Essential Ideas-Chemistry - K. A. Boudreaux, Angelo State University, Lecture notes of Chemistry

Download Essential Ideas-Chemistry - K. A. Boudreaux, Angelo State University and more Lecture notes Chemistry in PDF only on Docsity! Chapter Objectives: • Learn to use and manipulate units, and convert from one unit to another. • Learn to use the appropriate number of significant figures in measurements and calculations. • Learn how to classify matter by state and composition. Chapter 1 Essential Ideas Mr. Kevin A. Boudreaux Angelo State University CHEM 1311 General Chemistry Chemistry 2e (Flowers, Theopold, Langley, Robinson; openstax, 2nd ed, 2019) What Is Chemistry? • Chemistry is the science that seeks to understand the composition, properties, and transformations of matter by studying the behavior of atoms and molecules. • Chemistry is subdivided into different specialized fields: organic chemistry, inorganic chemistry, physical chemistry, biochemistry, analytical, chemistry, environmental chemistry, etc. • We study chemistry to provide ourselves with a better understanding of the underlying workings of nature, to learn how to make new materials with useful properties that satisfy particular needs. Chemistry intersects with other important fields, such as biology, molecular biology and genetics, medicine, physics, etc. 2 5 The SI System • In science, the most commonly used set of units are those of the International System of Units (the SI System, for Système International d’Unités). • There are seven fundamental units in the SI system. The units for all other quantities (e.g., area, volume, energy) are derived from these base units. Physical Quantity Unit Abbreviation Mass kilogram kg Length meter m Time second s Amount of substance mole mol Temperature kelvin K Electric current ampere A Luminous intensity candela cd The SI System — Mass, Length, and Time • The meter (m) is the SI unit of length. – Defined as the distance light travels in a vacuum in 1/2999,792,458 seconds. – 1 m = 39.37 in = 1.094 yards • The kilogram (kg) is the SI unit of mass. – 1 kg = 2.205 pounds (Avoirdupois) • The second (s) is the SI unit of time. • The mole (mol) is the SI unit for the amount of a substance (we’ll get into this much more later on). 6 7 The SI System — Temperature • In the SI system, temperature is measured in kelvins (K), but often the Celsius degree, °C, is used instead. • A kelvin is the same size as a Celsius degree, but with the zero point set at the coldest possible temperature, absolute zero (-273.15°C). • In most mathematical formulas, K must be used instead of °C. (Is It Hot In Here, Or Is It Me?) ( )32 - F C 32 C F 273.15 C K 9 5 5 9 = += += 10 Derived Units — Density • Density, the ratio of an object’s mass (m) to its volume (V), is given by the formula: V m d = – Density has units of mass over volume: g/mL, g/L, lb/gal, kg/m3, lb/ft3, etc. – Because volume changes with temperature, density is temperature-dependent. Density of water vs. temperature Densities of Various Substances at 20ºC Substance State Density (g/cm3) Oxygen gas 0.00133 Hydrogen gas 0.000084 Ethanol liquid 0.789 Benzene liquid 0.880 Water liquid 0.9982 Magnesium solid 1.74 Sodium chloride solid 2.16 Aluminum solid 2.70 Iron solid 7.87 Copper solid 8.96 Silver solid 10.5 Lead solid 11.34 Mercury liquid 13.6 Gold solid 19.32 11 Extensive and Intensive Properties • Extensive properties depend on the size of the sample (mass, volume, length, etc.). • Intensive properties are independent of the size of the sample (color, melting / boiling point, odor, etc.) – Despite the fact that the mass and volume of a sample are extensive properties, the density of a pure substance is an intensive property. 12 Larger and Smaller Units • In many instances, decimal multipliers are added to the units in cases where numbers are inconveniently large or small: – the diameter of a sodium atom: • long-hand: 0.000 000 000 372 m • scientific notation: 3.72×10-10 m • prefix units: 0.372 nm or 372 pm – the distance from the earth to the sun: • long-hand: 150,000,000,000 m • scientific notation: 1.5×1011 m • prefix units: 150 Gm 15 Scientific Notation • In science, we have the opportunity to work with numbers that are extremely large: 602,000,000,000,000,000,000,000 0.000000000000000000000000000911 6.02×1023 9.11×10-28 and numbers that are extremely small: • Numbers like this may be written more compactly using scientific notation: 16 Scientific Notation • To put a number into scientific notation, we move the decimal point behind the first significant figure, and multiply it by the appropriate power of ten. • For a number larger than 1, move the decimal point to the left, behind the first nonzero digit, and use a positive power of ten to indicate how many places the decimal point was moved: 6.02 = 6.02×100 60.2 = 6.02×101 602 = 6.02×102 6,020 = 6.02×103 6,020,000 = 6.02×106 Scientific Notation • For a decimal number (smaller than 1), move the decimal point to the right, behind the first nonzero digit, and use a negative power of ten to indicate how many places the decimal point was moved: 0.602 = 6.02×10-1 0.0602 = 6.02×10-2 0.00602 = 6.02×10-3 0.00000602 = 6.02×10-6 17 20 Examples: Conversions with SI Prefixes 2. A bar of aluminum has a mass of 1210 g. What is its mass in kilograms (kg)? = kg 1 g 1000 g 1210 g ? kg 1 = = g 1000 kg 1 g 1210 g 1000 kg 1 = 21 Examples: Metric Conversions 3. Convert 0.123 cm to mm. Answer: 1.23 mm 22 Examples: Scientific Notation 4. Round off each of the following numbers to the indicated number of significant digits, and write the answer in both long-hand and scientific notation: a. 0.07565 to two significant figures long-hand: _______________ scientific notation: _______________ b. 325801 to three significant figures long-hand: _______________ scientific notation: _______________ 25 Measurement and Significant Figures 26 Uncertainty in Measurement • For instance, on the scale below, the pointer is pointing between “13” and “14”: • We can “guess” that the pointer is about a third of the way to the 14, so we can estimate the reading as 13.3. Since the last digit is an estimation, anything further than that would be a wild guess. • All measurements have some degree of uncertainty to them. 27 Significant Figures • The total number of digits in a measurement is called the number of significant figures. • When reading a scale, the value you record should use all of the digits you are sure of, plus one additional digit that you estimate. This last estimated digit is the last significant figure in your reading. (On a digital readout, the last number on the screen is usually the last significant figure.) • The greater the number of significant figures, the greater the certainty of the measurement. 30 Measured Numbers vs. Exact Numbers • Exact numbers are relationships that are arrived at by counting discrete objects (3 eggs = 3.00000… eggs) or that are true by definition (12 inches = 1 foot, 60 s = 1 min, 5280 feet = 1 mile, 100 cm = 1 m, 2.54 cm = 1 inch, etc.). – There is no uncertainty in these numbers, and they have an infinite number of significant figures (i.e., they do not affect the number of significant figures in the result of a calculation). • All measured numbers will have some limit to how precisely they are known, and there is a limit to the number of significant digits in the number. – This must be taken into account when doing calculations with those numbers! 31 Counting Significant Figures Rules for Counting Significant Figures: a. All nonzero digits are significant. (42 has 2 sf’s.) b. Leading zeros are not significant; they are there to locate the decimal point. (0.00123 g has three sf’s.) c. Zeros in the middle of a number (middle zeros or captive zeros) are significant. (4.803 cm has 4 sf’s.) d. Trailing zeros are significant if the number contains a decimal point. (55.220 K has five sf’s; 50.0 mg has three sf’s, 5.100×10-3 has four sf’s.) e. Trailing zeros are not significant if the number does not contain a decimal. (34,200 m has three sf’s.) – Because trailing zeros can be ambiguous, it is a good practice to avoid potential errors by reporting the number in scientific notation. 32 Counting Significant Figures a. All nonzero digits are significant. 164.87 5 sf’s 395 3 sf’s b. Leading zeros are not significant. 0.766 3 sf’s 0.000033 2 sf’s 00591.3 4 sf’s c. Middle zeros are always significant. 2.028 4 sf’s 5107 4 sf’s 0.00304 3 sf’s 35 Manipulating Significant Figures 278 mi 11.70 gal = 23.8 mi/gal 3 sf's 4 sf's 3 sf's 3.18 + 0.01315 3.19315 2 decimal places 5 decimal places 2 decimal places3.19 36 Adding and Subtracting Measured Numbers • During addition or subtraction, the result has the same number of decimal places as the quantity with the fewest decimal places. – It is possible to either lose or gain significant figures in addition or subtraction. 6.14 + 0.0375 6.1775 6.18 0.002631 - 0.0014278 0.0012032 0.001203 5200 + 63.4 5263.4 5260 12.6198 - 12.5202 0.0996 96.52 + 3.48 100.00 37 Multiplying and Dividing Measured Numbers • During multiplication or division, the result has the same number of sf’s as the factor with the fewest sf’s. 3.0 tooff rounds 526...3.03684210 1.9 5.77 = 46 tooff rounds 68...46.2036001 (5.0) 54)0626)(128.(28.71)(0. = 40 Examples: Significant Figures 2. How many significant figures are in each of the following? a. 0.04450 m b. 1000 m = 1 km c. 0.00002 g d. 5.0003 km e. 1.00010-3 mL f. 10,000 m g. 4080 kg h. 1.4500 L i. 5280 feet = 1 mile j. 2.54 cm = 1 inch k. 0.000304 s 41 Examples: Significant Figures 3. Perform the following calculations to the correct number of significant figures. a. 1.10  0.5120  4.0015 ÷ 3.4555 = b. 0.355 + 105.1 – 100.5820 = c. 4.562  3.99870 ÷ (452.6755 – 452.33) = d. (14.84  0.55) – 8.02 42 Examples: Significant Figures 4. Do the following calculations, rounding the answers off to the correct number of significant digits. =                  in 1 cm 2.54 ft 1 in 12 mile 1 ft 5280 miles 43.12a. ( )( )( ) = cm .051 cm .270 cm 20.1 g 0.009 b. =++ 75.8060g g 35.5 g 8.233c. ( ) = mL 1.15 - mL 20.12 g 65.751 d. 2000414.592 cm 264.5502646 g/cm3 119.539 g 15 g/mL 45 Manipulating Units • Dimensional analysis (also known as the factor- label method) is a way of analyzing the setup of a problem by manipulating the units in the same way you would manipulate the numbers in the calculations. – If your final units are correct, there is a good chance the problem has been set up correctly. – If you end up with incorrect units (for instance, units of time when you’re measuring distance), or units which are clearly nonsense (for instance, cm•inches), the problem has been set up incorrectly. Examples: Unit Conversions 1. Convert 5.0 lb to kg (kilograms) – The conversion factor is 2.2 lb = 1 kg. – Do you multiply 5.0 by 2.2, divide 5.0 by 2.2, or divide 2.2 by 5.0? 46 kg 1 lb 2.2 lb 0.5 =      lb 0.5 1 kg 1 lb 2.2 =      lb 2.2 kg 1 lb 0.5 =      47 Examples: Unit Conversions 2. Convert 58 cm3 to gallons. Answer: 0.015 gal 1 cm3 = 1 mL 1 L = 1000 mL 1 L = 1.057 qt (not exact) 1 gal = 4 qt 50 Examples: Density 5. A common way to measure the density of an irregular solid is by the displacement of water. An irregularly-shaped piece of a shiny yellowish material having a mass of 51.842 g is submerged in a graduated cylinder containing 17.1 mL of water. The volume rose to 19.8 mL. (a) What is the density of this material? (b)Use the table of densities given earlier in this chapter to identify the material. Answer: (a) 19 g/cm3 51 Examples: Unit Conversions 6. The radius of a copper atom is 0.1280 nanometers (nm). What is its radius in picometers (pm)? Answer: 128.0 pm 52 Examples: Unit Conversions 7. How many square centimeters (cm2) are there in 2.00 square meters (m2)? = m 1 cm 100 m .002 2 cm 100 m 1 = =       m 1 cm 100 m .002 2 2 cmm 200 • 2cm 20000 55 Examples: Unit Conversions 10. Convert 37ºC to ºF and Kelvin. Answer: 98.6ºC (what’s wrong with this number?) 56 Examples: Unit Conversions 11. A small hole in the heat shield of a space capsule requires a 32.70 cm2 patch. If the patching material costs NASA $2.75/in2, what is the cost of the patch? (1 in = 2.54 cm [exact]) Answer: $13.94 57 Examples: Unit Conversions 12. A runner wants to run 10.0 km. She knows that her running pace is 7.5 miles per hour. How many minutes must she run? (1 mi = 1.6093 km) Answer: 50. min. 60 The States of Matter • Matter is anything that occupies space and has mass. Matter is classified by its state and by its composition: The States of Matter — Solids • Solids have a fixed shape and volume that does not conform to the container shape. – The atoms or molecules vibrate, but don’t move past each other, making solids rigid (more or less) and incompressible. 61 – In crystalline solids, the atoms and molecules are arranged with some kind of long-range, repeating order, as in diamonds, ice, or salt – In amorphous solids, there is no long- range order, as in charcoal, glass, plastics. 62 The States of Matter — Liquids • Liquids have fixed volumes that conform to the container shape, but forms an upper surface. – The particles are packed closely, as in solids, but are free to move past each other, making them fluid and incompressible (more or less). – e.g., liquid water 65 The Composition of Matter — Pure Substances • Pure Substances: – An element is the simplest type of matter with unique physical and chemical properties. • Elements consist of only one kind of atom. – A compound is a pure substance that is composed of atoms of two or more different elements. • Chemical compounds can be divided into ionic compounds and molecular compounds. • Compounds cannot be broken down by physical means into simpler substances, but can be broken down (although sometimes with difficulty) by chemical reactions. 66 The Composition of Matter — Mixtures • Mixtures: – Heterogeneous mixtures — the mixing is not uniform, and there are regions of different compositions — i.e., there are observable boundaries between the components. • e.g., ice-water, salad dressing, milk, dust in air. – Homogeneous mixtures (or solutions) — the mixing is uniform and there is a constant composition throughout — i.e., there are no observable boundaries because the substances are intermingled on the molecular level. • e.g., salt water, sugar water, metal alloys, air. 67 The Composition of Matter Matter Pure Substances Mixtures Elements Compounds Heterogeneous Mixtures Homogeneous Mixtures Ionic Compounds Molecular Compounds physical separation chemical reactions Atoms and Molecules • Atoms are the smallest particles of an element that has the properties of that element. • Molecules consist of two or more atoms joined by some kind of chemical bond (more later). 70 Gold Gold atoms (STM image) 71 Physical Changes and Physical Properties • A physical change occurs when a substance alters its physical form, but not its composition — the atoms and molecules in the sample retain their identities during a physical change (e.g., ice melting into liquid water, liquid water boiling to steam.) H2O(s) → H2O(l) • Physical properties are properties that do not involve a change in a substance’s chemical makeup (e.g., melting and boiling points, color, density, odor, solubility, etc.). Boiling of water: a physical change 72 Chemical Changes and Chemical Properties • A chemical change is one that alters the composition of matter — the atoms in the sample rearrange their connections in a chemical reaction, transforming the substance into a different substance (e.g., the rusting of iron, the formation of water from hydrogen and oxygen, etc.) 4Fe(s) + 3O2(g) ⎯→ 2Fe2O3(s) 2H2(g) + O2(g) ⎯→ 2H2O(l) • Chemical properties are properties that do involve a change in chemical makeup (e.g., flammability, corrosiveness, reactivity with acids, etc.). Rusting of iron: a chemical change 75 Physical and Chemical Properties of Copper • Copper — Physical Properties – reddish brown, metallic luster – it is malleable (easily formed into thin sheets) and ductile (easily drawn into wires) – good conductor of heat and electricity – can be mixed with zinc to form brass or with tin to form bronze – density = 8.95 g/cm3 – melting point = 1083°C – boiling point = 2570°C • Copper — Chemical Properties – slowly forms a green carbonate in moist air – reacts with nitric acid and sulfuric acid – forms a deep blue solution in aqueous ammonia 76 Examples: Physical and Chemical Changes 1. Which of the following processes are physical changes, and which are chemical changes? a. the evaporation of rubbing alcohol b. the burning of lamp oil c. the bleaching of hair with hydrogen peroxide d. the forming of frost on a cold night e. the beating of a copper wire into a sheet f. a nickel dissolving in acid to produce H2 gas g. dry ice evaporating without melting h. the burning of a log in a fireplace 77 Energy • Energy is defined as the ability to do work or produce heat. – Work is done when a force is exerted through a distance (w = Fd). – Heat is the energy that flows from one object to another because of a temperature difference. • Energy may be converted from one form to another, but it is neither created nor destroyed (the law of conservation of energy). • The total energy possessed by an object is the sum of its kinetic energy (energy of motion) and potential energy (stored energy due to position). • Energy is measured in Joules (J) (kg m2 s-2) or calories (cal).