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**BASIC CHEMISTRY CHM 138 CHAPTER 1: UNITS OF MEASUREMENTS**

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**Chemistry is the study of matter and the changes it undergoes**

Matter is anything that occupies space and has mass. A substance is a form of matter that has a definite composition and distinct properties. liquid nitrogen gold ingots silicon crystals

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**DIMENSIONS Defined as:**

Any physical quantities that can be measured by measuring devices. There are 2 types of dimension: 1.Basic dimension The simplest and the most basic physical quantities There are 7 basic dimensions: - Length (L) - mass (m) - time (t) - temperature (T) - amount of substances (n) - luminous intensity (I) - electrical current

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**Combination of Basic dimensions**

2. Derived dimensions The compound physical quantities that are obtained by combining basic physical by multiplying / dividing the basic dimensions. Derived dimensions Combination of Basic dimensions Area Length x length Volume Length x length x length Speed Length / time Force (mass x length) / (time x time) Density Mass / (length x length x length) Pressure Mass / (time x time)

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**International System of Units (SI)**

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**Meter per second squared m.s-2 Pressure Newton per meter squared, **

S.I system of units Dimension Symbol of dimension Name of unit Symbol of unit Definition of unit Basic SI Units length L meter m mass kilogram Kg time t second s temperature T Kelvin, degree celcius K, °C Amount of substance n mole mol Derived SI Units Energy E Joule J / Nm Kg.m2.s-2 Force F Newton N Kg.m.s-2 Power P watt W /J/s Kg.m2.s-3 Velocity V Meter per second m.s-1 Acceleration a Meter per second squared m.s-2 Pressure Newton per meter squared, Pascal N.m-2 Pa Density ƿ Kilogram per cubic meter Kg.m-3

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**American Engineering system of units Dimension Symbol of dimension **

Name of unit Symbol of unit Definition of unit Basic SI Units length L Feet ft mass m Pound mass Ibm force F Pound force Ibf time t Second S temperature T Degree Rankie, Degree Fahrenheit °R °F Amount of substance n mole Ib mol Derived SI Units Energy E Foot pound force Ft.Ibf Force Power P Horsepower hp Velocity V feet per second ft.s-1 Acceleration a feet per second squared ft.s-2 Pressure Pound force per second inch Psi Density ƿ Pound mass per cubic foot Ibm.ft-3

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**Quantity Equivalent Values**

Mass 1 kg = 1000 g = metric ton = Ibm = oz 1 Ibm = 16 oz = 5 x 10-4 ton = g = kg Length 1 m = 100 cm = 1000 mm = 106 microns (µm) = 1010 angstroms (A) = in. = ft = yd = mile 1 ft = 12 in. = 1/3 yd = m = cm Volume 1 m3 = 1000 L = 106 cm3 = 106 mL = ft3 = imperial gallons = gal = qt 1 ft3 = 1728 in.3 = gal = m3 = L = 28,317 cm3 Force 1 N = 1 kg.m/s2 = 105 dynes = 105 g.cm/s2 = Ibf 1 Ibf = Ibm.ft/s2 = N = x 105 dynes Pressure 1 atm = x 105 N/m2 (Pa) = kPa = bar = x 106 dynes/cm2 = 760 mm Hg = Ibf/in.2 (psi) = 33.9 ft = in. Hg Energy 1 J = 1 N.m = 107 ergs = 107 dyne.cm = x 10-7 kW.h = cal = ft-Ibf = x 10-4 Btu power 1 W = 1 J/s = cal/s = ft.Ibf/s = x 10-4 Btu/s = x 10-3 hp

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1 year = 365 days 1 day = 24 hours 1 hour = 60 minutes 1 minute = 60 seconds 30 days = 1 month 1 year = 12 month 1 qt = L 1 gal = L 1 L = fl oz 1 fl oz = ml 1 L = qt 1 in = 2.54 cm 1 m = in 1 mile = km 1 mile = 5280 feet 1 mile = 1760 yards 3 feet = 1 yard 12 inches = 1 feet 1 oz = g 1 lb = g 1 kg = lb 1 g = grains

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**MASS AND WEIGHT Matter - anything that occupies space and has mass.**

mass – measure of the quantity of matter SI unit of mass is the kilogram (kg) 1 kg = 1000 g = 1 x 103 g weight – force that gravity exerts on an object weight = c x mass A 1 kg bar will weigh 1 kg on earth 0.1 kg on moon on earth, c = 1.0 on moon, c ~ 0.1

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**VOLUME Volume – SI derived unit for volume is cubic meter (m3)**

1 cm3 = (1 x 10-2 m)3 = 1 x 10-6 m3 1 dm3 = (1 x 10-1 m)3 = 1 x 10-3 m3 1 L = 1000 mL = 1000 cm3 = 1 dm3 1 mL = 1 cm3

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**Density – SI derived unit for density is kg/m3**

1 g/cm3 = 1 g/mL = 1000 kg/m3 d = m V density = mass volume EXAMPLE: A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass? d = m V m = d x V = 21.5 g/cm3 x 4.49 cm3 = 96.5 g

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Ex. 1.1

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Ex. 1.2

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**A Comparison of Temperature Scales**

Three temperature scales: 0F – degrees Fahrenheit 0C – degrees Celcius K - Kelvin K = oC 9 oF oF = x oC oF 5 oC 5 oC oC = (oF – 32oF) x 9 oF

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Example: Convert F to degrees Celsius. 0C = x (0F – 32) 9 5 0C = x (172.9 – 32) = 78.3 9 5

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**Scientific Notation The number of atoms in 12 g of carbon:**

602,200,000,000,000,000,000,000 6.022 x 1023 The mass of a single carbon atom in grams: 1.99 x 10-23 N x 10n N is a number between 1 and 10 n is a positive or negative integer

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move decimal left move decimal right n > 0 n < 0 = x 102 = 7.72 x 10-6 Addition or Subtraction Write each quantity with the same exponent n Combine N1 and N2 The exponent, n, remains the same 4.31 x x 103 = 4.31 x x 104 = 4.70 x 104

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**(a x 10m) x (b x 10n) = (axb) x 10m+n**

Multiplication Multiply N1 and N2 Add exponents n1 and n2 (4.0 x 10-5) x (7.0 x 103) = ? = (4.0 x 7.0) x (10-5+3) = 28 x 10-2 = 2.8 x 10-1 (a x 10m) x (b x 10n) = (axb) x 10m+n Division Divide N1 and N2 Subtract exponents n1 and n2 8.5 x 104 ÷ 5.0 x 109 = ? = (8.5 ÷ 5.0) x 104-9 = 1.7 x 10-5 (a x 10m) ÷ (b x 10n) = (a ÷ b) x 10m-n

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Significant Figures - The meaningful digits in a measured or calculated quantity. RULES: Any digit that is not zero is significant 1.234 kg significant figures Zeros between nonzero digits are significant 606 m significant figures Zeros to the left of the first nonzero digit are not significant 0.08 L significant figure If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg significant figures

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If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant g 3 significant figures Numbers that do not contain decimal points, zeros after the last nonzero digit may or may not be significant. 400 cm or 2 or 3 significant figures 4 x significant figures 4.0 x significant figures

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**How many significant figures are in each of the following measurements?**

24 mL 2 significant figures 3001 g 4 significant figures m3 3 significant figures 6.4 x 104 molecules 2 significant figures 560 kg 2 significant figures

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**Significant Figures Addition or Subtraction**

The answer cannot have more digits to the right of the decimal point than any of the original numbers. 89.332 1.1 + 90.432 one significant figure after decimal point round off to 90.4 3.70 0.7867 two significant figures after decimal point round off to 0.79

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**Significant Figures Multiplication or Division**

The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x = = 16.5 3 sig figs round to 3 sig figs 6.8 ÷ = = 0.061 2 sig figs round to 2 sig figs

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**Significant Figures Exact Numbers**

Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures The average of three measured lengths; 6.64, 6.68 and 6.70? 3 = = 6.67 = 7 Because 3 is an exact number

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**Dimensional Analysis Method of Solving Problems**

Determine which unit conversion factor(s) are needed Carry units through calculation If all units cancel except for the desired unit(s), then the problem was solved correctly. given quantity x conversion factor = desired quantity desired unit given unit given unit x = desired unit

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**Dimensional Analysis Method of Solving Problems**

How many mL are in 1.63 L? Conversion Unit 1 L = 1000 mL 1L 1000 mL 1.63 L x = 1630 mL

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**The speed of sound in air is about 343 m/s**

The speed of sound in air is about 343 m/s. What is this speed in miles per hour? conversion units meters to miles seconds to hours 1 mi = 1609 m 1 min = 60 s 1 hour = 60 min 343 m s x 1 mi 1609 m 60 s 1 min x 60 min 1 hour x = 767 mi hour

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Chapter 3 Scientific Measurement Measurement In chemistry, #’s are either very small or very large 1 gram of hydrogen = 602,000,000,000,000,000,000,000.

Chapter 3 Scientific Measurement Measurement In chemistry, #’s are either very small or very large 1 gram of hydrogen = 602,000,000,000,000,000,000,000.

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