1. General, Organic and Biochemistry 7 th Edition GENERAL CHEMISTRY Copyright  The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. 1.1 The Discovery Process Chemistry - The study of matter… –Matter - Anything that has mass and occupies space A table A piece paper –What about air? Yes, it is matter

3. 1.1 The Discovery Process Chemistry: the study of matter its chemical and physical properties the chemical and physical changes it undergoes the energy changes that accompany those processes Energy - the ability to do work to accomplish some change

4. 1.1 The Discovery Process MAJOR AREAS OF CHEMISTRY Biochemistry - the study of life at the molecular level Organic chemistry - the study of matter containing carbon and hydrogen Inorganic chemistry - the study of matter containing elements, not organic Analytic chemistry - analyze matter to determine identity and composition 1

5. 1.1 The Discovery Process Physical chemistry - attempts to explain the way matter behaves CHEMISTRY medical practitioners pharmaceutical industry forensic sciences food science public health

6. Why study it? –Intellectual enterprise – provides a means of explaining our material world Make predictions Manipulate matter for specific purposes –Figures prominently in all sciences (biology, geology, astronomy, etc.) –Practical applications (e.g. life-saving medicines, food, sanitation, etc.) –Chemistry is everywhere, literally…

7. 1.1 The Discovery Process THE SCIENTIFIC METHOD The scientific method - a systematic approach to the discovery of new information Characteristics of the scientific process  Observation  Formulation of a question  Pattern recognition  Developing theories  Experimentation  Summarizing information 2

8. Technique or approach: –Investigating phenomena –Acquiring new knowledge –Correcting/integrating previous knowledge Identifiable features: –Gathering observable, empirical and measureable evidence –Collection of data through experimentation/observation

9. 1.1 The Discovery Process

10.  For scientists…?  Not exclusive to science, simply an approach of investigation  How does it affect our daily lives?  Science is everywhere  Technology  Medicine  Food  Environment  Economics  Energy

11.  Provides ‘tangible’ results  Responsible for all of the luxuries that we currently enjoy  Objective/self correcting  Verifiable (you can do the experiments yourself)  Effects laws/policies/guidelines/etc. (insofar as it contributes to the accumulation of scientific knowledge)  Civic duty to have a fundamental understanding of science  Which requires a thorough understanding of the scientific method and how it works

12. 1964 – Barnett Rosenberg et al., Michigan State University Studying the effects of electricity on bacterial growth Used Pt electrodes and noticed that cell division ceases when electrode is inserted Led to the discovery of the anticancer properties of cisplatin

13. Observation – cell division ceases when using Pt electrodes Question – why does this type of electrode kill cells? Hypothesis – certain Pt compounds inhibit cell division Prediction – Pt compounds will be found in these cell cultures Test – cisplatin identified (supports hypothesis) Further testing – added cisplatin to cells and observed its effects

14. 1.gas - particles widely separated, no definite shape or volume 2. liquid - particles closer together, definite volume but no definite shape 3. solid - particles are very close together, define shape and definite volume 1.2 Matter and Properties Three States of Matter 4

15. Three States of Water (a) Solid (b) Liquid (c) Gas

16. Comparison of the Three Physical States 1.2 Matter and Properties

17. Physical property - is observed without changing the composition or identity of a substance Physical change - produces a recognizable difference in the appearance of a substance without causing any change in its composition or identity -conversion from one physical state to another -melting an ice cube 5

18. Separation by Physical Properties Magnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry.

19. 1.2 Matter and Properties Chemical property - result in a change in composition and can be observed only through a chemical reaction Chemical reaction (chemical change) - a process of rearranging, removing, replacing, or adding atoms to produce new substances hydrogen + oxygen  water reactants products 5

20. 1.2 Matter and Properties Classify the following as either a chemical or physical property: a.Color b.Flammability c.Hardness d.Odor e.Taste

21. 1.2 Matter and Properties Classify the following as either a chemical or physical change: a.Boiling water becomes steam b.Butter turns rancid c.Burning of wood d.Mountain snow pack melting in spring e.Decay of leaves in winter

22. Intensive properties - a property of matter that is independent of the quantity of the substance -Density -Specific gravity Extensive properties - a property of matter that depends on the quantity of the substance -Mass -Volume 1.2 Matter and Properties 6

23. Pure substance - a substance that has only one component Mixture - a combination of two or more pure substances in which each substance retains its own identity, not undergoing a chemical reaction Classification of Matter 7

24. 1.2 Matter and Properties Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction Compound - a substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio Classification of Matter 7

25. 1.2 Matter and Properties Mixture - a combination of two or more pure substances in which each substance retains its own identity Homogeneous - uniform composition, particles well mixed, thoroughly intermingled Heterogeneous – nonuniform composition, random placement Classification of Matter 7

26. Classes of Matter 1.2 Matter and Properties

27. 1.3 Significant Figures and Scientific Notation Information-bearing digits in a number are significant figures The measuring device used determines the number of significant figures a measurement has The amount of uncertainty associated with a measurement is indicated by the number of digits or figures used to represent the information 8

28. 1.3 Significant Figures and Scientific Notation Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit

29. 1.3 Significant Figures and Scientific Notation Recognition of Significant Figures All nonzero digits are significant 7.314 has four significant digits The number of significant digits is independent of the position of the decimal point 73.14 also has four significant digits Zeros located between nonzero digits are significant 60.052 has five significant digits

30. 1.3 Significant Figures and Scientific Notation Use of Zeros in Significant Figures Zeros at the end of a number (trailing zeros) are significant if the number contains a decimal point. 4.70 has three significant digits Trailing zeros are insignificant if the number does not contain a decimal point. 100 has one significant digit; 100. has three Zeros to the left of the first nonzero integer are not significant. 0.0032 has two significant digits

31. 1.3 Significant Figures and Scientific Notation How many significant figures are in the following? 1.3.400 2.3004 3.300. 4.0.003040

32. 1.3 Significant Figures and Scientific Notation Scientific Notation Used to express very large or very small numbers easily and with the correct number of significant figures Represents a number as a power of ten Example: 4,300 = 4.3  1,000 = 4.3  10 3

33. For numbers greater than 1, the original decimal point is moved x places to the left, and the resulting number is multiplied by 10 x The exponent x is a positive number equal to the number of places the decimal point moved 5340 = 5.34  10 3 What if you want to show the above number has four significant figures? = 5.340  10 3 1.3 Significant Figures and Scientific Notation

34. To convert a number less than 1 to scientific notation, the original decimal point is moved x places to the right, and the resulting number is multiplied by 10 –x The exponent x is a negative number equal to the number of places the decimal point moved 0.0534 = 5.34  10 –2 1.3 Significant Figures and Scientific Notation

35. Convert the following to 2 SF using scientific notation… 0.0024 48.20 224 0.0180

36. Types of Uncertainty Error - the difference between the true value and our estimation –Random –Systematic Accuracy - the degree of agreement between the true value and the measured value Precision - a measure of the agreement of replicate measurements 1.3 Significant Figures and Scientific Notation

37. Exact and Inexact Numbers Inexact numbers have uncertainty by definition (any measured value) Exact numbers are a consequence of counting A set of counted items (beakers on a shelf) has no uncertainty Exact numbers by definition have an infinite number of significant figures

38. 1.3 Significant Figures and Scientific Notation =3.35  10 4 Rules for Rounding Off Numbers When the number to be dropped is less than 5 the preceding number is not changed When the number to be dropped is 5 or larger, the preceding number is increased by one unit Round the following number to 3 significant figures: 3.34966  10 4

39. Round off each number to three significant figures: 1.61.40 2.6.171 3.0.066494 1.3 Significant Figures and Scientific Notation

40. 1.4 Units and Unit Conversion Units - the basic quantity of mass, volume or whatever is being measured –A measurement is useless without units 9 English system - a collection of functionally unrelated units –Difficult to convert from one unit to another –1 foot = 12 inches = 0.33 yard = 1/5280 miles Metric System - composed of a set of units that are related to each other decimally, systematic –Units relate by powers of ten

41. 1.4 Units and Unit Conversion Basic Units of the Metric System Massgramg Lengthmeterm VolumeliterL Basic units are the units of a quantity without any metric prefix.

42. 1.4 Units and Unit Conversion

44. UNIT CONVERSION You must be able to convert between units -within the metric system -between the English system and metric system The method used for conversion is called the Factor-Label Method or Dimensional Analysis !!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!! 9

45. 1.4 Units and Unit Conversion Let your units do the work for you by simply memorizing connections between units. –For example: How many donuts are in one dozen? –We say: “Twelve donuts are in a dozen.” –Or: 12 donuts = 1 dozen donuts What does any number divided by itself equal? ONE!

46. 1.4 Units and Unit Conversion This fraction is called a unit or conversion factor Multiplication by a unit factor does not change the amount – only the unit

47. Example: How many donuts are in 3.5 dozen? You can probably do this in your head but try it using the Factor-Label Method. 1.4 Units and Unit Conversion

48. Start with the given information... 3.5 dozen Then set up your unit factor... See that the units cancel... Then multiply and divide all numbers... = 42 donuts

49. 1.4 Units and Unit Conversion Common English System Units Convert 12 gallons to quarts Convert 175 lbs to ounces Convert 10,000 ft to miles

50. 1.4 Units and Unit Conversion Intersystem Conversion Units Convert 4.00 ounces to kilograms Convert 175 lbs to kgs Convert 100 yards to meters

51. 1.4 Units and Unit Conversion Practice Unit Conversions 1.Convert 12 gallons to quarts 2.Convert 10 cm to meters 3.Convert 360 feet to miles 4.Convert 60 miles/hr to m/s 5.Convert 1.8 in 2 to cm 2

52. The speed of light is 299,792,458 m/s  convert this to miles per hour. The speed of sound is 340.29 m/s  convert this to miles per hour 1 mile = 1.6 km; 1 km = 1000 m; 60 s = 1 min; 60 min = 1 hour

53. 1.5 Experimental Quantities Mass - the quantity of matter in an object –not synonymous with weight –standard unit is the gram Weight = mass  acceleration due to gravity Mass must be measured on a balance (not a scale)

54. 1.5 Experimental Quantities Units should be chosen to suit the quantity described –A dump truck is measured in tons –A person is measured in kg or pounds –A paperclip is measured in g or ounces –An atom? For atoms, we use the atomic mass unit (amu) –1 amu = 1.661  10 -24 g

55. 1.5 Experimental Quantities Length - the distance between two points –standard unit is the meter –long distances are measured in km –distances between atoms are measured in nm, 1 nm = 10 -9 m Volume - the space occupied by an object –standard unit is the liter –the liter is the volume occupied by 1000 grams of water at 4 o C –1 mL = 1/1000 L = 1 cm 3

56. 1.5 Experimental Quantities The milliliter (mL) and the cubic centimeter (cm 3 ) are equivalent

57. 1.5 Experimental Quantities Time -metric unit is the second Temperature - the degree of “hotness” of an object 10

58. 1.5 Experimental Quantities 1. Convert 75 o C to o F 2.Convert -10 o F to o C Conversions Between Fahrenheit and Celsius

59. 1.5 Experimental Quantities K = o C + 273 Convert 20ºC to K Kelvin Temperature Scale The Kelvin scale is another temperature scale. It is of particular importance because it is directly related to molecular motion. As molecular speed increases, the Kelvin temperature proportionately increases.

60. 1.5 Experimental Quantities Energy - the ability to do work kinetic energy - the energy of motion potential energy - the energy of position (stored energy) Energy is also categorized by form: light heat electrical mechanical chemical Energy

61. 1.5 Experimental Quantities Characteristics of Energy Energy cannot be created or destroyed Energy may be converted from one form to another Energy conversion always occurs with less than 100% efficiency All chemical reactions involve either a “gain” or “loss” of energy

62. 1.5 Experimental Quantities Units of Energy Basic Units: calorie or joule 1 calorie (cal) = 4.184 joules (J) A kilocalorie (kcal) also known as the large Calorie. This is the same Calorie as food Calories. 1 kcal = 1 Calorie = 1000 calories 1 calorie = the amount of heat energy required to increase the temperature of 1 gram of water 1 o C.

63. 1.5 Experimental Quantities Concentration Concentration: –Either the number of particles of a substance; or the mass of those particles –that are contained in a specified volume Often used to represent the mixtures of different substances –Concentration of oxygen in the air –Pollen counts –Proper dose of an antibiotic –Concentration of Fluoride in toothpaste

64. 1.5 Experimental Quantities Density and Specific Gravity Density –the ratio of mass to volume –an intensive property –use to characterize a substance as each substance has a unique density –Units for density include: g/mL g/cm 3 g/cc 11

65. 1.5 Experimental Quantities liquid mercury brass nut water cork

67. Calculating the Density of a Solid 2.00 cm 3 of aluminum are found to weigh 5.40g. Calculate the density of aluminum in units of g/cm 3. 1.5 Experimental Quantities

68. Density Calculations Air has a density of 0.0013 g/mL. What is the mass of 6.0-L sample of air? Calculate the mass in grams of 10.0 mL if mercury (Hg) if the density of Hg is 13.6 g/mL. Calculate the volume in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams.

69. 1.5 Experimental Quantities Specific Gravity Values of density are often related to a standard Specific gravity - the ratio of the density of the object in question to the density of pure water at 4 o C Specific gravity is a unitless term because the 2 units cancel Often the health industry uses specific gravity to test urine and blood samples