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CHAPTER 1- CHEMISTRY OVERVIEW Chapter 1 introduces the concepts of science, scientific method, the concept of matter, energy and the physical means by which they are measured. INSTRUCTIONS, ASSIGNMENTS & QUIZ (for online class)

• Login to https://laulima.hawaii.edu • Complete Assignment #1 Scavenger Hunt Activity (mandatory, found online) • Study chapters 1 and 2 in Lecture Notes. Do self-assessment & learning checks. • Study chapters 1 and 2 multimedia in Laulima Modules • Do assignment#2 and upload in Laulima Assignments. • Take practice quiz #1 (unlimited times) to score 80% or better in Laulima TTS • Take actual quiz #1 in Laulima Tasks, Tests and Surveys. • Do extra credit and upload in Laulima Assignments.

OBJECTIVES After completing this chapter, you should be able to: 1. Describe the scientific method, and differentiate between observation, hypothesis, experiment, theory and law. 2. Explain why we should study chemistry. 3. Explain what is desirability quotient and apply risk-benefit analysis. 4. Explain what matter is, and the three states of matter 5. Explain the differences between physical and chemical properties and changes. 6. Classify matter into the correct category based on the following pairs: a) heterogeneous or homogenous, b) solution or pure substance, c) element or compound. 7. Understand the importance of measurement units and the units of the metric system. 8. Be able to convert measurements, including magnitude prefixes. 9. Solve problems on temperature and density. ORGANIZATION Chapter 1 – Chemistry is divided into four sections: A. Chemistry: A Science for All Seasons B. The Scientific Method

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C. Matter & Its Classification D. Measurement of Matter SELF-STUDY/REVIEW QUESTIONS The concepts outlined in the list below do not comprehensively cover every question you may expect to see on the quiz or final exam. Use the following questions as a self-assessment tool to help you gauge your understanding of the course material in this Module. 1.Know the steps in scientific method. Be able to distinguish between hypothesis,

theory, scientific law, experiments and control. 2.Be able to estimate desirability quotient using risk benefit analysis. 3.Classify whether a specific matter is pure or a mixture, whether it is a(n) element,

compound or mixture (homogeneous vs. heterogeneous). Note that the term “element” has two usages: a) pure substance b) a kind of atom

4.Classify a specific change as physical or chemical change. 5.Differentiate gas, liquid and solid phase models. 6.Use appropriate units when reporting measurements and convert among units of mass,

length, volume, and temperature (e.g. kg to g, mL to L, 0C to K, etc) 7.Know the meaning of density and how it is calculated. 8.Know the meaning of these terms (see Appendix) KEY  TERMS  applied  research  basic  research  Celsius  scale  (°C)  chemical  change  chemical  properties  chemical  symbol  chemistry  compound  density  element  gas  hypothesis  kelvin  (K)  kilogram  (kg)  liquid  liter  (L)  

mass  matter  meter  (m)  mixture  physical  change  physical  properties  risk–benefit  analysis  science  scientific  law  scientific  model  SI  units  solid  substance  temperature  theory  variable  weight  

___________________________________________________________________ Section A. Chemistry: A Science For All Seasons Chemistry is the study of matter and its changes. Matter is anything that has mass and occupies space (called ‘volume’). Examples are gold, air, gasoline, water, computer, etc. Examples of changes are: when gasoline burns to run your car, when the food we eat turn into energy, skin, bones, muscles, etc., when pesticides kill insects, etc.

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Chemistry affects our life every moment and transforms our world. Everything that we do involves chemistry. Look around you, such as the plastic materials that make up part of the computer, the paint on the walls, the clothing that you have on, the paper that you write on, the ink that you write with, and even the reactions that take place in your body which keep you alive. Please check out “What chemists do and where they work” http://www.dummies.com/how-to/content/what-chemists-do-and-where-they-work.html Science and Technology Technology is the direct application of knowledge to solve problems. Science is the process of seeking an understanding of underlying principles of nature. The kind of science that pursues knowledge for its own sake is called basic research, while the one that works toward the solution of a problem (e.g. environmental, health, or industrial) is called applied research. Science and technology involve risks and benefits. One approach is to analyze is the DQ (desirability quotient). If the benefit outweighs the risks, as in pasteurized milk to people of northern European descent, then milk has a high DQ. To people who have lactose intolerance, milk has a low DQ. Technologies such as nuclear power plants, artificial sweeteners and tamoxifen remain controversial. Risk benefit analysis and desirability quotient (DQ). Sample analysis: Saccharin  is  used  in  some  artificial  sweeteners.  Even  those  who  do  not  buy  packets  of  Sweet  and  Low  (brand  name),  use  it  because  it's in dozens of products, from Crest toothpaste and Listerine mouthwash to Robitussin cough syrup and Carefree chewing gum. Saccharin  has  been  shown  to  be  a  carcinogen  (a  substance  that  causes  cancer)  in  animal  tests,  but  there  is  little  evidence  of  carcinogenicity  in  humans.  Studies  have  also  shown  that  artificial  sweeteners  provide  little  benefit  to  those  who  want  to  lose  weight.  The  DQ  for  saccharin  is  uncertain. References:

1. Saccharin: Bittersweet. By: Corcoran, Leila, Jacobson, Michael, Nutrition Action Health Letter, 08857792, Apr98, Vol. 25, Issue 3.

2. Caloric threats from sugarfree drinks? By: J. R., Science News, 00368423, 7/10/2004, Vol. 166, Issue 2.

Note: To find references like this one, use the Health Source Consumer Edition online database accessible from WCC online library (EBSCO) by simply using your family name and student number (MyUH profile). It is recommended that you use this database in other chapters in this course. Science is an accumulation of knowledge about nature and our physical world based on experimentation. It is a unified whole. The various areas of science interact and support one another. Section B. Scientific Method. A process for forming and testing solutions to problems, or theorizing about how or why things work.

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Science has 5 characteristics: testable, reproducible, explanatory, predictive and tentative. Science is limited to studying those things that are observable, and natural processes in which variables can be controlled. The scientific method is dependent on the ability to control all aspects of a problem; this limits the method to “simple” systems. The intellectual process of science starts with experimental observations. There are different paths to scientific discovery. One possible process is shown in Figure 1. Scientific hypotheses (educated guesses) are testable explanations of observed data. These hypotheses are tested by designing and performing experiments. Experiments have two groups created by the procedure: the experimental group and the control group. The experimental group gets some treatment (variable) while the control group does not. The result of the experimental group will determine whether the hypothesis is valid or not. Scientific laws are statements of natural phenomena. (Law of Gravity, Law of Conservation of Mass/Matter, etc). Many scientific laws can be stated mathematically. Example: Boyle’s Law (PV = k) Figure 1. A possible scientific process. A scientific theory is a set of tested hypotheses that explain natural phenomena. Scientific theories are the best current explanations for natural phenomena. Theories are always tentative and may change as observations of nature change. Extra  Credit  #1.    Play  this  animation  and  answer  these  questions.  Worth  2  points.  http://www.sumanasinc.com/webcontent/animations/content/scientificmethod.html  or  just  go  to  Modules  and  click  the  link.  

1. What  is  Redi’s  hypothesis?  2. What  is  Redi’s  experiment?  3. What  is  a  control  group?  In  Redi’s  experiment,  which  jar  is  the  control?  4. What  is  an  experimental  group?  In  Redi’s  experiment,  which  jar  is  the  

experimental  jar?  5. What  is  Pasteur’s  conclusion?  

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Scientists often use models to explain complicated phenomena. Scientific models are tangible items or pictures used to represent invisible processes. One example, shown in Figure 2, illustrates water in an open beaker when left standing slowly disappears. Scientists explain evaporation with a theory based on tiny, invisible particles, called molecules, which are in constant motion. Some of the molecules (colored blue) gain enough energy to escape from the liquid and disperse with air (colored gray) molecules. The molecular picture explains the process of evaporation.

Figure 2. The evaporation of water explained in terms of a scientific model.

Learning Check 1. The first step of the scientific method involves ____________. 2. A pattern or relationship that has been established based on a large amount of

experimental data is a __________. 3. Are data and results two names of the same thing? T/F 4. A control group can differ from an test/experimental group by more than one

variable. T/F 5. A well-substantiated explanation of an aspect of the natural world is a _______.

Chemistry is useful not only in itself but is also fundamental to other scientific disciplines. The application of chemistry has revolutionized biology, medicine, provided materials for powerful computers, construction (engineering), agriculture, and has profoundly influenced other fields of science like psychology, geology, ecology, etc. is a central science. See Figure 3.

Ans. Observation, Law, F, F, Theory

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Figure 3. Chemistry has a central role among the sciences. Section C. Matter: Properties & Classification 1. Matter. Chemistry is the study of matter and its changes. Matter is anything that has mass and volume. Mass (grams) is the measure of the amount of matter in an object. Weight is the pull of gravity on the matter in an object. Weight is mass times gravitational attraction. Volume (liters or cc, cubic centimeters) is the space that matter occupies. Love, hate and other emotions do not have mass and volume. These are not matter. 2. States (or phases) of Matter: Solids, liquids, and gases and their properties.

State Shape Volume Solid Definite Definite

Liquid Not definite Definite Gas Not definite Not definite

A solid occupies a constant volume and has a fixed shape that does not change with an application of pressure. A liquid has a definite volume but its form/shape readily changes according to its container. A gas is diffuse, taking the shape and volume of its container.

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Any sample of matter has three possible physical states. For example, water can be found in the solid (ice), liquid and vapor (gaseous) forms simultaneously as shown in Fig. 4.

Figure 5. The three physical states of matter: solid, liquid and gas. The scientific model of each state is drawn in the circle above it.

On the particle level (see Fig 5), the three states are distinguished by how they are held together. The particles or molecules in the solid are packed and arranged in regular pattern (explaining why solids have definite shape and volume). The particles in the solid state have strong attractive forces. However, when the solid is heated the particles are able to vibrate and when the attractive forces are disrupted, they can move about. This is the liquid state, where the particles or molecules are touching each other (explaining definite volume) but have no regular pattern (explaining why liquids have no definite shape). Further heating causes the particles to move far apart and have no attraction for one another whatsoever. This is the gaseous state (have no fixed volumes and shapes). When you smell someone’s perfume or coffee aroma, this is undoubtedly due to the aromatic particles in the gas phase.

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Learning Check 1. The relative randomness of particles in the states of matter is best indicated by: gas (most random)>liquid>solid (least random). True or False? 2. The attractive forces between particles in solids tend to be stronger than in liquids.

True or False? 3. Physical and Chemical Properties of Matter. Matter exhibits chemical and physical properties. Physical properties are those properties of a substance that can be observed without changing the substance. These are directly observable: color, mass, weight, state (solid, liquid, or gas), and texture, are examples.

Chemical properties are those properties of a substance that can only be studied by forming new substances. Chemical properties tell us how matter will combine to form new and different substances (a chemical change).

Ans. True, True

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4. Physical and Chemical Changes of matter. An excellent demonstration is to burn paper (chemical change) and boil water (physical change). Listen to mp3 file in Laulima. Physical Changes are changes in which the composition of the substance is not changed. Examples: Chopping wood, Melting, Freezing, Boiling, Dissolving sugar in water Physical Changes. When you freeze water or melt ice or boil water to gas/vapor, the chemical composition remains the same. All three states (ice, water or water vapor) contain molecules that are composed of one atom of oxygen and two hydrogen atoms. All three states are represented as follows: water is H2O (l), ice is H2O (s) and steam is H2O (g) Chemical Change. When you pass electricity through water (called electrolysis), this causes the chemical bonds between oxygen and hydrogen in water (H2O) to break. The chemical composition changes from H2O to O2 (oxygen gas) and H2 (hydrogen gas). Conceptual Example: Physical vs. Chemical Change Which of the following events involve chemical changes and which involve physical changes? a.Your hair is cut. b.Lemon juice converts milk to curds and whey. c.Water boils. d.Water is broken down into hydrogen gas and oxygen gas. Solution: Ask yourself, “Is there a change in composition”?If yes, then it is a chemical change. a.Physical change: The composition of the hair is not changed by cutting. b.Chemical change: The compositions of curds and whey are different from the composition of the milk. c.Physical change: Liquid water and invisible water vapor formed when liquid water boils have the same composition; the water merely changes from a liquid to a gas. d.Chemical change: New substances, hydrogen and oxygen, are formed.

Learning Check. Physical or chemical change?

1. A piece of sodium metal fizzes, produces hydrogen gas, when added with water. 2. A piece of paper burns. 3. A red shirt turns white when bleach is added. 4. A stick of margarine melts. 5. A tree is chopped and ground up into sawdust. 6. Sweat on my skin evaporated, cooling my skin.

Ans. C, C, C, P, P, P

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5. Classification of matter based on composition. See Figure 6. Memorize this.

Figure 6. Classification of matter based on composition. Pure substances have constant composition while mixtures have variable composition. A pure substance has only one substance present. It cannot be separated into simpler substances by physical means, such as filtration, chromatography, crystallization or distillation. A pure substance can either be an element or a compound. An element cannot be broken down into simpler substances by chemical or physical means. A compound can be broken down into constituent elements by chemical, but not physical, means. A mixture contains two or more pure substances, and can be separated into those individual components by selection of the appropriate physical methods. Mixtures can be homogeneous: appear uniform throughout (milk, paint, and saltwater are examples) or heterogeneous: show variation in appearance and texture throughout (pizza, raisin bread, and chocolate chip cookies are examples). Elements Compounds Homogeneous Mixture

(Solutions) Heterogeneous Mixture

Gold, Au Salt, NaCl Air (N2, O2) Sand and Salt Carbon, C Water, H2O Salt water (H2O, NaCl) Oil and water Nitrogen, N2 Sugar, C12H22O11 White Gold (Au, Pt) Sand in water

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Elemental symbols: Elements are fundamental building blocks of all matter. All elements are found in the periodic table (Fig. 7). Each element is represented by a symbol containing one to three letters, with the initial letter capitalized. Elemental symbols are made up of one, two or three letters derived from the English (or Latin) name of the element. Only the first letter of the symbol is capitalized. Ex: Cu, C, Na. So far there are 115 elements known (found in the Periodic Table). So, if you see only one capital letter in the formula, then, that means it is representing an element.

Figure 7. The Periodic Table contains all known 115 elements (in symbols). Chemical formulas: Compounds are composed of two or more elements that are chemically combined in fixed ratios. A compound is represented by a chemical formula containing two or more capital letters with numbers to indicate the ratio. Examples: water (H2O), sugar (C12H22O11) and table salt (NaCl). A compound can be separated into simpler substances by chemical means. Water, sugar and salt (sodium chloride) can with more or less difficulty be subdivided into their elements. Heating sugar in a pan to a high degree leaves with a black carbonaceous deposit (carbon, C). Conceptual Examples:  1. Element vs. Compound. Which of the following represent elements and which represent compounds? C Ca HI BN In H2O HBr Answer Elements: C, Ca, and In (each is a single symbol, only one capital letter, meaning one element. Compounds: HI, BN, H2O and HBr (composed of more than one capital letter)

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2. Classify each hypothetical pure substance L and X, as an element or a compound. Indicate when such a classification cannot be made and explain your answer. a) Two elements when mixed combine to form only substance L. b) Substance X is not changed by heating it. Solution a) L is a compound. It is formed by combining two elements. b) X cannot be determined. The absence of a change is not conclusive.

3. Classify the following whether pure substance or mixture. a. sugar b. water c. sugar water  Solution A pure substance contains only one kind of basic building block or one kind of molecule. Every molecule in a pure substance is the same. If there are two or more kinds of molecule present, it is a mixture. a. Sugar has sugar molecules (pure substance). Its chemical formula is C12H22O11. b. Water has water molecules (pure substance). Its chemical formula is H2O. c. Sugar water has both sugar molecules and water molecules (mixture).

4. Classify the following whether homogeneous or heterogeneous mixtures. a. vodka b. pizza c. vinegar d. salt and pepper  Solution A homogeneous mixture has one phase all throughout. It means uniform or indistinguishable. You can’t pick out the components of a homogeneous mixture by looking at it. Vodka and vinegar are homogeneous. Vodka is made of water and alcohol mixed homogeneously. Vinegar is composed of acetic acid and water mixed uniformly. Heterogeneous means varied. Examples: Pizza and salt/pepper mixture. Learning Check. Classify each of the following (element, compound, homogeneous mixture or heterogeneous mixture?

1. Pure water 2. Copper wire 3. 12-carat gold ring 4. Oil and water 5. Carbon monoxide, CO

6. Cobalt, Co 7. Ozone, O3 8. Salad dressing 9. Gasoline 10. Air

Ans. C,E, Homo,Hetero,C,E,E,Hetero,Homo,Homo

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Section D. Measurement of Matter The standard SI system used in science worldwide is based on the metric system. The five most frequently used measurements in this course are mass, length, volume, temperature and time. The base units: kilograms (kg) for mass, meter (m) for length, liter (L) for volume, Kelvin (K) for temperature and seconds (s) for time. Please see Table 1.5.

 Demonstration. Familiarize yourself with these units. Length: meter (m) is slightly more than a yard. Example: The width of your pointy finger is 1 cm (centimeter) Mass: kilogram (kg) is 2.2 pounds (lbs) Example: An ordinary paper clip weighs about 1 gram. Volume: A liter is slightly more than a quart. Example: A sugar cube is approximately 1 mL (milliliter) Note: 1 mL = 1cc = 1 cm3 (cubic centimeter)  The  SI  system  is  easy  to  use  because  it  is  based  on  multiples  of  ten.  Common prefixes are used with the base units to indicate the multiple of ten that the unit represents. Some familiar examples are: 1 kilobyte = a thousand bytes 1 megabyte = a million bytes 1 gigabyte = a billion bytes The most common prefixes used.

Prefix Abbreviation Decimal Equivalent Exponent Equivalent mega M 1,000,000 106 kilo k 1,000 103 deci d 0.1 10-1 centi c 0.01 10-2 milli m 0.001 10-3 micro µ 0.000001 10-6

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Meaning  of  prefix:  The  prefixes  for  the  different  units  of  length,  volume,  and  mass  in  the  metric  system  obey  the  rules  in  the  table  above.  Let’s  apply  the  rules:  All  units  of  mass  in  the  metric  system  are  derived  from  the  gram.  For  example,  the  prefix  cent-­‐  signifies  one-­‐hundredth;  therefore,  one  centigram  is  one  one-­‐hundredth  of  a  gram.            kilo  =1,000     1  kilogram  (kg)  =  1,000  grams  (g)                                                                                    1  gram  (g)  =  1  g          deci  =  0.1         1  decigram  (dg)  =  0.1  g          centi  =  0.01     1  centigram  (cg)  =  0.01  g          milli  =  0.001     1  milligram  (mg)  =  0.001  g          micro  =  0.000001        1  microgram  (1µg)  =  0.000001  g    Conversion  of  units.    If  we  are  converting  to  a  smaller  unit,  simply  move  the  decimal  point  to  the  right.    If  we  are  converting  to  a  bigger  unit  we  move  the  decimal  to  the  left.  Listing  the  units  in  order  from  largest  to  smallest  will  indicate  how  many  places  to  move  the  decimal  point  and  in  which  direction.   Examples of Metric Unit Conversion. Please watch the movie. Convert (a) 1.83 kg to grams (b) 4.16  m  to  cm (c) 729 mL to L

Example  a:    To  convert  1.83  kg  to  g,  write  the  metric  units  in  order,  from  largest  to  smallest.                     kg  __  __  g  dg  cg  mg  __  __  µg      Because  grams  are  three  places  to  the  right  of  kilograms,  converting  kg  to  g  requires  moving  three  decimal  positions  to  the  right;  therefore,  1.83  kg  =  1830  g.  Use  zeroes  to  fill  empty  positions.      Conversion  between  units  of  length  and  volume  in  the  metric  system  follow  the  same  procedure.    Example  b:    Convert  4.16  m  to  cm.    Because  centimeters  are  two  places  to  the  right  of  meters,  converting  4.16  m  to  cm  requires  moving  two  decimal  positions  to  the  right;  therefore,  4.16  m  =  416  cm    Example  c:  To  convert  729  mL  to  L  (liters),  list  the  metric  units  in  order,  from  largest  to  smallest.                    kL  __  __  L  dL  cL  mL  __  __  µL    Since  L  is  three  places  to  the  left  of  mL,  converting  mL  to  L  requires  moving  three  decimal  places  to  the  left;  therefore,  729  mL  =  0.729  L    

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Learning Check. 1. How many meters are equivalent to 2000 mm? 2. How many liters are there in 0.05 mL? 3. Is 0.03 grams the same as 30 mg? 4. How many cm are there in 0.04 mm? More examples on metric conversions. Please watch the following two interactive videos. You need to input your answers to get the video to continue to the very end. http://videone.ivytech.edu/flv/School_of_Health_Sciences/MEAS239/DKNOX/mea_metric_conversions_1/mea_metric_conversions_1.html  http://videone.ivytech.edu/flv/School_of_Health_Sciences/MEAS239/DKNOX/mea_metric_conversions_samples/mea_metric_conversion_samples.html   Mass, Volume and Density Common units of volume (amount of space an object occupies) are liter (L) and cm3 or milliliter (mL). Common unit of mass is grams. Objects of the same mass may have different volumes, while objects of the same volume may have different masses.

1. Density is defined as the amount of matter in a given amount of space. Mathematically, density is the mass-to-volume ratio of matter. Density = Mass ÷ Volume The density of copper is 8.94 g/cm3. The density of gold is 19.3 g/cm3 If you have 1 cm3 of each, which would weigh more, Cu or Au?

Ans. 2.000 m, 0.00005 L, Yes, 0.004 cm

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2. Density of water is 1.0 g/mL (memorize this). Density is usually expressed in g/mL or g/cc units. Note that 1 mL is equivalent to 1 cm3 so the units (g/mL) = (g/cm3) Practice Problem: Density Calculate the density of a metal sample with a mass of 18.96 g and a volume of 4.31 cm3. d = m ÷ v = 18.96 g ÷ 4.31cm3 = 4.40 g/cm3 (use calculator). Demonstration: Understanding the concept of density Density of alcohol = 0.789 g/mL Density of ice = 0.917 g/cm3 Density of water = 1.00 g/cm3 Place a glass of water and a glass of ethyl alcohol side by side and add an ice cube to each one. What do you observe? a. Ice floats on water because the density of ice is less than that of water. b. Ice (floats or sinks?) ____________ on/in ethyl alcohol. Write your conclusion: Conceptual Examples. Answer the following without doing a detailed calculation. (a) Which has the greater volume, a 50.0-g block of copper or a 50.0-g block of gold? (b) Which has the greater mass, 225 mL of ethyl alcohol or 225 mL of hexane? a. It takes a larger block of copper to have a mass of 50.00 g than of gold. A 50.0-g block of copper has a greater volume than a 50.0-g block of gold. b. The density of ethyl alcohol (0.789 g/mL) is greater than that of hexane (0.660 g/mL). Because it is more dense (that is, it has more mass in each unit of volume), 225 mL of ethyl alcohol has a greater mass than does 225 mL of hexane.

Learning Check 1. If we use the units of grams (g) for mass and milliliters (mL) for volume, then the

units for density will be _______. 2. Will a substance with a mass of 1.3 g, and a volume of 1.8 mL float on pure

water? Yes or No? Why? Specific gravity - the ratio of the density of the object in question to the density of pure water at 4oC. Specific gravity is a unitless term. Often the health industry uses specific gravity to test urine and blood samples to diagnose illnesses.

Ans. g/mL, Yes, its density (0.72g/mL) is less.

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Energy: Heat and Temperature 1. Energy is the ability to do work or transfer heat. Energy exists in two major forms: a. Potential energy is stored energy. b. Kinetic energy is energy in motion. 2. Heat vs. Temperature Heat is energy that is transferred from hotter objects to cooler objects. Temperature measure of the average kinetic energy of an object. 3. Temperature Celsius scale unit is (0C). Water freezes at 00C and boils at 1000C. (memorize this) Kelvin scale unit is (K). absolute scale. 0 K = -273 0C Conversion: K = 0C + 273

Celsius to Kelvin Conversion (will be used in Chapter 6)

Practice Problem. Human body temperature is 37 oC. Convert this to Kelvin. K = oC + 273 = 37 + 273 = 310 K

Reminder: please watch all multimedia in Laulima Modules.

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CHAPTER 2- ATOMS OVERVIEW Chapter 2 begins with the Greek concept of matter and moves chronologically forward through the development of the atomistic model of matter. OBJECTIVES After completing this chapter, you should be able to:

1. Discuss  Dalton’s  atomic  theory  and  how  it  explains  experimental  results  2. Compare  atoms  versus  molecules  

ORGANIZATION Chapter 2 – Atoms is divided into four sections: A. History & Three Laws of Matter B. John Dalton & the Atomic Theory of Matter C. Periodic Table D. Atoms vs. Molecules SELF-STUDY/REVIEW QUESTIONS The concepts outlined in the list below do not comprehensively cover every question you may expect to see on the quiz or final exam. Use the following questions as a self-assessment tool to help you gauge your understanding of the course material in this Module. Recall from Chapter 1the following: 1.    A  scientific  law  (a  statement  or  math  equation)  summarizes  experimental  data.  2.    A  theory  is  a  model  that  consistently  explains  observations.  3.    A  compound  contains  two  or  more  elements  chemically  bonded  together.    In  this  chapter,  know  the  following:  1.    A  molecule  is  a  group  of  atoms  that  are  chemically  bonded  together.  2.    Substances  of  different  compositions  are  different  compounds.  3.    The  periodic  table  is  a  systematic  arrangement  of  the  elements.    4.    What  are  the  three  laws  governing  chemical  reactions?  5.    Dalton’s  atomic  theory.  6.    The  key  terms:    atom    atomic  theory  law  of  conservation  of  mass  

law  of  definite  proportions  molecule  periodic  table

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Section  A.    History  A.1.  Atoms:  ~  450  B.C.  The Greek philosophers had two main ideas regarding matter:

1. The continuous view: that matter has endless divisibility 2. The atomistic view (Leucippus and Democritus)                    Atomos:    means  point  at  which  matter  can  no  longer  be  subdivided.

 A.2.  Lavoisier:    The  Law  of  Conservation  of  Mass,  early  1700’s    

Lavoisier was the first great experimentalist. More than anything, he introduced experimental chemistry to Western civilization. Through his experiments he formulated the law of conservation of mass, which states that matter is neither created nor destroyed during a chemical reaction. He is often called the father of modern chemistry.

 

                                   

                                                                       100  g        =        92.61  g    +  7.39  g  During  a  chemical  change,  matter  is  neither  created  nor  destroyed.  In  the  reaction  shown,  the total mass of reactant (HgO, mercuric oxide) is the same as the total mass of products (add 92.61 g Hg, mercury and 7.39 g of Oxygen). Note the “arrow à “ separates the reactants from products.

A.3.  Proust:  The  Law  of  Definite  Proportions,  1799  1. Proust concluded from analyses that elements combine in definite proportions to form compounds. He formulated the law of definite proportions (also called the law of constant composition). Many other scientists further proved the law. This is a strong evidence for the atomistic nature of matter. If matter were continuous, any random ratios by mass would be possible.

 A  compound  always  contains  the  same  elements  in  certain  definite  proportions.  

F2      +    H2    à  2HF  

The  combining  ratio  between  F2  and  H2  is  1:1  (one  F2  is  to  one  H2).  The  chemicals,  F2  and  H2,  do  not  combine  in  any  random  ratio.    In  the  example  above,  since  there  are  only  two  H2,  only  two  F2  reacted.    The  extra  F2  remains  unreacted.      

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   2.  The  law  of  definite  proportions  is  the  basis  for  chemical  formulas.    Because  each  compound  has  a  definite  proportion  of  elements  (same  chemical  formula),  it  does  not  matter  how  a  compound  was  formed.    As  an  example,  regardless  of  the  source,  copper  carbonate  has  the  same  composition  (Cu2(OH)2CO3)  whether  it  a)  occurs  in  nature  as  mineral,  malachite  b)  is  formed  as  patina  on  copper  roofs  or  c)  is  synthesized  in  the  laboratory.    The  law  of  definite  proportions  also  means  that  compounds  have  constant  properties  in  addition  to  constant  composition.  

                 3.    Berzelius  experiment  (below)  illustrates  the  Law  of  Definite  Proportions.    

                         a. The  first  row  indicates  that  10.00  g  of  lead  combines  with  1.55  g  of  sulfur  to  form  

11.55  g  of  lead  sulfide.    b. The  second  row  indicates  that  you  will  still  get  11.55  g  of  lead  sulfide  from  10.00  g  

of  lead  even  if  you  combine  it  with  3.00  g  of  sulfur.    Some  of  the  sulfur  will  not  react.  

c. The  third  row  indicates  that  you  will  still  get  11.55  g  of  lead  sulfide  from  1.55  g  of  sulfur  even  if  you  combine  it  with  18.00  g  of  lead.  Some  of  the  lead  will  not  react.  

d. Moral  lesson:  the  combining  ratio  is  fixed!    

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A.4.  John  Dalton:    Law  of  Multiple  Proportions,  1803  

– Elements  may  combine  in  more  than  one  set  of  proportions;  each  set  corresponds  to  a  different  compound.  Example  below  shows  carbon  and  oxygen  can  form  two  different  compounds:  carbon  monoxide  (CO)  and  carbon  dioxide  (CO2).      

 

   

– The  Law  of  Multiple  Proportions  states  that  the  masses  of  one  element  which  combine  with  a  fixed  mass  of  the  second  element  are  in  a  ratio  of  whole  numbers.    

– The  mass  of  oxygen  that  combined  with  12  g  of  carbon  in  CO  is  16  g,  while  the  mass  of  O  that  combined  with  12  g  of  C  in  CO2  is  32  g.    Dividing  32g  by  16g  is  2.    This  means  the  masses  of  O  that  combine  with  C  are  in  a  ratio  of  2:1  (consistent  with  Law  of  Multiple  Proportions.).  

Carbon monoxide Carbon dioxide

CO CO2

Produced in small amount

when fuels burn

Becomes solid (dry ice)

when cooled at -800C

Deadly even at 0.2% in air Fizz in softdrinks

Section  B.    John  Dalton’s  Atomic  Theory  of  Matter.  John  Dalton  proposed  the  atomic  theory  to  explain  the  three  laws  of  matter  (stated  above).      

1.    All  matter  is  composed  of  extremely  small  particles  called  atoms.  2.    All  atoms  of  a  given  element  are  alike  and  differ  from  the  atoms  of  any  other  element.  3.    Compounds  are  formed  when  atoms  of  different  elements  combine  in  fixed  proportions.  4.    A  chemical  reaction  involves  the  rearrangement  of  atoms.    

According to Dalton, elements are distinguished from each other by the weights of their atoms. His explanations about the following are still accepted nowadays:

• Elements are composed of only one kind of atom.

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• Compounds  are  composed  of  two  or  more  kinds  of  atoms  chemically  combined  in  definite  proportions.

• Matter  must  be  atomic  to  account  for  the  law  of  definite  proportions. • The  rearrangement  of  atoms  explains  the  law  of  conservation  of  mass. • The  existence  of  atoms  explains  how  multiple  proportions  can  exist.

These are the modern modifications of Dalton’s Atomic Theory:

1. Atoms can be divided, as we shall see in Chapter 11-Nuclear Chemistry. An atom is broken apart in a nuclear reaction producing an atom of different identity.

2. Atoms of the same element can have different masses (called isotopes) as we will see in the next chapter.

 Section  C.  The  Periodic  Table  

By the mid 1800’s 55 elements were known, but no successful way existed to classify them. Mendeleev (1869) arranged a table of elements according to increasing atomic weights, placing elements with similar properties in the same column. He left gaps for yet undiscovered elements. He predicted the properties of those elements. When those elements were eventually discovered, many of his predictions were found to be accurate. This was the beginning of the modern periodic table. At present there are 115 known elements.

 Section  D.    Atoms  Versus  Molecules     1. Atoms are the smallest particles of an element. It is possible

to observe computer-enhanced images of atoms. Atoms are not destroyed in chemical reactions thus, can be recycled. 2. A molecule is a group of atoms chemically bonded together. Hence, molecules can be decomposed or divided into atoms by chemical means. Example: A molecule of water (chemical formula, H2O) is composed of two atoms of hydrogen (H) bonded to an atom of oxygen (O). When electric current is passed through water, a chemical reaction occurs, splitting water into hydrogen and oxygen.

Elements and compounds. Elements are composed of one type of atom. Example: all the invisible particles of gold are called gold atoms, which are all of the same size, shape and color. Atoms are the smallest particles of an element. Many compounds exist as groups of atoms bonded together as units called molecules. A molecule is considered the smallest constituent particle of a compound. Important exceptions to the rule are the eight (8) elemental molecules: hydrogen (H2), nitrogen (N2), oxygen (O2), ozone (O3), fluorine (F2), chlorine (Cl2), bromine (Br2), iodine (I2).

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Conceptual Examples. a. Which of the following are atoms and which are molecules? b. Which of the following are elements and which are compounds?

Solution a. Each chemical above is composed of two or more atoms. Hence, all the above are molecules. b. The top row (hydrogen, oxygen, nitrogen and chlorine molecules) are elements. The bottom row molecules are compounds. Physical and chemical changes: Revisited using atomic scientific models.

Physical Change composition of molecules remain the same

only the distance between molecules change

Chemical Change incurs a change in composition of the

molecules

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Learning Check a. Which of the following are atoms and which are molecules? b. Which of the following are elements and which are compounds?

Ans. Ar, Ne, He. Kr and Xe are atoms (and elements). N2, H2 and O2 are molecules (and elements). CH4 and CO2 are molecules/compounds.

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Self-Assessment for Chaps 1 & 2. (Answers are in the Appendix)

1. Which of the following is NOT matter? a. Heart b. Love c. Food d. Water

2. Which of the following is

NOT a compound?

3. No scientist has yet been able to decompose oxygen. Oxygen must then be a. An element b. A compound c. Homogeneous mixture d. Heterogeneous mixture

4. The statement, “ in a chemical reaction, the mass of the products is never less than or

more than the total mass of reactants” is an example of a. Hypothesis b. Theory c. Law d. Experiment

5. A number of people become ill after eating in a restaurant. Which of the following is a hypothesis?

a. Everyone who ate the oysters got sick. b. Symptoms include nausea and dizziness. c. Bacteria in the oysters may have caused the illness. d. Identifying the bacteria in the oyster is the logical thing to do.

6. What is the mass in grams of an object having a mass of 1.2 kg?

a. 0.0012 g b. 12 g c. 120 g d. 1200 g

7. A solid with a volume of 6.0 mL and a mass of 12 g has a density of ____?

a. 0.5 g/mL b. 60 g/mL c. 2.0 g/mL d. 2.0 mL/g

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8. Which of the following has the lowest density? Refer to the drawing. a. Cork b. water c. brass d. mercury

9. What class of matter is breakfast coffee? a. element b. compound c. homogeneous mixture d. heterogeneous mixture 10. Hair shampoo (check the ingredients on the label) is an example of

a. element b. compound c. solution d. heterogeneous mixture

11. Classify the type of change occurring when a cake is baked from flour, baking powder, sugar, milk, eggs and butter.

a. physical b. chemical 12. The smallest particle of an element is a. atom b. compound c. mixture

13. The chemical combination of two or more atoms in fixed amounts is called a a. mixture b. compound c. element 14. Copper sulfate can be further subdivided into simpler substances by chemical means only. Therefore, it is a/an ________ a. mixture b. compound c. element 15. Two chemicals, A and B, react completely to yield products C and D. If 4 g of A and 8 g of

B are reacted (with no excess of either), and 5 g of C are produced, how much D is produced? a. 5 g b. 9 c. 7 g d. 3 g 16. Which of these depicts a chemical change in the submicroscopic level? Why? a. b.

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Chaps 1 & 2 Worksheet Homework (by Prof. B. Reeves). 1. Chemistry studies ___________ and its ________. 2 A ____________ is an educated guess. 3. A ___________ is a set of tested hypotheses that explain natural phenomena and can be changed with new research. 4. A scientific ______________ are descriptions of natural phenomena that are found to be true to a long period of time. 5. A scientific _____________ can be pictures which help us understand an invisible process. 6. Describe solid, liquid and gas in terms of shape and volume. 7. Identify as a physical change or a chemical change a. tearing paper b. burning wood. c. melting ice d. boiling water e. iron rusts f. silver tarnishes 8. Identify as a physical property or chemical property a. tarnishing b. melting c. freezing d. explodes e. red f. hard

9. Identify as a compound or element a. CO b. Co c. Pb d. S e. Si f. H2O2 g. HCN h. NO i. N j. CO2

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10. Use the chart to convert a. 35 mm to m b. 25 cm to m c. 450 km to m d. 0.003 mm to m e. 0.003 m to mm f. 452 km to cm g. 425 mm to km h. 500 k to mm i. 500 mm to km j. 250 m to mm 11. Identify as a homogeneous or heterogeneous mixture. a. oil and water b. sand c. salt water

d. salt and pepper e. Kool Aid drink f. breakfast tea

12. What is the density? a. mass = 25 g and volume = 3mL b. mass = 2.4 lbs and volume = 40 square inches. 13. Convert 15oC to K and 298K to oC 14. Fill in the blanks. 1. The Law of ____________ _____________ states that elements combine in definite proportions to form compounds. 2. The Law of ____________ _________________ says that elements may combine in more than one set of proportions, each set corresponds to a different compound. 3. Dalton’s Atomic Theory says that ____________ are made of one kind of atom, _____________ are formed by 2 or more kinds of atoms chemically combined, and atoms can be rearranged during chemical _____________ 4. Two things that are wrong in Dalton’s Theory are due to ________________ reactions where atoms can be changed to a different identity and atoms exist as _______________ because they have different masses. 5.Elements consist of _______________ one type of atom and ______________ consist of more than one type of atom chemically combined.