An important characteristic of mixtures is that the individual components retain their physical and chemical properties. Thus, it is possible to separate the components based on their different properties. For example, we can separate ethanol from water by making use of their different boiling temperatures, in a process known as distillation. Energy is the capacity to do work or transfer heat. Work is the energy transferred when a force exerted on an object causes a displacement of that object. Heat is the energy used to cause the temperature of an object to increase. Force is any push or pull on an object. Forms of energy 1. Kinetic energy is the energy of motion. Its magnitude depends on the object’s mass and its velocity: KE = ½mv2 2. Potential energy of an object depends on its relative position compared to other objects. Potential energy also refers to the composition of an object, including the energy stored in chemical bonds. One of the goals in chemistry is to related the energy changes in the macroscopic world to the kinetic or potential energy of substances at the molecular level........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 22 Activity # 2 Matter: Elements and compounds, Physical and Chemical Changes OBJECTIVES At the end of the activity, the students shall be able to: 1. characterize the physical changes in matter 2. differentiate physical from chemical changes in matter 3. identify the indices of chemical change MATERIALS/APPARATUS Test tubes, alcohol lamp, test tube rack, evaporating dish, beaker, graduated cylinder, nichrome wire, mortar and pestle, crucible tong PROCEDURE A. Physical properties of matter 1. Complete the table below and take note of the color, state, classification and solubility of each sample in water Sample Color State Classification Solubility in water Salt Sugar Charcoal Copper wire Monosodium glutamate Calcium carbonate Soy sauce Coconut oil Vinegar B. Chemical and Physical changes 1. Using a crucible tong, burn a piece of copper wire using an alcohol lamp until red hot. Observe what happened to the copper wire. 2. Place a piece of copper wire on a watch glass. Add about 10 drops of silver nitrate. Stand for 15 to 20 minutes. Observe changes that happened to the copper wire. 3. Burn a matchstick. Take note of the changes that happened to the matchstick. 4. Place pieces of ice cubes into a beaker with water. Stand for 5 minutes. What was formed on the sides of the beaker. C. Separation techniques for Mixtures 1. Grinding- Grind half teaspoon each of salt and benzoic acid using a mortar and pestle. Observe the mixture. List down if it is a homogenous of heterogenous mixture........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 23 2. Hydration- Transfer the mixture from C.1 into a beaker containing 15mL of water and stir. Observe the type of mixture that results from the reaction. 3. Filtration- Line a glass funnel with filter paper. Pour the mixture from C.2 and filter. Collect the filtrate and residue. Note your observations. 4. Evaporation- Transfer the filtrate that was collected from C.3 into an evaporating dish. Evaporate the liquid to dryness. Observe what was formed in the evaporating dish after heating........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 24 Report: Names of members: Matter: Elements and Activity compounds, Physical #2 and Chemical Changes Section: Date: Group: Instructor: DATA AND RESULTS: A. Properties of Matter Sample Color State Classification Solubility in water Salt Sugar Charcoal Copper wire Monosodium glutamate Calcium carbonate Cupric sulfate Sulfur Vinegar B. Physical and Chemical Changes in Matter Reaction Observations Type of Change 1. Heating of copper wire 2. Reaction of copper wire with silver nitrate 3. Burning of matchstick 4. Melting of ice.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 25 C. Separation techniques for Mixtures Separation techniques and Observation Type of mixture their definitions 1. Grinding 2. Hydration 3. Filtration 4. Evaporation Questions for Research 1. How do physical change differ from chemical change? 2. What are the indices of chemical changes in matter? 3. In the melting of ice, what was formed on the sides of the container? Where did it come from? 4. How do you classify rusting of iron in terms of changes in matter? Explain in detail the processes that take place. What effect does this process impose upon the economy as a whole. How do you prevent such process to occur in iron........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 26 UNIT 3 ATOMIC STRUCTURE AND THE PERIODIC TABLE SUBATOMIC PARTICLES 1. Electron - Sir William Crookes (1832-1919) passed an electric current through a gas-filled tube (Crookes tube) with electrodes sealed at both ends - as gas was passed out of a tube, a pressure was reached at which the remaining gas glowed  glow produced by negative particles, called cathode rays, passing from the negative electrode (cathode) to the negative electrode (anode)  streams of negatively charged particles are called electrons, which travel in straight lines - 1897, English physicist Sir J.J. Thomson determined the ratio of the charge, e, to the mass, m, of cathode-ray particles  e/m found to be identical in all three subatomic particles (protons, electrons, neurons)  e was measured by American physicist R.A. Millikan in 1909 e = 1.6022 x 10-19 Coulomb e/m = charge to mass ratio for electron = 1.75881962 x 1011 Coulomb/g m = e_ = 1.6022 x 10-19 Coulomb (e/m) 1.75881962 x 1011 Coulomb/g  mass of electron (me-) = 9.1096 x 10-28 g 2. Proton - 1886, German physicist Eugen Goldstein, using Crookes tube, observed that a fluorescence (in the inner surface of a cathode) was emitted from the anode and passed through the holes in the cathode  indicates movement of positive rays, called protons  caused by the loss of electrons from the gas molecules in the tube when they are struck by high-speed cathode rays e/m = 9.5791 x 104 C/g e = 1.6022 x 10-19 C mp+ = 1.6726 x 10-24 g 3. Neutron - 1932, English physicist James Chadwick - neutrally (zero) charged particle whose mass is close to that of a proton.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 27 Particle Mass (g) Mass (amu) Relative Charge electron (e-) 9.1096 x 10-28 0.00055 -1 proton (p+) 1.6726 x 10-24 1.00728 +1 neutron (n) 1.6749 x 10-24 1.00867 0 1 amu (atomic mass unit) = 1.6606 x 10-24 g Facts: 1. Protons and neutrons are found at the center of the atom called the nucleus. Electrons are found outside the nucleus in shells or energy levels. 2. Number of protons in the nucleus is defined as the atomic number 3. Atomic weight is the sum of the number of protons and neutrons in the nucleus 4. Number of electrons in the electron cloud of an atom = Number of protons in the nucleus  atom is neutral A A = atomic weight/mass number (in amu) E E = symbol of the element Z Z = atomic number Isotopes - atoms of a given element which have the same number of protons and electrons (i.e. same atomic number, Z) but different atomic masses (A  diff. no. of neutrons) - e.g. seven isotopes of C, but only 12C and 13C are stable ELECTRONIC STRUCTURE OF THE ATOM - current chemical theory indicates that as electrons move about an atom’s nucleus, they are restricted to specific regions around the nucleus of the atom - such restrictions are determined by the amount of energy the electrons possess - electron energies are limited to certain values and that a specific “behavior” is associated with each allowed energy value - electrons are arranged in shells or orbits according to theories of quantum mechanics.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 28 Quantum Mechanical Model of the Atom - based on the mathematical expression called the wave function 1. location of an electron cannot be determined exactly; all that can be identified is the region of space where there is a relatively high probability of finding the electron  the orbital occupied by the electron 2. orbitals are characterized by the principal quantum number, n, an integer (n = 1, 2, 3…) (shells)  as n increases, the orbitals extend farther from the nucleus, energy of the orbitals also increases n: 1 2 3 4 5 6 7 shell: K L M N O P Q 3. orbitals with the same value of n may have different shapes  different shapes of orbitals are distinguished by the second quantum number, l, the azimuthal or subsidiary quantum number (subshells) l = 0 to n-1 ex: n=1 n=3 l=n–1=0 n – 1 = 2  l = 0, 1, 2 for l = 0  s-orbital (draw!) l = 1  p-orbital l = 2  d-orbital l = 3  f-orbital Shell Principal Quantum Azimuthal Quantum Orbitals No. (n) No. (l) K 1 0 1s L 2 0 2s 1 2p M 3 0 3s 1 3p 2 3d N 4 0 4s 1 4p 2 4d 3 4f.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 29 4. third quantum number, m, magnetic quantum number (orbitals)  may have integral value from –l to + l Principal Quantum Azimuthal Quantum Magnetic Quantum No. No. (n) No. (l) (m) 1 0 0 2 0 0 1 -1, 0, +1 3 0 0 1 -1, 0, +1 2 -2, -1, 0, +1, +2 4 0 0 1 -1, 0, +1 2 -2, -1, 0, +1, +2 3 -3, -2, -1, 0, +1, +2 5. fourth quantum number, s, spin quantum number  specifies the direction of spin of an electron  either clockwise (s = + ½) or counterclockwise (S = - ½) 6. maximum number of electrons that may be found in a shell = 2n2; number of atomic orbitals (s, p, d, f) in a shell = n2 ; each atomic orbital can hold up to 2 electrons Summary of Allowed Values for Each Quantum Number Shell n l orbital m s K 1 0 1s 0 ±½ L 2 0 2s 0 ±½ 1 2p -1 ±½ 1 2p 0 ±½ 1 2p +1 ±½ M 3 0 3s 0 ±½ 1 3p -1 ±½ 1 3p 0 ±½ 1 3p +1 ±½ 2 3d -2 ±½ 2 3d -1 ±½ 2 3d 0 ±½ 2 3d +1 ±½ 2 3d +2 ±½.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 30 The Electronic Structure of Atoms The arrangement of electrons of an atom that are arranged outside the nucleus of an atom determines the chemical properties of an atom. The General notation below shows the mass number (A) a.k.a. atomic weight or atomic mass, represents the number of protons and neutrons. The atomic number (Z) represents the number of protons. Since the atoms is electrical neutral in the free state, this number also represents the number of electrons. The number of neutrons is taken by subtracting the atomic umber from the mass number. For oxygen, the general notation is represented as 16 O8 the number of protons = 8 the number of electrons = 8 the number of neutrons =8 The number of electrons is equal to 8 and is arranged and distributed using the Aufbau diagram. The Periodic Law and the Periodic Table Periodic Law – When elements are arranged in order of increasing atomic number, elements with similar chemical properties occur at periodic (regularly recurring) intervals. Periodic Table – Tabular arrangement of the elements in order of increasing atomic number such that elements having similar chemical properties are positioned in vertical columns........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 31 Dmitri Ivanovich Mendeleev: The Father of the Periodic Table. He arranged the elements according to increasing atomic numbers. The Periodic Table Periods – horizontal rows of elements Groups – elements in the same vertical columns; have similar chemical properties The Periodic Table of Elements Electron Configurations A statement of how many electrons an atom has in each of its electron subshells. An oxygen atom as an electron arrangement of two electrons in the 1s subshell, two electrons in the 2s subshell, and four electrons in the 2p subshell. Oxygen: 1s22s22p4.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 32 The Aufbau Diagram a.k.a. the Building Up Principle Shell diagram This shows the arrangement of electrons that are distributed to each of the shells found outside the nucleus of the atom. For oxygen with electronic configuration of 1s22s22p4, the shell diagram is represented as; O 2) 6) Orbital Diagrams A notation that shows how many electrons an atom has in each of its occupied electron orbitals. Oxygen: 1s22s22p4 Oxygen: 1s 2s 2p Distinguishing Electron Last electron added to the electron configuration for an element when electron subshells are filled in order of increasing energy. This last electron is the one that causes an element’s electron configuration to differ from that of an element immediately preceding it in the periodic table. The Electronic Basis for the Periodic Law and the Periodic Table The electron arrangement in the outermost shell is the same for elements in the same group. This is why elements in the same group have similar chemical properties.  Group 1A – very reactive.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 33 Li: 1s22s1 Na: 1s22s22p63s1 K: 1s22s22p63s23p64s1 Classification of Elements 1. A system based on selected physical properties of the elements, in which they are described as metals or nonmetals. 2. A system based on the electron configurations of the elements, in which elements are described as noble-gas, representative, transition, or inner transition elements........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 34 The properties of metals and non-metals Identifying the properties and classification of elements using their electronic configuration 1. Distribute the electrons of an element using the Aufbau diagram. 16 O8 the number of protons = 8 the number of electrons = 8 the number of neutrons =8.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 35 2. The distinguishing electron a.k.a. differentiating electron is the last electronic configuration that enters an orbital. If the distinguishing electron fills up the s or p blocks, then it belong to Group A (Representative elements), while if it fills the d or f blocks, then it belongs to Group B. Oxygen: 1s22s22p4 Oxygen fills the p block, so it belongs to Group A. 3. To identify the Family of the element, distribute the electrons using a shell diagram. If the second to the last shell shows stable configuration (2, 8 or 18) , then the number of electrons in the valence shell represents the number of the family. For oxygen with electronic configuration of 1s22s22p4, the shell diagram is represented as; O 2) 6) Family VIA 4. To identify the period/series where the element belongs, count the number of shells which corresponds to the period/series of the element. O 2) 6) Period/Series: 2nd period 5. To classify the elements, the valence electrons (the number of electrons in the outermost shell) represents the classification of the element. If the valence electron = 1,2 or 3, the element is a metal If the valence electron = 5,6 or 7, the element is a non metal If the valence electron = 4, the element is a metalloid If the valence electron = 8, the element is an inert gas/noble gas O 2) 6) With valence electron of 6, oxygen is classified as a non-metal.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 36 Activity # 3 Atomic Structure: Flame Test OBJECTIVES At the end of the activity, the students shall be able to: 1. identify metals by observing their visible spectrum 2. explain the color of flame emitted by metal in terms of electronic transition of atoms. MATERIALS/APPARATUS test tubes, test tube rack, alcohol lamp, nichrome wire EXPERIMENTAL PROCEDURE A. Flame Test 1. Place about 1 mL of the different metal chlorides in different test tubes. Label properly. 2. Get a nichrome wire and sterilize by dipping it with hydrochloric acid and heating it with the alcohol lamp until it is red hot. 3. Take a sample of the sample metal chlorides by dipping the tip of the nichrome wire to a test tube and subject to the flame. Take note of the color of the flame 4. Sterilize the nichrome wire by dipping it in the HCl and flaming until red hot. 5. Let the nichrome wire cool down and dip the nichrome wire in another test tube and take note of the color of the flame. 6. Repeat steps 4-5 for the rest of the samples. Reagents Potassium chloride Barium chloride Copper II chloride Sodium chloride Strontium chloride Lithium chloride Calcium chloride Refer to the following youtube link for the experiment: MegaLab Flame Test: https://www.youtube.com/watch?v=NEUbBAGw14k.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 37 Report: Names: Activity Atomic Structure: Section: Date: #3 Flame Test Group: Instructor: DATA AND RESULTS: A. Flame test Reagents Metal tested Color of Flame Potassium chloride Barium chloride Copper II chloride Sodium chloride Strontium chloride Lithium chloride Calcium chloride B. Questions for research 1. Explain the principle behind the ability of metals to emit different colors of flame. 2. Give examples of the practical applications of Flame test. 3. What is Atomic Absorption Spectroscopy (AAS)? How is the principle of AAS related to flame test? 4. Give examples of the practical uses of AAS........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 38 UNIT 4 MOLECULES, CHEMICAL BONDS AND CHEMICAL EQUATIONS Why do atoms combine? Atoms combine to become stable and follow the inert gas configuration. Some substances are chemically bonded molecules and others are an association of ions. This depends upon the electronic structures of the atoms and the nature of the chemical forces within the compounds. Classification of chemical forces/bonds 1. Ionic bonds 2. Covalent bonds 3. Metallic bonds Ionic bonds - electrostatic forces that exist between ions of opposite charge. This type of bond typically involves a reaction between metal with a nonmeat. l Covalent bonds - results from the sharing of electrons between two atoms. This, on the other hand typically involves one non-metallic element with another. Examples: Hydrogen with another hydrogen Hydrogen with another non-metal Non-metal with another non-metal C/Si with another non-metal Metallic bonds are found in solid metals (copper, iron, aluminum) and each metal bonded to several neighboring groups of metals. The bonding electrons free to move throughout the 3-dimensional structure. Lewis Symbols and the Octet Rule Valence electrons reside in the outer shell and are the electrons which are involved in chemical interactions and bonding (valence comes from the Latin valere, "to be strong"). Electron-dot notation (Lewis symbols)  consists of the chemical symbol for the element plus a dot for each valence electron.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 39 Note in writing Lewis dot symbol:  The dot (representing electrons) are placed on the four sides of the atomic symbol (top, bottom, left, right)  Each side can accommodate up to 2 electrons (Pauli exclusion principle)  The number of valence electrons in the table below is the same as the column number of the element in the periodic table (for representative elements only) IONIC BONDING Sodium with atomic number of 11 Electron configuration is 1s22s22p63s1, thus there is 1 valence electron. Its shell diagram is represented by: Na 2) 8) 1) With 1 valence electron, its Lewis symbol would therefore be: Na Chlorine with atomic number of 17 Electron configuration is 1s22s22p63s23p5 thus there are 7 valence electrons. Its shell diagram is represented by: Cl 2) 8) 7) With 7 valence electrons, its Lewis symbol would therefore be: Atoms often gain, lose, or share electrons to achieve the same number of electrons as the inert gas closest to them in the periodic table Because all noble gasses (except He) have filled s and p valence orbitals (8 electrons), many atoms undergoing reactions also end up with 8 valence electrons. This observation has led to the Octet Rule: Atoms tend to lose, gain, or share electrons until they are surrounded by 8 valence electrons Sodium tends to lose its lone valence electron, while chlorine tends to gain 1 more electron to follow the octet rule........................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 40 The compound NaCl contains an ionic bond between Na and Cl. Na gets a +1 charge while Cl gets a -1 charge due to the complete transfer of electrons from sodium to chlorine forming NaCl which is electrically neutral when in compound crystal form. Note: the first four elements in the Periotic Table follows the Duet rule (Rule of 2) that made them exceptions to the octet rule (H, Li, He and Be). COVALENT BONDING Hydrogen with atomic number of 1 Electron configuration is 1s1 thus there is 1 valence electron. Its shell diagram is represented by: H 1) One H atom reacting with another H atom is represented by: H.. H  H-H or H2 The bond that exists between the diatomic hydrogen is a covalent bond which is mutually shared between the two hydrogen atoms. One hydrogen atom reacting with another non-metal like Cl, is represented by: H.  H-Cl The bond that exists between the hydrogen and chlorine is a covalent bond which is mutually shared between chlorine and hydrogen atom. Bond Polarity and Electronegativity The electron pairs shared between two atoms are not necessarily shared equally Extreme examples: 1. In Cl2 the shared electron pairs is shared equally.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 41 2. In NaCl the 3s electron is stripped from the Na atom and is incorporated into the electronic structure of the Cl atom - and the compound is most accurately described as consisting of individual Na+ and Cl- ions For most covalent substances, their bond character falls between these two extremes Bond polarity is a useful concept for describing the sharing of electrons between atoms  A nonpolar covalent bond is one in which the electrons are shared equally between two atoms  A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. If this relative attraction is great enough, then the bond is an ionic bond Electronegativity A quantity termed 'electronegativity' is used to determine whether a given bond will be nonpolar covalent, polar covalent, or ionic. Electronegativity is defined as the ability of an atom in a particular molecule to attract electrons to itself (the greater the value, the greater the attractiveness for electrons) Electronegativity is a function of:  the atom's ionization energy (how strongly the atom holds on to its own electrons)  the atom's electron affinity (how strongly the atom attracts other electrons) Fluorine is the most electronegative element (electronegativity = 4.0), the least electronegative is Cesium (notice that are at diagonal corners of the periodic chart).......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 42 General trends:  Electronegativity increases from left to right along a period  For the representative elements (s and p block) the electronegativity decreases as you go down a group  The transition metal group is not as predictable as far as electronegativity Electronegativity and bond polarity The difference in electronegativity between two atoms can used to gauge the polarity of the bonding between them Compound F2 HF LiF Electronegativity 4.0 - 4.0 = 0 4.0 - 2.1 = 1.9 4.0 - 1.0 = 3.0 Difference Type of Bond Nonpolar covalent Polar covalent Ionic (non-covalent)  In F2 the electrons are shared equally between the atoms, the bond is nonpolar covalent  In HF the fluorine atom has greater electronegativity than the hydrogen atom. The H-F bond can thus be represented as:.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 43  The 'd+' and 'd-' symbols indicate partial positive and negative charges.  The arrow indicates the "pull" of electrons off the hydrogen and towards the more electronegative atom  In lithium fluoride the much greater relative electronegativity of the fluorine atom completely strips the electron from the lithium and the result is an ionic bond (no sharing of the electron) A general rule of thumb for predicting the type of bond based upon electronegativity differences:  If the electronegativities are equal (i.e. if the electronegativity difference is 0), the bond is non-polar covalent  If the difference in electronegativities between the two atoms is greater than 0, but less than 2.0, the bond is polar covalent  If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic CHEMICAL NOMENCLATURE - system of names used to distinguish compounds from each other and the rules needed to devise these names - old system, rule was “anything goes” o ex. quicksilver (mercury), gypsum (calcium sulfate), laughing gas (nitrous oxide) - International Union of Pure and Applied Chemistry (IUPAC) Rules o set of compound-naming rules produced by committees of the IUPAC Types of Compounds based on the number of elements present: 1. binary compound o contains just two different elements ex. NH3, H2O, CO2 o any number of atoms of the two elements may be present in a molecule or formula unit, but only two elements may be present 2. ternary compound o contains three different elements ex. HNO3, H2SO4, NaOH Ionic and Molecular Compounds (for nomenclature purposes): 1. metal + nonmetal  ionic 2. combination of nonmetals  covalent  molecular 3. polyatomic ions  ionic  whole polyatomic ion reacts like a monoatomic ion 4. metalloids are considered to be nonmetals  metalloid + nonmetal is molecular.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 44 BINARY IONIC COMPOUNDS - simplest type of ionic compound  only monoatomic ions are present - monoatomic positively-charged metallic ions + monoatomic negatively-charged nonmetallic ions a. binary ionic compounds containing a fixed-charge metal b. binary ionic compounds containing a variable-charge metal Fixed-Charge Metal: - always exhibit the same behavior in ion formation  they always lose the same number of e-s - forms only one type of ion, which always has the same magnitude - group IA metals  +1 ions - group IIA metals  +2 ions - Al, Ga from group IIIA  +3 ions - Zn (group IIB, +2), Cd (group IIB, +2) and Ag (group IB, +1)  reason for the charges is periodic-table correlation  the octet rule Variable-Charge Metal: - forms more than one type of ion, with different charges - all metals, except for the 15 fixed-charged metals, are variable-charge  cannot be easily related to periodic-table position  presence of d or f electrons complicates octet rule - see Table 1 Nomenclature for Binary Ionic Compounds: 1. compounds containing a fixed-charge metal  full name of the metallic element is given first, followed by a separate word consisting of the root of the nonmetallic element name and the suffix –ide (see Table 2) ex. NaF  sodium fluoride 2. compounds containing a variable-charge metal  charge of the metal must be incorporated into the name of the compound  magnitude of the charge is indicated by using a Roman numeral, inside parentheses, placed immediately after the name of the metal ex. Fe+2: iron (II) ion Fe+3: iron (III) ion Au +: gold (I) ion  Roman numeral is considered to be part of the metal’s name, but not of the formula ex: Name the following binary ionic compounds:.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 45 FeO  get charge of the Fe first by calculating from the oxidation number of O  net charge must be equal to 0. - other method for indicating the charge on metal ions uses suffixes rather than Roman numerals  more complicated and less precise, currently being abandoned  when a metal has two common ionic charges: o suffix –ous  used for the ion of lower charge o suffix –ic  used for the ion of higher charge (see Table 3) IONIC COMPOUNDS CONTAINING POLYATOMIC IONS (see Table 4) - names are derived in the same way as those of binary ionic compounds  metallic ion first, then nonmetallic ion ending in -ide  if polyatomic ion is positive, its name is substituted for that of the metal  if polyatomic ion is negative, its name is substituted for the nonmetal stem plus –ide  if both positive and negative ions are polyatomic, name includes just the names of the polyatomic ions ex. K2CO3  potassium carbonate (H3O)2S  hydronium sulfide Co(NO3)3  cobalt (III) nitrate BINARY MOLECULAR COMPOUNDS - covalently bonded compound in which just two nonmetallic elements are present 1. named in the order in which they appear in the formula  least electronegative nonmetal is usually written first in the formula  name of the first nonmetal is used in full, name of second nonmetal treated like in binary ionic compounds  stem plus –ide 2. number of atoms of each element present in a molecule is explicitly incorporated into the name of the compound by using Greek numerical prefixes (see Table 5)  prefix precedes name of each nonmetal note: in ionic compounds, formula subscripts are not mentioned in the name  prefixes are needed in naming binary molecular compounds because numerous different compounds exist for many pairs of nonmetallic elements ex. NO, NO2, N2O, N2O3, N2O4, N2O5 nitrogen monoxide, nitrogen dioxide, dinitrogen oxide, dinitrogen trioxide, dinitrogen tetraoxide, dinitrogen pentaoxide.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 46 NOMENCLATURE FOR ACIDS Acids - substances that produce H+ in aqueous solutions 1. non-oxyacids  acids composed of hydrogen and one or more nonmetals other than oxygen  all common non-oxyacids, except HCN, are binary compounds a. the word hydrogen is replaced by the prefix hydro- b. suffix –ide on the stem of the name of the nonmetal is replaced with the suffix –ic c. the word acid is added to the end of the name (as a separate word) ex. HCl(sol’n)  hydrochloric acid H2Se(sol’n)  hydroselenic acid HCN(sol’n)  hydrocyanic acid 2. oxyacids  acids composed of hydrogen , oxygen, and another nonmetal  names are derived from the names of the polyatomic ions produced when the acid molecules break into ions in solution ex. H2SO4(sol’n)  H+ + SO4-2 a. when polyatomic ion produced ends in –ate, –ate is replaced with –ic, then acid b. when polyatomic ion produced ends in –ite, –ite is replaced with –ous, plus acid ex. H2SO4(sol’n)  sulfuric acid HNO3(sol’n)  nitric acid HNO2(sol’n)  nitrous acid TABLE 1. Common Variable-Charge Metallic Element Ions and Their Charges Element Symbol Ions Formed chromium Cr Cr+2, Cr+3 cobalt Co Co+2, Co+3 copper Cu Cu+, Cu+2 gold Au Au+, Au+3 iron Fe Fe+2, Fe+3 lead Pb Pb+2, Pb+4 manganese Mn Mn+2, Mn+3 tin Sn Sn+2, Sn+4.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 47 TABLE 2. Names of Some Common Nonmetal Ions Element Stem/Root Name of Ion Formula bromine brom- bromide ion Br- carbon carb- carbide ion C-4 chlorine chlor- chloride ion Cl- fluorine fluor- fluoride ion F- hydrogen hydr- hydride ion H- iodine iod- iodide ion I- nitrogen nitr- nitride ion N-3 oxygen ox- oxide ion O-2 phosphorus phosph- phosphide ion P-3 sulfur sulf- sulfide ion S-2 TABLE 3. Comparison of IUPAC and Old System Names for Selected Metal Ions Element Ions Preferred Name Old System copper Cu+ copper (I) ion cuprous ion Cu+2 copper (II) ion cupric ion iron Fe+2 iron (II) ion ferrous ion Fe+3 iron (III) ion ferric ion tin Sn+2 tin (II) ion stannous ion Sn+4 tin (IV) ion stannic ion lead Pb+2 lead (II) ion plumbous ion Pb+4 lead (IV) ion plumbic ion gold Au+ gold (I) ion aurous ion Au+3 gold (III) ion auric ion TABLE 4. Formulas and Names of Some Common Polyatomic Ions Key Element Present Formula Name of Ion nitrogen NO3- nitrate ion NO2- nitrite ion NH4+ ammonium ion N-3 azide ion sulfur SO4-2 sulfate ion HSO4- hydrogen sulfate/bisulfate ion SO3-2 sulfite ion HSO3- hydrogen sulfite/bisulfite ion S2O3-2 thiosulfate ion phosphorus PO4-3 phosphate ion HPO4-2 hydrogen phosphate/biphosphate ion H2PO4- dihydrogen phosphate ion.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 48 PO3-3 phosphate ion carbon CO3-2 carbonate ion HCO3- hydrogen carbonate/bicarbonate ion C2O4-2 oxalate ion CH3COO- acetate ion /C2H3O2- CN- cyanide ion OCN- cyanate ion SCN- thiocyanate ion chlorine ClO4- perchlorate ion ClO3- chlorate ion ClO2- chlorite ion ClO- hypochlorite ion oxygen O2-2 peroxide ion boron BO3-3 borate ion hydrogen H 3O + hydronium ion OH- hydroxide ion metals MnO4 - permanganate ion CrO4-2 chromate ion Cr2O7 -2 dichromate ion * most frequently encountered polyatomic ions are in italics Facts Concerning the Polyatomic Ions in Table 4: 1. most of the ions are negatively-charged (only H3O+ and NH4+ are positive) 2. four of the polyatomic ions have names ending in –ide: hydroxide, cyanide, azide, peroxide 3. –ate and –ite pairs of ions  ion in the pair w/ the higher number of oxygen is always the –ate ion  -ite ion always contains one less oxygen than the –ate ion 4. pairs of ions where one member of the pair differs from the other by having a hydrogen atom present ex. CO3-2 (carbonate) and HCO3- (hydrogen carbonate or bicarbonate)  charge on the hydrogen-containing ion is always one less than the charge on the other ion 4. SO4-2 and S2O3-2, OCN- and SCN-  a sulfur atom has replaced an oxygen atom in one member of the pair  prefix –thio is used to denote this replacement  sulfate-thiosulfate, cyanate-thiocyanate.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 49 TABLE 5. Greek Numerical Prefixes (from 1 to 10) Greek Prefix Number mono- 1 di- 2 tri- 3 tetra- 4 penta- 5 hexa- 6 hepta- 7 octa- 8 nona- 9 deca- 10.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 50 CHEMICAL EQUATIONS Chemical Reaction - process in which at least one new substance is produced as a result of a chemical change  color change, emission of heat and/or light, evolution of gas, precipitation, etc. *reactants - starting materials for a chemical reaction - consumed/used up as chemical reaction proceeds *products - substances produced as a result of a chemical reaction Law of Conservation of Mass - The law states that “Mass is neither created nor destroyed in any transformation of matter” - sum of the masses of the products is always the same as the sum of the masses of the reactants - Antoine Laurent Lavoisier (1743-1794) - in a chemical reaction, total mass of reactants = total mass of products Writing Chemical Equations Chemical Equation - written statement that uses symbols and formulas instead of words to describe the changes that occur in a chemical reaction ex. magnesium oxide reacts with carbon to produce carbon monoxide and magnesium chemical equation: MgO + C  CO + Mg Conventions: 1. the correct formulas of the reactants are always written on the left side of the equation 2. the correct formulas of the products are always written on the right side of the equation 3. the reactants and products are separated by an arrow pointing toward the products 4. plus signs are used to separate different reactants or products from each other 5. if the physical states of the substances involved are of interest, they.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 51 may be indicated in parentheses after each formula  the common states encountered are i. (g) for gas ii. (l) for liquid iii. (s) for solid iv. (aq) for aqueous solution Note: “+” signs on the reactant side of the equation means “reacts with” “” means “to produce” “+” signs on the product side means “and” - in writing chemical equations, two conditions must be followed: o it must be consistent with experimental facts  only reactants and products that are actually involved in the reaction are included  correct formulas for the reactants and products are used o it must be consistent with the law of conservation of mass  chemical equation must be balanced Balancing Chemical Equations - there must be the same number of atoms of each element involved in the reaction on each side of the equation - use of coefficients o number placed before the formula of a substance to denote the amount of that substance ex. 3Cu(s) + 8HNO3(aq)  3Cu(NO3)2(l) + 2NO(g) + 4H2O(l) PCl3 + 3H2O  H3PO3 + 3HCl NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq) Types of Chemical Reactions 1. Synthesis reaction/Direct union - single product is produced from two (or more) reactants X + Y  XY - XY is always a compound ex. H2 + Cl2  2HCl S + O2  SO2 SO3 + H2O  H2SO4 2NO + O2  2NO2.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 52 2. Decomposition reaction/Analysis - single reactant is converted (broken down or decomposed) into two or more simpler substances - opposite of the synthesis reaction XY  X + Y ex. 2CuO  2Cu + O2 2H2O 2H2 + O2 CaCO3  CaO + CO2 3. Single-replacement reaction - one element within a compound is replaced by another element - there are always two reactants (one element, one compound), and two products (also an element and a compound) X + YZ  Y + XZ ex. Zn + H2SO4  H2 + ZnSO4 Ni + 2HCl  H2 + NiCl2 Mg + Ni(NO3)2  Ni + Mg(NO3)2 4. Double-replacement reaction - two compounds exchange parts with each other and form two different compounds AX + BY  AY + BX - generally involve ionic compounds in aqueous solution - “partner swapping” ex: AgNO3 + NaCl  NaNO3 + AgCl NaF + HCl  NaCl + HF AgNO3 + HCl  AgCl + HNO3 1. double displacement reactions are also called “metathesis” reactions 2. occur when a precipitate (insoluble solid), an insoluble gas, or a weak electrolyte is formed ex: AgNO3 + HCl  AgCl + HNO3  AgCl is an insoluble salt 2HCl + Na2S  H2S(g) + 2NaCl  H2S(g) is an insoluble gas HCl + NaOH  H2O + NaCl  H2O is a weak electrolyte Solubility Rules (see Table 6) 3. rules in table 6 apply to compounds of the following cations: a. +1 cations: Li+, Na+, K+, Rb+, Cs+, NH4+, Ag+.......................................................... BY: RUTH T. LIBAG, RMT, LPT, MAE CHEMISTRY 53 b. +2 cations: Mg+2, Ca+2, Sr+2, Ba+2, Mn+2, Fe+2, Co+2, Ni+2, Cu+2, Zn+2, Cd+2, Hg+2, Hg2+2, Sn+2, Pb+2 c. +3 cations: Fe+3, Al+3, Cr+3 4. common inorganic solids (HCl, HBr, etc.) are soluble in water Table 6. Mainly Water-Soluble Ionic Compounds NO3- all nitrates are soluble CH3COO- all acetates are soluble ClO3- all chlorates are soluble Cl- all chlorides are soluble except AgCl, Hg2Cl2, and PbCl2* Br- all bromides are soluble except AgBr, Hg2Br2, PbBr2* and HgBr2* I- all iodides are soluble except AgI, Hg2I2, PbI2, and HgI2 SO4-2 all sulfates are soluble except CaSO4*, SrSO4, BaSO4, PbSO4, Hg2SO4 and Ag2SO4* Mainly Water-Insoluble Ionic Compounds S-2 all sulfides are insoluble except those of the IA and IIA elements and (NH4)2S CO3-2 all carbonates are insoluble except those of the IA elements and (NH4)2CO3 SO3 -2

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