SURVEY OF CHEMISTRY I CHEM 1151 CHAPTER 1 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

CHAPTER 1 MEASUREMENT

- Is the determination of the dimensions, capacity, quantity, or extent of something - Is a quantitative observation and consists of two parts: a number and a scale (called a unit) - Is the tool chemists use most Examples mass, volume, temperature, pressure, length, height, time MEASUREMENT

SIGNIFICANT FIGURES Precision - Provides information on how closely individual (repeated) measurements agree with one another Accuracy - Refers to how closely individual measurements agree with the true value (correct value) Precise measurements may NOT be accurate

SIGNIFICANT FIGURES Exact Numbers - Values with no uncertainties - There are no uncertainties when counting objects or people (24 students, 4 chairs, 10 pencils) - There are no uncertainties in simple fractions (1/4, 1/7, 4/7, 4/5) Inexact Numbers - Associated with uncertainties - Measurement has uncertainties (errors) associated with it - It is impossible to make exact measurements

SIGNIFICANT FIGURES Measurements contain 2 types of information - Magnitude of the measurement - Uncertainty of the measurement Only one uncertain or estimated digit should be reported Significant Figures digits known with certainty + one uncertain digit

RULES FOR SIGNIFICANT FIGURES 1. Nonzero integers are always significant 4732 (4 sig. figs.)875 (3 sig. figs.) 2. Leading zeros are not significant (2 sig. figs.) (4 sig. figs.) The zeros simply indicate the position of the decimal point 3. Captive zeros (between nonzero digits) are always significant (5 sig figs.) (8 sig figs)

RULES FOR SIGNIFICANT FIGURES 4. Trailing zeros (at the right end of a number) are significant only if the number contains a decimal point (5 sig figs) (3 sig figs) 5. Exact numbers (not obtained from measurements) are assumed to have infinite number of significant figures

RULES FOR SIGNIFICANT FIGURES How many significant figures are present in each of the following? What is the uncertainty in each case? significant figuresuncertainty 1.24 g3± 0.01 g L 2± L mL 5± mL kg 5± kg

RULES FOR SIGNIFICANT FIGURES Rounding off Numbers 1. In a series of calculations, carry the extra digits through to the final result before rounding off to the required significant figures 2. If the first digit to be removed is less than 5, the preceding digit remains the same (round down) Round to two significant figures 2.53 rounds to 2.5 and 1.24 rounds to 1.2

RULES FOR SIGNIFICANT FIGURES Rounding off Numbers 3. If the first digit to be removed is greater than 5, the preceding digit increases by 1 (round up) (2.56 rounds to 2.6 and 1.27 rounds to 1.3) 4. If the digit to be removed is exactly 5 (round even) - The preceding number is increased by 1 if that results in an even number (2.55 rounds to 2.6 and rounds to 1.4) - The preceding number remains the same if that results in an odd number (2.45 rounds to 2.4 and rounds to 1.2)

RULES FOR SIGNIFICANT FIGURES Multiplication and Division - The result contains the same number of significant figures as the measurement with the least number of significant figures x 4.02 = = ÷ 1.2 = = The certainty of the calculated quantity is limited by the least certain measurement, which determines the final number of significant figures

RULES FOR SIGNIFICANT FIGURES Addition and Subtraction - The result contains the same number of decimal places as the measurement with the least number of decimal places - The certainty of the calculated quantity is limited by the least certain measurement, which determines the final number of significant figures = 5.5 = 4.03= 6.00 (not 6)

SCIENTIFIC NOTATION - Used to express too large or too small numbers (with many zeros) in compact form - The product of a decimal number between 1 and 10 (the coefficient) and 10 raised to a power (exponential term) 24,000,000,000,000 = 2.4 x coefficient exponential term exponent (power) = 4.58 x 10 -7

SCIENTIFIC NOTATION - Provides a convenient way of writing the required number of significant figures to 4 significant figures = x to 3 significant figures = 2.40 x to 2 significant figures = 3.0 x 10 -4

SCIENTIFIC NOTATION - Add exponents when multiplying exponential terms (5.4 x 10 4 ) x (1.23 x 10 2 ) = (5.4 x 1.23) x = 6.6 x Subtract exponents when dividing exponential terms (5.4 x 10 4 )/(1.23 x 10 2 ) = (5.4/1.23) x = 4.4 x 10 2

MEASUREMENT SYSTEMS Two measurement systems: English System of Units (commercial measurements): pound, quart, inch, foot, gallon Metric System of Units (scientific measurements) SI units (Systeme International d’Unites) liter, meter, gram More convenient to use

FUNDAMENTAL (BASE) UNITS Physical Quantity Mass Length Time Temperature Amount of substance Electric current Luminous intensity Name of Unit Kilogram Meter Second Kelvin Mole Ampere Candela Abbreviation kg m s (sec) K mol A cd

Area = length x length = m x m = m 2 Volume = m x m x m = m 3 Volume may also be expressed in LITERS (L) 1L = 1000 mL = 1000 cm 3 or cubic centimeters (c.c.) Implies 1mL = 1c.c. mL is usually used for volumes of liquids and gases c.c. is usually used for volumes of solids Density = kg/ m 3 DERIVED UNITS

Physical Quantity Force Pressure Energy Power Frequency Name of Unit Newton Pascal Joule Watt Hertz Abbreviation N (m-kg/s 2 ) Pa (N/m 2 ; kg/(m-s 2 ) J (N-m; m 2 -kg/s 2 ) W (J/s; m 2 -kg/s 3 ) Hz (1/s)

Prefix Giga Mega Kilo Deci Centi Milli Micro Nano Pico Femto Abbreviation G M k d c m µ n p f Notation UNIT CONVERSIONS

1 gigameter (Gm) 1 megameter (Mm) 1 kilometer (km) 1 decimeter (dm) 1 centimeter (cm) 1 millimeter (mm) 1 micrometer (µm) 1 nanometer (nm) 1 picometer (pm) 1 femtometer (fm) = 10 9 meters = 10 6 meters = 10 3 meters = meter = meter = meter = meter = meter = meter = meter

UNIT CONVERSIONS Length/Distance 2.54 cm = 1.00 in. 12 in. = 1 ft 1 yd = 3 ft 1 m = 39.4 in. 1 m = 1.09 yd 1 km = mile 1 km = 1000 m Time 1 min = 60 sec 1 hour = 60 min 24 hours = 1 day 7 days = 1 week Volume 1 gal = 4 qt 1 qt = L 1 L = 1.06 qt 1 L = gal 1 mL = fl. oz. Mass 1 Ib = 454 g 1 Ib = 16 oz 1 kg = 2.20 lb 1 oz = 28.3 g

UNIT CONVERSIONS 1 km = 1000 mor» Conversion Factors 1 L = 1000 mL 24 hours = 1 day 1 kg = 2.20 lb » » » or

UNIT CONVERSIONS given number · unitnew unit unit to be converted = new number · new unit quantity to be expressed in new units conversion factor quantity now expressed in new units given datadesired unit unit of given data = answer in desired unit

Convert 34.5 mg to g How many gallons of juice are there in 20.0 liters of the juice? UNIT CONVERSIONS Convert 4.0 gallons to quarts

Convert 2.64 μg to kg Convert m 2 to km 2 UNIT CONVERSIONS Convert 4.0 cm 3 to μm 3

DENSITY - The amount of mass in a unit volume of a substance Ratio of mass to volume =Density = Units Solids: grams per cubic centimeter (g/cm 3 ) Liquids: grams per milliliter (g/mL) Gases: grams per liter (g/L) - Density of 2.3 g/mL implies 2.3 grams per 1 mL - Density usually changes with change in temperature

For a given liquid: - Objects with density less than that of the liquid will float - Objects with density greater than that of the liquid will sink - Objects with density equal to that of the liquid will remain stationary (neither float nor sink) DENSITY

The liquid level in a graduated cylinder reads mL. The level rises to mL when g of piece of gold is added to the cylinder. What is the density of gold? Volume of the piece of gold = mL – mL = 6.70 mL Mass of the piece of gold = g Density = mass/volume = g/6.70 mL = 19.3 g/mL or 19.3 g/cm 3 DENSITY

TEMPERATURE - The degree of hotness or coldness of a body or environment 3 common temperature scales Metric system: Celsius and Kelvin English system: Fahrenheit Celsius Scale ( o C): Reference points are the boiling and freezing points of water (0 o C and 100 o C) degree interval Kelvin Scale (K): Is the SI unit of temperature (no degree sign) The lowest attainable temperature on the Kelvin scale is 0 (-273 o C) referred to as the absolute zero

TEMPERATURE Fahrenheit Scale: Water freezes at 32 o F and boils at 212 o F degree interval or 10 o, 40 o, 60 o may be considered as 2 significant figures 100 o may be considered as 3 significant figures or

TEMPERATURE Convert 29 o C to K Convert 29 K to o C Convert 29 o F to o C Convert 29 o C to o F

TEMPERATURE Heat A form of energy necessary to raise the temperature of a substance Units: Calorie (cal) or joules (J) [1 cal = J] Specific Heat The quantity of heat energy necessary to raise the temperature of 1 gram of a substance by 1 o C Units: cal/g. o C

TEMPERATURE Calorie The amount of heat energy needed to raise the temperature of 1 gram of water by 1 degree Celsius

PERCENTAGE - per one hundred The chemistry class at CSU is made up of 39 females and 12 males. What percentage of the class are females and males