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Chapter 1 Chemical Foundations AP Chemistry

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Objectives Recall units of measure Describe uncertainty in measurement Use scientific notation for numbers Apply significant figure rules

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Units of Measure SI Base Units –Mass Kilogram – kg 1 kilogram is about 2.20 pounds –Length Meter – m 1 meter is about 3 feet

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Units of Measure Time –Second – s –The time needed for a cesium-133 atom to perform 9,192,631,770 complete oscillations. Temperature –Kelvin – K –273 K = 0 degrees C Amount –Mole – mol –1 mol = 6.022x10 23 particles

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Units of Measure Current –Ampre – A Luminous Capacity –Candela – cd First five are the most commonly used in chemistry

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Volume Volume is not an SI Base Unit –Metric system –Powers of 10 1 Liter is 1/1000 of a cubic meter 1 Liter (L) = 1000 cm 3 = 1000 mL

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Cubic Meter Liter Milliliter

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Volume? Uncertainty is in the last digit.

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Accuracy and Precision Accuracy –The nearness of a measurement to its accepted value Precision –The agreement between numerical values –You can be precise without being accurate

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Accurate & Precise NeitherPrecise Loss of accuracy due to systematic errors –Error in same direction every time Random Error give erratic results –Poor technique

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Significant Figures All known digits plus one estimated digit in a measurement

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What is the length? 2 Sig. Fig 1 Known Digit 1 Estimated Digit

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Significant Figures Rule #1 –All Nonzero digits are significant –Ex. 76.44 mL –Ex. 285.85 s

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Significant Figures Rule #2 Captive Zeros Zeros appearing between nonzero digits are significant Ex. 308.2001 g =

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Significant Figures Rule #3 Leading Zeros Zeros appearing in front of nonzero digits are not significant Ex. 0.007036 g Takes care of unit changes

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Significant Figures Rule #4 Ending Zeros Ending zeros are significant if there is a decimal place Ex. 53.00 m Ex. 40000. m Ex. 40000 m 40000. is much more precise than 40000

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What is the length?

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Significant Figures Rule #5 Exact Numbers Exact number have an unlimited number of significant figures Exact numbers are counting numbers or definitions 2 cars or 1000g/1kg

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Significant Figures Rule #6 Scientific Notation All numbers that come before the x10 n are significant –Must be in proper form Ex. 3.33x10 5 Ex. 2.04x10 -4

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Rounding 5 and larger round up 4 and smaller round down Round the following 34.567 to 2 SF = 756.44 to 4 SF = 0.004325 to 3 SF = 3436543 to 2 SF =

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Addition and Subtraction w/ SF Addition and Subtraction –The answer must have the same number of digits to the right of the decimal as there are in the measurement having the fewest digits to the right of the decimal point –Ex. 12.11 m + 15 m = –Number of SFs does not matter!

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Multiplication and Division w/ SF Multiplication and Division –The answer can have no more SFs than are in the measurement with the fewest total SFs –Ex. 55 m / 11.34 s =

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Scientific Notation A method of representing very large or very small numbers M x 10 n –M is a number 1 or larger and less than 10 –n is an integer (positive or negative) –All digits in M are significant (If in proper form)

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Converting to Sci. Notation Move decimal so that M is between 1 and 10 Determine n by counting the number of places the decimal point was moved –Moved to the left, n is positive –Moved to the right, n is negative

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Examples 340,000,000 = 5.04x10 5 = 0.00000300 = 2.212x10 -4 =

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Sci. Notation on Calculators Enter digits in you calculator using the EE key. For TI 83s it is the 2 nd of the comma For TI 30s it is a key Saves key strokes Fewer OOR mistakes 3.4x10 6 = 3.4E6 7.4x10 -5 = 7.4E-5

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Sci. Notatation Math Operations Multiply and Divide –Multiply or divide first number –Add exponents (Multiply) –Subtract exponents (Divide) Addition and Subtraction –Exponents must be the same –Then add or subtract first number –Exponents stay the same

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Calculations 3.0x10 5 + 4.0x10 5 = 4.0x10 3 – 2.0x10 2 = 7.0x10 5 * 2.0x10 4 = 8.0x10 6 / 4.0x10 -3 = _____4.5x10 4 ____ = 6.2x10 6 * 3.1x10 -8

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Objectives Recall metric prefixes Convert numbers from one unit to another Describe different temperature scales Explain density and perform calculations Classify matter into groups

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The Metric System A system based on powers of ten Uses SI Units Allows easy work with both large and small numbers Prefixes tell us which power of 10 we are using

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SI Prefixes (10x larger) Page 9 in your book Tera Giga Mega Kilo Hecto Deca Base 10 12 10 9 10 6 10 3 10 2 10 1 10 0 1000000000000 1000000000 1000000 1000 100 10 1 T G M k h da

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SI Prefixes (10x smaller) Base Deci Centi Milli Micro Nano Pico 10 0 10 -1 10 -2 10 -3 10 -6 10 -9 10 -12 1.1.01.001.000001.000000001.000000000001 d c m μ n p

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Conversions To convert between units set up conversion factors –Ratios of equality

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Convert 67 kg to g

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Convert 450 cL to dL

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Convert 3.4x10 8 ng to kg

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Converting From Metric To English Find ratios that are true –Page 18 has some equivalents

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Convert 763 cm to yd

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Convert 1.2 mi/hr to ft/s

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Convert 3.8 m 2 /hr to cm 2 /s

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Temperature Many different temp. scales All 0 marks based on different ideas 0 ºF Coldest saltwater stays a liquid 0 ºC Normal Freezing Point of water 0 K Molecular motion stops 1 K = 1 ºC = 1.8 ºF

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Temperature Conversion Temp K = 273 + Temp C Temp C = Temp K – 273 0 ºC = 273 K If you need any others look up the equ. T C = (T F – 32)(5/9) T F = T C (9/5) + 32

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Density Ratio of mass to volume Density = __Mass__ Volume Periodic Trend Units –Solids – g/cm 3 –Liquids – g/mL –Gases – g/L

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Density Determination Mass is determined on a balance Volume is measured in two ways –Regular objects can be measured –All objects can use water displacement

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Density Physical Property –Can be used to identify a substance Lead 11.35 Iron 7.87 Magnesium 1.74 Zinc 7.13 Copper 8.96

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Example: A metal cube has sides measuring 3.00 cm. It has a mass of 242.13g. What is the density? What is the metal?

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Density Physical Property –Can be used to identify a substance Lead 11.35 Iron 7.87 Magnesium 1.74 Zinc 7.13 Copper 8.96

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Homework p.33 #'s 33a-f,36a-d,42 47,57,68

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