1. Chemistry Chapter 1 Chemical Foundations Dr. Daniel Schuerch

2. Scientific Method 1.Start by asking a question regarding a phenomena, and then making observations of the phenomena 2.As prior research is evaluated, observations are made, and data is collected –Data may be quantitative (numerical) or qualitative (descriptive) 3.Scientist then use the data to make inferences regarding the phenomena observed –An inference is logical interpretation of the phenomena based on prior knowledge or experience

3. The Scientific Method 4.After initial observations, the scientist will propose one or more hypotheses –A hypothesis is proposed scientific explanation for a set of observations THAT CAN BE TESTED! –Hypotheses are generated using prior knowledge, logical inferences, and informed creative imagination –Hypotheses may compete to explain the same phenomena –Must be proposed in a way that can be tested Controlled experiments Collecting more data 5.Scientist will design and implement controlled experiments to test their hypotheses –After testing, some hypotheses will be ruled out, some will be confirmed

4. The Scientific Method 6.A series of tested hypotheses may then be used to develop a scientific theory or model that explains various observations about a phenomena –A scientific theory is a well-substantiated explanation of some aspect of the natural world, based on a body of facts that have been repeatedly confirmed through observation and experiment Theories (models) result in more predictions that are tested through experimentation, that may be used to modify the theory as needed –Sometimes testing of hypotheses lead to statement called a natural law –A natural law is a summary of observed (measurable) behavior A law tells what happens; a theory (model) is our attempt to explain why it happens

5. Using and expressing measurements A measurement is a quantity that has both a number and a unit –Examples of measurements include? –Measurements are fundamental to the experimental sciences –Thus, it is important to be able to make measurements and to decide whether a measurement is correct –In science, the units of measurement are those of International System of Measurements (SI)

6. Scientific Notation AKA Exponential Notation In science/chemistry, we often encounter very small or very large numbers that are cumbersome to work with if we write out the entire number –Example, one gram of Hydrogen contains 602,000,000,000,000,000,000,000 atoms –One atom of Gold has a mass of 0.000000000000000000000327 gram Using scientific notation: –One gram of Hydrogen contains 6.02 x 10 23 atoms 6.02 x 10 23 is read six point zero two times ten to the twenty- third –One atom of Gold has a mass of 3.27 x 10 -22 grams 3.37 x 10 -22 is read three point three seven time 10 to the negative twenty-second

7. Scientific Notation Continued Numbers greater than One 6.02 X 10 23 CoefficientExponent 602,000,000,000,000,000,000,000 1.The coefficient, in scientific notation, is always a number between 1 and 10, thus move the decimal place to the left for number greater than 1, until you have a number between 1 and 10 2.To determine the exponent, count how many places the decimal moved to the left

8. Scientific Notation Continued Numbers less than One 3.27 X 10 -22 CoefficientExponent 0.000000000000000000000327 1.The coefficient, in scientific notation, is always a number between 1 and 10, thus move the decimal place to the right for numbers less than 1, until you have a number between 1 and 10 2.To determine exponent, count how many places the decimal moved to the right and add a negative sign before the exponent

9. Practice Give the scientific notation for: 30,000 12 0.2.0000132 1,000,320 Write out the number for the following scientific notations 1.23x10 1 1.23x10 -1 1.23x10 -4 1.23x10 4 Are the numbers 1.230x10 4 and 123.0x10 2 different numbers?

10. When adding or subtracting numbers written in scientific notation, you must be sure that the exponents are the same before doing the arithmetic. Adding and Subtracting Using Scientific Notation Add 2 x 10 2 + 2 x 10 3 Subtract 2 x 10 3 - 1 x 10 2

11. Multiplying and dividing also involve two steps, but in these cases the quantities being multiplied or divided do not have to have the same exponent. Multiplying and Dividing Using Scientific Notation For multiplication, you multiply the coefficients. Then, you add the exponents. For division, you divide the coefficients. Then, you subtract the exponent of the divisor from the exponent of the dividend.

12. Suppose you are asked to solve the following problems. Multiplying and Dividing Numbers in Scientific Notation

13. Calculator In most calculators, scientific notation is given as something 1.0E11 were E stands for exponent (x10 # ) The button for exponent may be labeled exp or E, EE, or EXP Come see your teacher if you have any doubts about what button on your calculator is the exponent key

14. On your own, use your calculator to calculate the following: 1.1.35x10 5 · 1.65x10 4 2.1.35x10 5 - 1.65x10 4 3.1.35x10 5 + 1.65x10 4 4.1.35x10 5 ÷ 1.65x10 4 5.1.35x10 -4 ÷ 1.65x10 -10 6.1.35x10 -4 ÷ 1.65x10 -2

15. On your own, use your calculator to calculate the following: 1.1.35x10 5 · 1.65x10 4 = 2.23x10 9 2.1.35x10 5 - 1.65x10 4 = 1.19x10 5 3.1.35x10 5 + 1.65x10 4 = 1.52x10 5 4.1.35x10 5 ÷ 1.65x10 4 = 8.18 5.1.35x10 -4 ÷ 1.65x10 -10 = 8.18x10 5 6.1.35x10 -4 ÷ 1.65x10 -2 = 8.18x10 -3

16. Accuracy and Precision Accuracy is a measure of how close a measurement comes to the actual or true value of whatever is measured –To evaluate the accuracy of a measurement, the measured value must be compared to the correct value Precision is a measure of how close a series of measurements are to one another –To evaluate the precision of a measurement, you must compare the values of two or more repeated measurements

17. Accuracy and Precision

18. Accuracy and Precision Practice Method OneMethod TwoMethod ThreeMethod Four 100.1 cm108 cm108.0 cm110.9 cm 99.9 cm105 cm108.0 cm130.9 cm 100.0 cm95 cm108.3 cm95.0 cm 99.9 cm100 cm107.7 cm95.7 cm 100.1 cm92 cm108.0 cm127.0 cm

19. Systematic and Random Error Method OneMethod TwoMethod ThreeMethod Four 100.1 cm108 cm108.0 cm110.9 cm 99.9 cm105 cm108.0 cm130.9 cm 100.0 cm95 cm108.3 cm95.0 cm 99.9 cm100 cm107.7 cm95.7 cm 100.1 cm92 cm108.0 cm127.0 cm

20. Determining Error Accepted value – the correct value of a measurement based on reliable references Experimental value – the actual value you measure Error – is the difference between the experimental value and the accepted value Error = Experimental value – Accepted value Percent error is the absolute value of the error divided by the accepted value, multiplied by 100% Percent Error = Error Accepted Value X 100% Percent Error = Experimental value – Accepted value Accepted Value X 100% Or

21. What is the percent error of a measured value of 114 lb if the person’s actual weight is 107 lb? Percent Error = Experimental value – Accepted value Accepted Value X 100%

22. Significant Figures Significant figures in a measurement include all of the digits that are known, plus a last digit that is estimate –Example on white board Measurements must always be reported in significant figures because calculated answers often depend on the number of significant figures in the values used in the calculation

23. Rules for determining whether a digit in a measured value is significant (its all about the zeros, almost): Every non zero (0) digit is assumed to be significant –123.1 has four significant digits Zeros appearing between nonzero digits are significant –1000.1 has five significant digits Left most zeros in front of nonzero digits are not significant –0.0001 has one significant digit Zeros at the end of a number and to the right of a decimal point are always significant –1.0000 has five significant digits Zeros at the rightmost end of a measurement that lie to the left of an understood decimal place are not significant if they serve as placeholders to show the magnitude of the number –23000 has two significant digit, however, 23000.0 has 6 significant figures There are two situations in which numbers have an unlimited number of significant figures 1.Absolute values such as 23 people in the class (no estimation or fractions of people, you either have 23 people or don’t have 23 people) 2.Exactly defined quantities such as those found within a system of measurements (100 cm in one meter or 1000 ml in 1 liter)

24. Significant Figures in Calculations Calculated answers cannot be more precise than the least precise measurement from which it was calculated –When adding or subtracting, the result should be rounded to the same number of decimal places as the measurement with the least number of decimal places –In calculations involving multiplication and division, you need to round the answer to the same number of significant figures as the measurement with the least number of significant figures Work some example problems on chalk board

25. The International System of Units The international system of units (SI) is a revised version of the metric system adopted by international agreement in 1960 –Why is it important for us to have an international system of units?

26. Base Units There are seven base units in SI A base unit is a defined unit in a system of measurement that is based on an object or event in the physical world A base unit is independent of other units

27. The Seven Base SI Units QuantitySI Based UnitSymbol LengthMeterm MassKilogramKg TemperatureKelvinK TimeSeconds Amount of SubstanceMolemol Luminous intensityCandelacd Electric currentAmpereA All other SI units are derived from the base units above

28. Metric Prefixes PrefixAbbreviationMeaningExample MegaM10 6 1 megameter = 1,000,000 m Kilok10 3 1 kilometer = 1,000 m Decid10 -1 1 decimeter = 0.1 m Centic10 -2 1 cenitmeter = 0.01 m *Millim10 -3 1 millimeter = 0.001 m Microµ10 -6 1 micrometer = 0.000001 m Nanon10 -9 1 nanometer = 0.000000001 m Picop10 -12 1 picometer = 0.000000000001 m Femtof10 -15 1 femtometer = 1 x 10 -15 Range of human eye?

29. Units of Length The SI base unit for length is the meter (m) A meter is the distance that light travels through a vacuum in 1/299 792 458 of a second A meter, which is close in length to a yard, is useful for measuring the length and width of a room Common metric units of lengths include the centimeter, meter, and kilometer

30. Mass The SI base unit for mass is the kilogram (kg) Mass is a measure of the amount of matter A kilogram of matter on earth exerts a force of approx. 2.2 pounds –How is mass different from weight? Weight is a force that measures the pull on a given mass by gravity The kilogram is defined by a platinum-iridium metal cylinder Common metric units of mass include the kilogram, gram, milligram, and microgram

31. The SI base unit for time is the second (s) The frequency of microwave radiation given off by a cesium-133 atom is the physical standard used to establish the length of a second Time

32. Temperature Temperature is a measure of how hot or cold an object is –Determines the direction of heat transfer Always from object of higher temperature to lower temperature –Almost all substances expand when heated and contract when cooled –Scientist use two equivalent units of temperature, the degree Celsius and Kelvin

33. Units of Temperature The Celsius scale was devised by Anders Celsius, a Swedish astronomer –He used the temperatures at which water freezes and boils to establish his scale because these temperatures are easy to reproduce He defined the freezing point as 0 and the boiling point as 100 Then he divided the distance between these points into 100 equal units, or degrees Celsius

34. Temperature Scales The Kelvin scale was devised by a Scottish physicist and mathematician, William Thomson, who was known as Lord Kelvin A kelvin (K) is the SI base unit of temperature –Notice no degree sign used –A change of one kelvin is equal to one degrees C –0 K is known as absolute zero or -273 ºC –On the Kelvin scale, water freezes at about 273 K and boils at about 373 K

35. Temperature Conversions It is easy to convert from the Celsius scale to the Kelvin scale To convert temperatures reported in degrees Celsius into kelvins, you just add 273 K = ºC + 273 To convert kelvin to degrees C ºC = K - 273

36. Temperature Conversion

37. Energy Energy is the capacity to do work or to produce heat –The Joule (J) and the calorie (cal) are common units of energy –The Joule is SI unit of energy, named after English physicist James Prescott Joule –The calorie is the quantity of heat that raises the temperature of 1 gram of pure water by 1ºC 1J = 0.2390 cal1cal = 4.184J

38. Derived Units Not all quantities can be measured with base units. For example, the SI unit for speed is meters per second (m/s). Notice that meters per second includes two SI base units—the meter and the second. A unit that is defined by a combination of base units is called a derived unit –Two other quantities that are measured in derived units are volume and density

39. Derived Units Volume Volume is the space occupied by an object The derived unit for volume is the cubic meter, which is represented by a cube whose sides are all one meter in length For measurements that you are likely to make, the more useful derived unit for volume is the cubic centimeter (cm 3 ) also known as the ml.

40. Derived Units Volume Liter

42. Derived Measurements Density Density is a ratio that compares the mass of an object to its volume The units for density are often grams per cubic centimeter (g/cm 3 ). You can calculate density using this equation: Density is a property that can be used to identify an unknown sample of matter. Every sample of pure aluminum has the same density D M V

43. Derived Measurements Density Density is an intensive property that depends only on the composition of a substance, not on the size of the sample The density of a substance generally decreases as its temperature increases –Water is an exception

44. Dimensional Analysis Dimensional analysis is a method of problem-solving that focuses on the units used to describe matter

45. Dimensional Analysis Dimensional Analysis often uses conversion factors  A conversion factor is a ratio of equivalent values used to express the same quantity in different units  Conversion factors are derived from equivalence statements Equivalence statement Conversion Factors

46. Dimensional Analysis A conversion factor is always equal to 1. Because a quantity does not change when it is multiplied or divided by 1, conversion factors change the units of a quantity without changing its value.

47. Dimensional Analysis Suppose you want to know how many meters are in 48 km. –You need an equivalence statement that relates kilometers to meters –Then, you need to make the conversion factor that will solve your problem

48. Dimensional Analysis

49. Conversion Factors Equivalence Statements

50. Graphing Graph the following data: Concentration of X (Mole/L) Absorbance 650 nm 0.500.201 1.000.400 1.500.610 2.000.790 2.500.971 3.001.221 3.501.401 4.001.611 One ml of solution containing an unknown concentration of X was diluted 1/10 with distilled water. The absorbance 650 of the diluted solution was 1.10. What was the original concentration of X in the solution.

51. Graphing Concentration of X in the 1/10 dilution is 2.75 mol/L. The concentration of X in the undiluted sample is 27.5 mol/L

52. Substances Matter that has a uniform and definite composition is called a substance –Examples of substances? –Does a substance have to be pure? Every sample of a given substance has identical intensive properties because every sample has the same composition –A given substance has identical physical properties –Physical properties are a quality or condition of a substance that can be observed or measured without changing the substance’s composition Examples of physical properties?

53. States of Matter There are three states of Matter –Solid – has a definite shape and volume and the shape does not depend on containment –Liquid – has an indefinite shape, flows, yet has a fixed volume –Gas – matter that takes both the shape and volume of its container Vapor is different from a gas –Vapor describes the gaseous state of a substance that is generally a liquid or solid at room temperature –Gases exist in gaseous state at RT What about plasma?

54. Physical Changes Physical Change a change in which some properties of a material change, but the composition of the material does not change –Examples of physical changes include: freezing, melting, boiling, condensing, magnetism, conductivity, hydrophobicity, density, color –What are some other physical changes? Physical changes can be classified as reversible or irreversible –Examples of reversible and irreversible physical changes are?

55. Mixtures A mixture is a physical blend of two or more components –Mixtures can be classified as heterogeneous or homogeneous based on the distribution of their components Heterogeneous mixtures are those that do not have a uniform composition throughout, more than one phase –Heterogeneous mixtures are often called suspensions Homogeneous mixtures have a uniform composition throughout, one phase –Homogeneous mixtures are often solutions consisting of solutes dissolved in a solvent »Solutions can be solid, liquid, or gas »Phase is used to describe any part of a sample with uniform composition and properties –What are some examples of homogeneous and heterogeneous mixtures?

56. Separating Mixtures The components of a mixture can be separated from each based each components different physical properties –Examples Solids can be separated from liquids or gases using filtration Liquids can be separated from each other by distillation, which takes advantage of the different boiling points of two liquids What are some other ways you could separate mixtures? –How would you separate a mixture of iron filling and salt?

57. Chromatography Utilizes a mobile and stationary phase –Mobile phase is a liquid or gas –Stationary phase is a solid The mobile phase caries the mixture to be separated over the stationary phase which separates the mixture based on various physical properties –Size –Solubility –Hydrophobicity –Charge

58. Paper Chromatography Distance solvent has traveled up paper Ink Band Chromatography Paper Solvent Time = 0 minTime = 10 min Time = 20 min 1.What is the mobile phase? 2.What is the stationary phase? 3.How many different substances was the ink made of? 4.Which color substance interacted with stationary phase the least? 5.Which color substance moved the slowest?

59. Elements and Compounds An element is simplest form of matter that has a unique set of properties –What are some examples of elements? A compound is a substance that contains two or more elements chemically combined in a fixed proportion –What does fixed proportion mean? –What are some examples of compounds Compounds can be broken down into simpler substances by chemical means, but elements can’t –What about nuclear fission?

60. Breaking down Compounds Physical means used to separate mixtures cannot be used to break compounds into simpler substances –Breaking down compounds involves chemical change Chemical change is a change that produces matter with a different composition and physical properties than the original matter Breaking down compounds is finite –A chemical can only be broken down to its elements, after which, no further break down can occur

61. Properties of Compounds The properties of compounds are different from those of their component elements + Sodium Metal (Will cause explosion if mixed with water) Chlorine Gas Deadly if inhaled Salt (Sodium Chloride) Required for you survival Element Compound

62. Distinguishing Substances and Mixtures If the composition of a material is fixed, the material is a substance If the composition of a material my vary, the material is a mixture Matter Substance Definite composition (homogeneous) Mixture of Substances Variable composition CompoundElement Homogeneous Mixture Uniform, AKA solution Heterogeneous Mixture Non-uniform, distinct phases Can be separated chemically