Chemistry: The Study of Change

Chemistry is the study of matter and the changes it undergoes Matter is anything that occupies space and has mass. Matter may exist as a (pure) substance? 1.4

What is a pure substance? Group Discussion What is a pure substance?

Pure Substance A (pure) substance is a form of matter that has a definite composition and distinct properties. Examples: water, ammonia, sucrose, gold, oxygen

soft drink, milk, salt water A mixture is a combination of two or more substances in which the substances retain their distinct identities. Homogenous mixture – composition of the mixture is the same throughout. soft drink, milk, salt water Heterogeneous mixture – composition is not uniform throughout. Oil and water, wood, iron filings in sand, sand in water 1.4

Physical means can be used to separate a mixture into its pure components. distillation magnet 1.4

115 elements have been identified An element is a substance that cannot be separated into simpler substances by chemical means. 115 elements have been identified 83 elements occur naturally on Earth gold, aluminum, lead, oxygen, carbon 32 elements have been created by scientists technetium, americium, seaborgium 1.4

A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions. Compounds can only be separated into their pure components (elements) by chemical means. Water (H2O) Glucose (C6H12O6) Ammonia (NH3) 1.4

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Three States of Matter 1.5

hydrogen gas burns in oxygen gas to form water Physical or Chemical? A physical change does not alter the composition or identity of a substance. ice melting sugar dissolving in water A chemical change alters the composition or identity of the substance(s) involved. hydrogen gas burns in oxygen gas to form water 1.6

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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 A 1 kg bar will weigh 2.2 lb on earth 0.4 lb on moon 1.7

Measurements

All measured quantities make known three pieces of information. Measurements All measured quantities make known three pieces of information. The quantity or number The unit The uncertainty in the measurement.

Table 1.2 SI Base Units Base Quantity Name of Unit Symbol Length meter Mass kilogram kg Time second s Temperature kelvin K Amount of substance mole mol 1.7

Derived SI Units Quantity Definition of Quantity SI unit Area Length squared m2 Volume Length cubed m3 Density Mass per unit volume kg/m3 Force Mass times acceleration of object kg * m/s2 ( =newton, N) Pressure Force per unit area kg/(ms2) ( = pascal, Pa) Energy Force times distance traveled kg * m2/s2 ( = joule, J) Speed Distance traveled per unit time m/s Acceleration Speed changed per unit time m/s2

Table 1.3 Prefixes Used with SI Units Symbol Meaning Tera- T 1012 Giga- G 109 Mega- M 106 Kilo- k 103 Deci- d 10-1 Centi- c 10-2 Milli- m 10-3 Micro- 10-6 Nano- n 10-9 Pico- p 10-12 1.7

Volume – is length cubed or L3 SI derived unit for volume is cubic meter (m3) 1 L = 1 dm3 (Definition) 1 mL = 1 x 10-3 L or 1L = 1 x 103 mL 1 dm3 = (10 cm)3 = 1000 cm3 1 L = 1000 mL = 1000 cm3 = 1 dm3 1 mL = 1 cm3 1.7

Density – SI derived unit for density is kg/m3 1 g/cm3 = 1 g/mL = 1000 kg/m3 density = mass volume d = m V 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 1.7

K = 0C + 273.15 273 K = 0 0C 373 K = 100 0C 0F = x 0C + 32 9 5 32 0F = 0 0C 212 0F = 100 0C 1.7

Convert 172.9 0F to degrees Celsius. 0F = x 0C + 32 9 5 0F – 32 = x 0C 9 5 x (0F – 32) = 0C 9 5 0C = x (0F – 32) 9 5 0C = x (172.9 – 32) = 78.3 9 5 1.7

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: 0.0000000000000000000000199 1.99 x 10-23 N x 10n N is a number between 1 and 10 n is a positive or negative integer 1.8

Scientific Notation Addition or Subtraction 568.762 0.00000772 move decimal left move decimal right n > 0 n < 0 568.762 = 5.68762 x 102 0.00000772 = 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 104 + 3.9 x 103 = 4.31 x 104 + 0.39 x 104 = 4.70 x 104 1.8

Scientific Notation Multiplication Division (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 Multiply N1 and N2 Add exponents n1 and n2 Division 8.5 x 104 ÷ 5.0 x 109 = (8.5 ÷ 5.0) x 104-9 = 1.7 x 10-5 Divide N1 and N2 Subtract exponents n1 and n2 1.8

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Significant Figures

SIGNIFICANT FIGURES Except when all numbers are integers it is impossible to measure the exact value of a quantity. The uncertainty in a measurement is indicated by the number of significant figures, which are the meaningful digits in a measured or calculated quantity. The last digit is understood to be uncertain by + or - 1.

Measurement Uncertainty 6 mL (+/- 1 mL) 6.0 mL (+/- 0.1 mL) 6.00 mL (+/- 0.01 mL) As a rule of thumb estimate 1 figure beyond the smallest subdivision on the scale.

The Number of Significant Figures in a Measurement Depends Upon the Measuring Device Fig 1.15A As a rule of thumb estimate 1 figure beyond the smallest subdivision on the scale.

Fig. 1.6

Significant Figures Any digit that is not zero is significant 1.234 kg 4 significant figures Zeros between nonzero digits are significant 606 m 3 significant figures Zeros to the left of the first nonzero digit are not significant 0.08 L 1 significant figure If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg 2 significant figures 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 0.00420 g 3 significant figures (4.20 mg) 1.8

How many significant figures are in each of the following measurements? 24 mL 2 significant figures 3001 g 4 significant figures 0.0320 m3 3 significant figures 6.4 x 104 molecules 2 significant figures 560 kg 2 significant figures 1.8

Do this practice exercise with your group For numbers that do not contain decimal points, the trailing zeros may or may not be significant. Use scientific notation to avoid ambiguity. Example 1.3 478 cm b. 6.01 cm c. 0.825 m d. 0.043 kg e. 1.310 x 1022 atoms f. 7000 mL Do this practice exercise with your group

Significant Figures Addition or Subtraction – use decimal places 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 -2.9133 0.7867 two significant figures after decimal point round off to 0.79 1.8

Significant Figures Multiplication or Division – use # of sig. figs. 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 3.6666 = 16.536366 = 16.5 3 sig figs round to 3 sig figs 6.8 ÷ 112.04 = 0.0606926 = 0.061 2 sig figs round to 2 sig figs 1.8

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? 6.64 + 6.68 + 6.70 3 = 6.67333 = 6.67 = 7 Because 3 is an exact number 1.8

Example 1.4 Practice Exercise 26.5862 L + 0.17 L 9.1 g – 4.682 g 7.1 x 104 dm x 2.2654 x 102 dm 6.54 g / 86.5542 mL 7.55 x 104 m – 8.62 x 103 m

Accuracy – how close a measurement is to the true value Precision – how close a set of measurements are to each other accurate & precise precise but not accurate not accurate & not precise 1.8

Factor-Label Method of Solving Problems Write down starting number with units. 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. How many mL are in 1.63 L? 1 L = 1000 mL 1L 1000 mL 1.63 L x = 1630 mL 1L 1000 mL 1.63 L x = 0.001630 L2 mL 1.9

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? 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 1.9

Common SI-English Equivalent Quantities Quantity English to SI Equivalent Length 1 mile = 1.61 km 1 yard = 0.9144 m 1 foot (ft) = 0.3048 m 1 inch = 2.54 cm (exactly!) Volume 1 cubic foot = 0.0283 m3 1 gallon = 3.785 dm3 1 quart = 0.9464 dm3 1 quart = 946.4 cm3 1 fluid ounce = 29.6 cm3 Mass 1 pound (lb) = 0.4536 kg 1 pound (lb) = 453.6 g 1 ounce = 28.35 g

Example 1.5 Practice exercise A roll of Al foil has a mass of 1.07 kg. What is the mass in pounds. Example 1.6 Practice exercise The density of silver is 10.5 g/cm3. Convert this to kg/m3.

Sample Problem The volume of an irregularly shaped solid can be determined from the volume of water it displaces. A graduated cylinder contains 245.0 mL water. When a small piece of Pyrite, an ore of Iron, is submerged in the water, the volume increases to 315.8 mL. What is the volume of the piece of Pyrite in cm3 and in liters. Vol (mL) = 315.8 mL - 245.0 mL = 70.8 mL Vol (cm3) = 70.8 mL x 1 cm3/ 1 mL = 70.8 cm3 Vol (liters) = 70.8 mL x 10 –3 L / mL = 7.08 x 10 -2 L

Density Density = mass/volume (an intensive quantity) A small rectangular slab of lithium has a mass of 1.49 g and measures 2.09 cm by 1.11 cm by 1.19 cm. Find its density in g/mL and lb/in3. What is the volume of 15.3 g of Li? What is the mass of 134 mL of Li?

Densities of Some Common Substances Substance Physical State Density (g/cm3) Hydrogen Gas 0.000089 Oxygen Gas 0.0014 Grain alcohol Liquid 0.789 Water Liquid 1.0 Table salt Solid 2.16 Aluminum Solid 2.70 Lead Solid 11.3 Gold Solid 19.3