Chemistry. The study of matter and the changes it undergoes

Transcription

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

2 5 Major Areas of Chemistry Analytical Chemistry- composition of substances. Inorganic Chemistry- substances without carbon Organic Chemistry- most substances containing carbon Biochemistry- Chemistry of living things Physical Chemistry- describes the behavior of chemicals (ex. stretching)

3 1.1 Chemistry s impact on society: Health & Medicine Biotechnology Energy & Environment Materials & Technology Agriculture- world s food supply The Environment- both risks and benefits involved in discoveries Astronomy and Space Exploration

4 1.3 The Scientific Method A logical approach to solving problems or answering questions. Starts with observation- noting and recording facts hypothesis- an educated guess as to the cause of the problem, or a proposed explanation

5 Scientific Method controlled experiment- designed to test the hypothesis only two possible answers hypothesis is right hypothesis is wrong Generates data & observations from experiments. Modify hypothesis - repeat the cycle

6 Observations Hypothesis Experiment Cycle repeats many times. The hypothesis gets more and more certain. Becomes a theory A thoroughly tested model that explains why things behave a certain way.

7 Data Collection: Qualitative vs. Quantitative Examples??

8 1.4 What is Matter? Matter is anything that takes up space and has mass. Everything, but energy! Mass- amount of material or stuff in an object Weight is due to gravity, and changes from location to location; mass is always constant.

9 Types of Matter Substance- a particular kind of matter - pure; is uniform (all the same) and has a definite composition (examples are elements & compounds) water; gold Mixture- more than one kind of matter; has a variable composition

10 Substances Elements- simplest kind of matter cannot be broken down any simpler all one kind of atom. Compounds are substances that can be broken down only by chemical methods When broken down, the pieces have completely different properties than the original compound. Made of two or more atoms, chemically combined (not physical blend!)

11 10 Most Common Elements

12 Mixtures Physical blend of at least two substances; variable composition Heterogeneous- mixture is not uniform in composition Chocolate chip cookie, gravel, soil. Homogeneous- same composition throughout; called solutions Kool-aid, air, salt water Every part keeps its own properties.

13 Compound or Mixture Compound Mixture Made of one kind of material Made by a chemical change Definite composition Made of more than one kind of material Made by a physical change Variable composition

14 Classification of Matter

15 Which is it? Compound Element Mixture

16 1.5 States of Matter Solid- matter that can not flow, vibrational movement, low kinetic energy Liquid- flows, more kinetic energy than solid, Gas- flows, independent molecular motion, fills container, highest kinetic energy Vapor- a substance that is currently a gas, but normally is a liquid or solid at room temperature. (water vapor)

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19 Solid Definite Volume? YES States of Matter Definite Shape? YES Temp. increase Small variation Compressible? NO Liquid YES NO Small variation. NO Gas NO NO Large Variation YES

20 Freeze Melt Condense Evaporate or Boil Solid Liquid Gas

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22 1.6 Properties Words that describe matter (adjectives) Physical Properties- a property that can be observed and measured without changing the composition. Examples- color, hardness, m.p., b.p. Chemical Properties- a property that can only be observed by changing the composition of the material. Examples-volatile, flammable

23 Types of Properties: Extensive Properties: Dependent on quantity of matter ex: mass, volume Intensive Properties: Independent of quantity ex: density, boiling point

24 Physical Changes A change that changes appearances, without changing the composition. Ex. Boil, melt, cut, bend, split, crack Boiled water is still water. Chemical changes - a change where a new form of matter is formed. Ex. Rust, burn, decompose, ferment

25 Filtration: Physical Separation

26 Distillation: Physical Separation

27 1.7 Measurement: International System of Units The number is only part of the answer; it also need UNITS The standards of measurement used in science are those of the Metric System

28 International System of Units Metric system is now revised as the International System of Units (SI), as of 1960 Simplicity and based multiples of base units (Know them p.17 Table 1.3)

29 International System of Units Sometimes, non-si units are used Liter, Celsius, calorie Some are derived units Made by joining other units Speed (miles/hour) Density (grams/ml)

30 Most Commonly Used Prefixes

31 Volume The space occupied by any sample of matter Calculated for a solid by multiplying the length x width x height SI unit = cubic meter (m 3 ) Everyday unit = Liter (L), which is non-si

32 Solid Volume Calculations 1 cm 3 = 1 ml 1 dm 3 = 1000 ml = 1 L 1 m 3 = 1,000,000 ml = 1,000 L

33 Volume Measuring Instruments Graduated cylinders Graduated Pipet Buret Volumetric Flask Syringe

34 Volume from Water Displacement

35 Units of Mass Mass is a measure of the quantity of matter Weight is a force that measures the pull by gravity- it changes with location Mass is constant, regardless of location

36 Working with Mass The SI unit of mass is the kilogram (kg), even though a more convenient unit is the gram Measuring instrument is the balance or scale

37 Density The formula for density is: Density = mass volume Common units are g/ml, or possibly g/cm 3, (or g/l for gas) Density is an intensive, physical property

38 Density and Temperature What happens to density as the temperature increases? Mass remains the same Most substances increase in volume as temperature increases Thus, density generally decreases as the temperature increases

39 Density and water Water is an important exception Over certain temperatures, the volume of water increases as the temperature decreases Does ice float in liquid water? Why?

40 Temperature Heat moves from warmer object to the cooler object Remember that most substances expand with a temp. increase Basis for thermometers

41 Temperature scales Celsius scale- named after a Swedish astronomer Uses the freezing point(0 o C) and boiling point (100 o C) of water as references Divided into 100 equal intervals, or degrees Celsius

42 Temperature scales Kelvin scale (or absolute scale) Named after Lord Kelvin K = o C A change of one degree Kelvin is the same as a change of one degree Celsius No degree sign is used

43 3 most common temp. scales 0 K is called absolute zero, and equals C

44 Conversion Formulas K = 0 C F = 9/5( 0 C) C = 5/9 ( 0 F 32)

45 1.8 Handling Numbers How do you know what number to round your calculation answers to? Significant figures: Determining which numbers are meaningful in a measurement or calculated quantity.

46 Working with Scientific Notation Regardless of magnitude: all numbers can be expressed in formula N x 10 n N = number between 1 and 9.9 N = positive or negative whole # Ex: Express in scientific notation Answer: 4.56 X 10-7

47 Working with Scientific Notation Do you know how to use your calculator when you have numbers in scientific notation????

48 Significant Figures Significant figures in a measurement include all of the digits that are known, plus a last digit that is estimated or uncertain. **This is what you did when you read the volume from the glassware in lab.

49 Significant Figures Rules: (page 24: Know 1-5) 1.) Any non-zero digit is significant Ex: = 5 sig. figs 2.) Zeros between non-zero digits are significant Ex: = 8 sig. figs 3.) Zeros to the left of the first non-zero digit are not significant Ex: = 6 sig. figs

50 Significant Figures 4.) All zeros to the right of the decimal point are significant if they follow a non-zero number Ex: = 7 sig. figs = 3 sig. figs 5.) # s without decimals present ambiguous info. Always use scientific notation to clear up problems. Ex: 56,700 =? sig. figs x 10 4 = 4 sig. figs

51 Significant Figures Try These: How many sig. figs? A.) 1,245 B.) 1,245, C.) D.) E.) x 10-5

52 Sig. fig. calculations Addition and Subtraction The answer should be rounded to the same number of decimal places as the least number in the problem

53 Addition & Subtraction Examples: 1.) 4.56 cm cm= Answer: 7.7 cm 2.) L L = Answer: L 3.) 450 g g= Answer: 451 g

54 Sig. Fig. calculations Multiplication and Division Round the answer to the same number of significant figures as the least number in the measurement

55 Multiply & Divide Examples: 1.) 451 x 3.2 = Answer: 1,400 or 1.4 x ) /5.60 = Answer: or 6.16 x ) x 12 = Answer: 2.4

56 Significant Figures **Exact numbers obtained from definition or by counting of objects can be considered to have an infinite # of sig. figs. They are not considered in the calculation. Only use measurements!! Ex: 12 eggs in a dozen

57 Uncertainty in Measurements Need to make reliable measurements in the lab Accuracy how close a measurement is to the true value Precision how close the measurements are to each other (reproducibility) **Do multiple trials for experiments!!

58 Uncertainty in Measurements Accepted value correct value based on reliable references Experimental value the value measured in the lab Error the difference between the accepted and experimental values

59 Uncertainty in Measurements Error = accepted experimental Can be positive or negative Percent error = the absolute value of the error divided by the accepted value, times 100% % error = error accepted value x 100%

60 % Error Can you think of an easier way to calculate this???

61 1.9 Dimensional Analysis A way to analyze and solve problems by using units (or dimensions) of the measurement Based on conversion factors Conversion factors are fractions that are equal to one. Both the top and bottom measurements are identical; they just use different units. Examples: 1ft/12 in 5,280 ft/1 mi

62 Dimensional Analysis Give me some more examples!

63 Dimensional Analysis Example Problems A ruler is 12.0 inches long. How long is it in cm? ( 1 inch = 2.54 cm) in meters? A race is 10.0 km long. How far is this in miles? Pikes peak is 14,110 ft. above sea level. What is this in meters?

64 Dimensional Analysis Another measuring system has different units of measure: 6 ft = 1 fathom 100 fathoms = 1 cable length 10 cable lengths = 1 nautical mile 3 nautical miles = 1 league Problem: Jules Verne wrote a book 20,000 leagues under the sea. How far is this in feet?

65 Problem solving 1. ANALYZE a) Identify the unknown Both in words and what units it will be measured in. Write it down! May need to read the question several times.

66 Problem Solving b) Identify what is given (the known ) Write it down! Unnecessary information may also be given

67 Problem solving c) Plan a solution Break it down into steps. Look up needed information: *Tables *Formulas *Constants, or conversion factors *Choose an equation

68 Problem solving 2. CALCULATE doing the arithmetic use a calculator

69 Problem Solving 3. EVALUATE Round off to proper # of sig. figs. Proper units? Need Scientific Notation? Check your work! Reread the question, did you answer it? Is it reasonable? Estimate an approximate answer

70 Converting Complex Units Units expressed as a ratio or raised to a power speed is: miles/hour gas mileage is: miles/gallon density is: g/cm 3 Volume is: cm 3, dm 3, m 3

71 Examples: The density of silver is 10.5 g/cm 3. Convert the density to kg/m 3. Answer: 1.05 x 10 4 kg/m 3 The density of the lightest metal, lithium, is 5.34 x 10 2 kg/m 3. Convert the density to g/cm 3. Answer: g/cm 3

72 Lastly. What makes you perfect or close to it? Practice Practice & More Practice!!!