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Mathematical Relationships in Chemistry CP Chemistry.

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Presentation on theme: "Mathematical Relationships in Chemistry CP Chemistry."— Presentation transcript:

1 Mathematical Relationships in Chemistry CP Chemistry

2 What You’ll Learn in this Unit Measurement Dimensional Analysis Scientific Notation Significant Figures Error Density

3 Measurement Every measurement has two parts Number with the correct sig - figs Scale (unit) SI system (le Systeme International) based on the metric system Prefix + base unit Prefix tells you the power of 10 to multiply by - decimal system -easy conversions

4 Measurement We use the SI (not the Sports Illustrated) It is called the Systeme Internationale. COMMON SI UNITS Symbol Unit Name Quantity Definition mmeterlengthbase unit kgkilogrammassbase unit ssecondtimebase unit Kkelvin temperature base unit °C degree Celsiusdegree Celsius** temperature m 3 cubic metervolumem3m3 Lliterliter**volume dm 3 = m 3 Nnewtonforcekg·m/s 2 JjouleenergyN·m WwattpowerJ/s PapascalpressureN/m 2 Hzhertzfrequency1/s

5 Metric SI Units Mass - kilogram (kg) Length- meter (m) Time - second (s) Temperature - Kelvin (K) Electric current - ampere (amp, A) Amount of substance - mole (mol)

6 PrefixSymbolMagnitude GigaG1,000,000,000 (10 9 ) MegaM1,000,000 (10 6 ) KiloK1,000(10 3 ) HectoH100 (10 2 ) Decada10 (10 1 ) base unitg, L, m1 (10 0 ) Decid0.1 (10 -1 ) Centic0.01 (10 -2 ) Millim0.001 (10 -3 ) Microμ (10 -6 ) Nanon (10 -9 ) Picop ( )

7 Dimensional Analysis Using Units to solve problems

8 Dimensional Analysis Use conversion factors to change the units Conversion factors = 1 1 foot = 12 inches (equivalence statement) 12 in = 1 = 1 ft. 1 ft. 12 in 2 conversion factors multiply by the one that will give you the correct units in your answer.

9 Example Problem There are 2.2 lb in 1 kg If you weigh 158 lbs, how many kg do you weigh?

10 Example Problems 11 yards = 2 rod 40 rods = 1 furlong 8 furlongs = 1 mile 1 mile = 1.6 km The Kentucky Derby race is 1.25 miles. How long is the race in rods, furlongs, meters, and kilometers?

11 Example Problems Convert: –475m to km –35daL to mm

12 Scientific Notation 100 = 1.0 x = 1.0 x This provides a way to show significant figures.

13 TOO QUICK FOR YOU! So here are the rules.. slowly! 1.Place decimal point after 1 st real non-zero integer. (ex) 1.0 NOT Raise 10 to the exponential which equals the number of places you moved.

14 Sample Problems

15 Answers x X x X X X 10 -6

16 Uncertainty Basis for significant figures All measurements are uncertain to some degree Precision- how repeatable Accuracy- how correct - closeness to true value. Random error - equal chance of being high or low- addressed by averaging measurements - expected

17 Uncertainty Systematic error- same direction each time Want to avoid this Better precision implies better accuracy You can have precision without accuracy, and vice versa

18 Precision vs. Accuracy Precision- the degree of agreement among several measurements of the same quantity. Accuracy- the agreement of a particular value with the true value


20 Significant Figures Meaningful digits in a MEASUREMENT The number of significant figures in your measurement will tell the reader how exact the instrumentation is If it is measured or estimated, it has sig figs. If not, it is exact.

21 Significant Figures All numbers except zero are significant. Some zeros are, some aren’t

22 Which Zeros Count? In between other sig figs does Before the first number doesn’t After the last number counts if it is after the decimal point the decimal point is written in sig figs sig figs

23 Doing the Math Multiplication and division, same number of sig figs in answer as the least in the problem Addition and subtraction, same number of decimal places in answer as least in problem.

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

25 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

26 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 scale

27 Density  Which is heavier- lead or feathers?  It depends upon the amount of the material  A truckload of feathers is heavier than a small pellet of lead  The relationship here is between mass and volume- called Density

28 Density  The formula for density is: mass volume Common units are g/mL, or possibly g/cm 3, (or g/L for gas) Density =

29 Density Useful for identifying a compound Useful for predicting weight An intrinsic property- does not depend on what the material is Intensive Property Density is a physical property, and does not depend upon sample size

30 Things related to density  Corn oil density – 0.921g/mL  Water density – 1.000g/mL  What happens when corn oil and water are mixed?  Why?  Will lead float in water?

31 Example Problem An empty container weighs g. Filled with carbon tetrachloride (density=1.53 g/cm 3 ), the full container weighs g. What is the volume of the container?

32 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

33 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?

34 Specific Gravity  A comparison of the density of an object to a reference standard (which is usually water) at the same temperature  Water density at 4 o C = 1 g/cm 3

35 Specific Gravity Formula D of substance (g/cm 3 ) D of water (g/cm 3 ) Note there are no units left, since they cancel each other Measured with a hydrometer Uses? Tests urine, antifreeze, battery SG =

36 Temperature A measure of the average kinetic energy Different temperature scales, all are talking about the same height of mercury. In lab take the reading in ºC then convert to our SI unit Kelvin ºC = K

37 Temperature  Heat moves from warmer object to the cooler object  Glass of iced tea gets colder?  Remember that most substances expand with a temp. increase?  Basis for thermometers

38 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

39 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

40 Temperature scales  Water freezes at 273 K  Water boils at 373 K  0 K is called absolute zero, and equals –273 o C

41 100ºC = 212ºF 0ºC = 32ºF 100ºC = 180ºF 1ºC = (180/100)ºF 1ºC = 9/5ºF For Calculations: °F = 9/5 (°C) + 32 °C = 5/9 (°F - 32)

42 Error Calculations Error = Experimental value - accepted value % error = [error] accepted value X 100

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