Introduction to AP Chemistry: Matter and Measurement

Transcription

1 Introduction to AP Chemistry: and Mr. Matthew Totaro Legacy High School AP Chemistry

2 Chemistry The study of matter and the changes it undergoes.

3 Scientific Method A systematic approach to solving problems. Procedure designed to test an idea Tentative explanation of a single or small number of observations Careful noting and recording of natural phenomena Generally repeating occurrence in nature General explanation of natural phenomena

4 Anything that has mass and takes up space.

5 Atoms are the building blocks of matter.

6 Atoms are the building blocks of matter. Each element is made of the same kind of atom.

7 Atoms are the building blocks of matter. Each element is made of the same kind of atom. A compound is made of two or more different kinds of elements.

8 States of Temperature Dependence of the Phases

9 Classification of

10 Classification of

11 Classification of

12 Classification of

13 Classification of

14 Classification of

15 Classification of

16 Classification of

17 Classification of

18 Classification of

19 Classification of

20 Properties and Changes of

21 Properties of Physical Properties: Can be observed without changing a substance into another substance. Boiling point, density, mass, volume, etc. Chemical Properties: Can only be observed when a substance is changed into another substance. Flammability, corrosiveness, reactivity with acid, etc.

22 Some Physical Properties Mass Volume Density Solid Liquid Gas Melting point Boiling point Volatility Taste Odor Color Texture Shape Solubility Electrical Thermal Magnetism conductance conductance Malleability Ductility Specific heat capacity

23 Some Physical Properties of Iron Iron is a silvery solid at room temperature with a metallic taste and smooth texture. Iron melts at 1538 C and boils at 4428 C. Iron s density is 7.87 g/cm 3. Iron can be magnetized. Elemental Iron Iron conducts electricity, but not as well as most other common metals. 23

24 Some Chemical Properties Acidity Causticity Reactivity Inertness (In)Flammability Oxidizing ability Basicity (aka alkalinity) Corrosiveness Stability Explosiveness Combustibility Reducing ability 24

25 Some Chemical Properties of Iron Iron is easily oxidized in moist air to form rust. When iron is added to hydrochloric acid, it produces a solution of ferric chloride and hydrogen gas. Iron is more reactive than silver, but less reactive than magnesium. 25

26 Properties of Intensive Properties: Independent of the amount of the substance that is present. Density, boiling point, color, etc. Extensive Properties: Dependent upon the amount of the substance present. Mass, volume, energy, etc.

27 Changes of Physical Changes: Changes in matter that do not change the composition of a substance. Changes of state, temperature, volume, etc. Chemical Changes: Changes that result in new substances. Combustion, oxidation, decomposition, etc.

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

29 Chemical Changes in The rusting of iron is a chemical change. The iron atoms in the nail combine with oxygen atoms from O 2 in the air to make a new substance, rust, with a different composition.

30 Chemical Reactions In the course of a chemical reaction, the reacting substances are converted to new substances.

31 Chemical Reactions

32 Compounds Compounds can be broken down into more elemental particles.

33 Electrolysis of Water

34 Separation of Mixtures

35 Distillation Separates homogeneous mixture on the basis of differences in boiling point.

36 Distillation

37 Filtration Separates solid substances from liquids and solutions.

38 Chromatography Separates substances on the basis of differences in solubility in a solvent.

39 Units of

40 SI Units Système International d Unités Uses a different base unit for each quantity

41 Metric System Prefixes convert the base units into units that are appropriate for the item being measured.

42 Volume The most commonly used metric units for volume are the liter (L) and the milliliter (ml). A liter is a cube 1 dm long on each side. A milliliter is a cube 1 cm long on each side.

43 Temperature A measure of the average kinetic energy of the particles in a sample.

44 Temperature In scientific measurements, the Celsius and Kelvin scales are most often used. The Celsius scale is based on the properties of water. 0 C is the freezing point of water. 100 C is the boiling point of water.

45 Temperature The Kelvin is the SI unit of temperature. It is based on the properties of gases. There are no negative Kelvin temperatures. K = C

46 Temperature The Fahrenheit scale is not used in scientific measurements. F = 9/5( C) + 32 C = 5/9( F 32)

47 Practice Convert 0.0 F into Kelvin Sort information Strategize Given: Find: Concept Plan: 0.0 F Kelvin F C K Follow the concept plan to solve the problem Sig. figs. and round Check Equations: Solution: Round: Check: K = 255 K Because kelvin temperatures are always positive and generally between 250 and 300, the answer makes sense

48 Density Intensive physical property of a substance d = m V

49 Example: Decide if a ring with a mass of 3.15 g that displaces cm 3 of water is platinum Write down the given quantities and the quantity you want to find Given: Find: mass = 3.15 g volume = cm 3 density, g/cm 3 Find the equation that relates the given quantity to the quantity you want to find Solve the equation for the quantity you want to find, check the units are correct, then substitute and compute Equation: Compare to accepted value of the intensive property Density of platinum = 21.4 g/cm 3 therefore not platinum

50 and Significant Figures

51 What Is a? Quantitative observation Comparison to an agreed standard Every measurement has a number and a unit

52 A The unit tells you what standard you are comparing your object to The number tells you 1. what multiple of the standard the object measures 2. the uncertainty in the measurement Scientific measurements are reported so that every digit written is certain, except the last one, which is estimated

53 Uncertainty in s Different measuring devices have different uses and different degrees of accuracy.

54 Estimating the Last Digit For instruments marked with a scale, you get the last digit by estimating between the marks if possible Mentally divide the space into ten equal spaces, then estimate how many spaces over the indicator the mark is

55

56 Uncertainty in

57 Significant Figures The term significant figures refers to digits that were measured. When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

58 Significant Figures 1. All nonzero digits are significant. 2. Zeroes between two significant figures are themselves significant. (trapped zeros) 3. Zeroes at the beginning of a number are never significant. (leading zeros) 4. Zeroes at the end of a number are significant if a decimal point is written in the number. (trailing zeros) 5. Exact numbers and constants do not have significant figures.

59 Example: Determining the Number of Significant Figures in a Number How many significant figures are in each of the following? m km 10 dm = 1 m s mm 10,000. m 4 sig. figs.; the digits 4 and 5, and the trailing 0 5 sig. figs.; the digits 5 and 3, and the interior 0 s infinite number of sig. figs., exact numbers 4 sig. figs.; the digit 1, and the trailing 0 s 1 sig. figs.; the digit 2, not the leading 0 s Tricky, the dot makes all 5 digits significant

60 Significant Figures When addition or subtraction is performed, answers are rounded to the least significant decimal place. When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

61 Multiplying or Dividing with Sig Figs When multiplying or dividing measurements with significant figures, the result has the same number of significant figures as the measurement with the lowest number of significant figures = = 45 3 sig. figs. 5 sig. figs. 2 sig. figs. 2 sig. figs = = sig. figs. 3 sig. figs. 3 sig. figs.

62 Addition and Subtraction with Significant Figures When adding or subtracting measurements with significant figures, the result has the same number of decimal places as the measurement with the lowest number of decimal places = = 5.7

63 Both Multiplication/Division and Addition/Subtraction with Significant Figures When doing different kinds of operations with measurements with significant figures, do whatever is in parentheses first, evaluate the significant figures in the intermediate answer, then do the remaining steps ( ) = 2 dp 1 dp = 12 4 sf 1 dp & 2 sf 2 sf

64 Example: Perform the Following Calculations to the Correct Number of Significant Figures

65 Example: Perform the Following Calculations to the Correct Number of Significant Figures

66 Accuracy versus Precision Accuracy refers to the proximity of a measurement to the true value of a quantity. Precision refers to the proximity of several measurements to each other.

67 Accuracy vs. Precision Suppose three students are asked to determine the mass of an object whose known mass is g The results they report are as follows Looking at the graph of the results shows that Student A is neither accurate nor precise, Student B is inaccurate, but is precise, and Student C is both accurate and precise.

68 Measuring Accuracy Accuracy is measured by calculating Percent Error

69 Example: A student was asked to determine the density of a sample of nickel metal. When she finished, she reported the density of nickel as 5.59 g/ml. The density of nickel was 6.44 g/ml. What is her percent error?