Chapter 1 Matter, Measurement, and Problem Solving

1. Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro Chapter 1Matter,Measurement, and Problem Solving Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA 2008, Prentice Hall

2. Scientific Method a test of a hypothesis or theory a tentative explanation of a single or small number of natural phenomena a general explanation of natural phenomena the careful noting and recording of natural phenomena a generally observed natural phenomenon

3. Limitations to the Scientific Method • The scientific method is limited • by what can be observed with the five senses • to the present • how, not why a process works • in that it cannot make moral judgments • cannot deal with the unique

4. Chemistry • Chemistry is the science that seeks to understand the behavior of matter by studying the behavior of atoms and molecules • atoms • are submicroscopic particles • are the fundamental building blocks of all matter • molecules • two or more atoms attached together • attachments are called bonds • attachments come in different strengths • molecules come in different shapes and patterns

5. carbon dioxide carbon monoxide • composed of one carbon atom and two oxygen atoms • colorless, odorless gas • incombustible • does not bind to hemoglobin • composed of one carbon atom and one oxygen atom • colorless, odorless gas • burns with a blue flame • binds to hemoglobin Structure Determines Properties • the properties of matter are determined by the atoms and molecules that compose it

6. Classification of Matter • matter is anything that has mass and occupies space • we can classify matter based on whether it’s solid, liquid, or gas

7. Solids • the particles in a solid are packed close together and are fixed in position • though they may vibrate • the close packing of the particles results in solids being incompressible • the inability of the particles to move around results in solids retaining their shape and volume when placed in a new container, and prevents the particles from flowing

8. Crystalline Solids • some solids have their particles arranged in an orderly geometric pattern – we call these crystalline solids • salt and diamonds

9. Amorphous Solids • some solids have their particles randomly distributed without any long-range pattern – we call these amorphous solids • plastic • glass • charcoal

10. Liquids • the particles in a liquid are closely packed, • have some ability to move around • incompressible • take the shape of their container and to flow • don’t have enough freedom to escape or expand to fill the container

11. Gases • particles do not interact with each other • the particles are constantly moving, bumping into each other and the container • there is a lot of empty space between the particles

12. Gases • compressible • expand to fill and take the shape of their container • will flow

13. Classification of Matterby Composition • made of one type of particle • all samples show the same intensive (independent of amount) properties • made of multiple types of particles • samples may show different intensive properties

14. Classification of Pure Substances • made of one type of atom (some elements found as multi-atom molecules in nature) • combine together to make compounds • made of one type of molecule, or array of ions • molecules contain 2 or more different kinds of atoms

15. Classification of Mixtures • made of multiple substances, whose presence can be seen • portions of a sample have different composition and properties • made of multiple substances, but appears to be one substance • all portions of a sample have the same composition and properties

16. Different Physical Property Technique Boiling Point Distillation State of Matter (solid/liquid/gas) Filtration/Decanting Adherence to a Surface Chromatography Volatility Evaporation Density Centrifugation & Decanting Separation of Mixtures • separate mixtures based on different physical properties of the components

17. Distillation

18. Filtration

19. Chromatography • Separation based upon adherence to a surface. • Stationary phase • Moving phase

20. Evaporation Liquid vaporizes leaving less volatile liquid or solid.

21. Centrifugation • Separation based density. • Centrifugal motion causes more dense materials to go to the bottom of the tube

22. Decanting • Separation based state • Carefully pour liquid leaving precipitate • Separation based density. • Carefully pour less dense liquid

23. Properties of Matter • physical properties are the characteristics of matter that can be changed without changing its composition • characteristics that are directly observable • chemical properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy • characteristics that describe the behavior of matter

24. Physical Changes in Matter The boiling of water is a physical change. The water molecules are separated from each other, but their structure and composition do not change.

25. Dissolving of Sugar Subliming of Dry Ice C12H22O11(s) CO2(g) Dry Ice CO2(s) C12H22O11(aq) Common Physical Changes • processes that cause changes in the matter that do not change its composition • state changes • boiling / condensing • melting / freezing • subliming • dissolving

26. Chemical Changes in Matter The rusting of iron is a chemical change. The iron atoms in the nail combine with oxygen atoms from O2 in the air to make a new substance, rust, with a different composition.

27. C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(l) Common Chemical Changes • processes that cause changes in the matter that change its composition • rusting • processes that release lots of energy • burning

28. Energy Changes in Matter • changes in matter, both physical and chemical, result in the matter either gaining or releasing energy • energy is the capacity to do work • work is the action of a force applied across a distance • a force is a push or a pull on an object • electrostatic force is the push or pull on objects that have an electrical charge

29. Energy of Matter - Kinetic • kinetic energy is energy of motion • motion of the atoms, molecules, and subatomic particles • thermal (heat) energy is a form of kinetic energy because it is caused by molecular motion

30. Energy of Matter - Potential • potential energy is energy that is stored in the matter • due to the composition of the matter and its position in the universe • chemical potential energy arises from electrostatic forces between atoms, molecules, and subatomic particles

31. Conversion of Energy • you can interconvert kinetic energy and potential energy • whatever process you do that converts energy from one type or form to another, the total amount of energy remains the same • Law of Conservation of Energy

32. Potential to Kinetic Energy

33. Spontaneous Processes • materials that possess high potential energy are less stable • processes in nature tend to occur on their own when the result is material(s) with lower total potential energy • processes that result in materials with higher total potential energy can occur, but generally will not happen without input of energy from an outside source • when a process results in materials with less potential energy at the end than there was at the beginning, the difference in energy is released into the environment

34. The Standard Units • Scientists have agreed on a set of international standard units for comparing all our measurements called the SI units • Système International = International System

35. Length • SI unit = meter • About a yard • Commonly use centimeters (cm) • 1 m = 100 cm • 1 cm = 0.01 m = 10 mm • 1 inch = 2.54 cm (exactly)

36. Mass • Measure of the amount of matter present in an object • weight measures the gravitational pull on an object, which depends on its mass • SI unit = kilogram (kg) • about 2 lbs. 3 oz. • Commonly measure mass in grams (g) or milligrams (mg) • 1 kg = 2.2046 pounds, 1 lbs. = 453.59 g • 1 kg = 1000 g = 103 g • 1 g = 1000 mg = 103 mg • 1 g = 0.001 kg = 10-3 kg • 1 mg = 0.001 g = 10-3 g

37. Time • measure of the duration of an event • SI units = second (s)

38. Temperature Scales • Fahrenheit Scale, °F • used in the U.S. • Celsius Scale, °C • used in all other countries • Kelvin Scale, K • The SI unit for temperature

39. Prefix Multipliers in the SI System

40. Volume • Derived unit • any length unit cubed • Measure of the amount of space occupied • SI unit = cubic meter (m3) • Commonly measure solid volume in cubic centimeters (cm3) • 1 m3 = 106 cm3 • 1 cm3 = 10-6 m3 = 0.000001 m3 • Commonly measure liquid or gas volume in milliliters (mL) • 1 L is slightly larger than 1 quart • 1 L = 1 dm3 = 1000 mL = 103 mL • 1 mL = 0.001 L = 10-3 L • 1 mL = 1 cm3

41. Mass & Volume • mass and volume are extensive properties • the value depends on the quantity of matter • extensive properties cannot be used to identify what type of matter something is • if you are given a large glass containing 100 g of a clear, colorless liquid and a small glass containing 25 g of a clear, colorless liquid - are both liquids the same stuff?

42. Mass vs. Volume of Brass

43. What Is a Measurement? • quantitative observation • every measurement has a number and a unit • every digit written is certain, except the last one which is estimated

45. Estimation in Weighing • What is the uncertainty in this reading?

46. Uncertainty in Measured Numbers uncertainty comes from: • limitations of the instruments used for comparison, • the experimental design, • the experimenter, • nature’s random behavior

47. Precision and Accuracy • accuracy is an indication of how close a measurement comes to the actual value of the quantity • precision is an indication of how reproducible a measurement is

48. Accuracy vs. Precision

49. Precision • imprecision in measurements is caused by random errors • errors that result from random fluctuations • no specific cause, therefore cannot be corrected • we determine the precision of a set of measurements by evaluating how far they are from the actual value and each other • even though every measurement has some random error, with enough measurements these errors should average out

50. Accuracy • inaccuracy in measurement caused by systematic errors • errors caused by limitations in the instruments or techniques or experimental design • can be reduced by using more accurate instruments, or better technique or experimental design • we determine the accuracy of a measurement by evaluating how far it is from the actual value • systematic errors do not average out with repeated measurements because they consistently cause the measurement to be either too high or too low