Chemistry: The Central Science PDF

Chapter 1 Chemistry: The Central Science 1  1.1 The Study of Chemistry Chemistry – the study of all aspects of matter and the changes that matter undergoes Matter – anything that has mass and occupies space...

Chapter 1 Chemistry: The Central Science 1  1.1 The Study of Chemistry Chemistry – the study of all aspects of matter and the changes that matter undergoes Matter – anything that has mass and occupies space 2  Chemistry you may already know – Familiar terms: molecules, atoms, and chemical reactions Familiar chemical formula: H2O 3  Molecules can be represented several different ways including molecular formulas and molecular models. – Molecular models can be “ball-and-stick” or “space-fill.” Each element is represented by a particular color 4 5  1.2: Classification of Matter Matter is either classified as a substance or a mixture of substances. Substance – Can be either an element or a compound – Has a definite (constant) composition and distinct properties – Examples: sodium chloride, water, oxygen 6  Substances Element: cannot be separated into simpler substances by chemical means. – Examples: iron, mercury, oxygen, and hydrogen Compounds: two or more elements chemically combined in definite ratios – Cannot be separated by physical means – Examples: salt, water and carbon dioxide 7  Mixtures Mixture: physical combination of two or more substances – Substances retain distinct identities – No universal constant composition – Can be separated by physical means Examples: sugar/iron; sugar/water 8 Molecular Comparison of Substances and Mixtures Atoms of an element Molecules of an element Molecules of a compound Mixture of two elements and a compound 9  Types of Mixtures – Homogeneous: composition of the mixture is uniform throughout Example: sugar dissolved in water – Heterogeneous: composition is not uniform throughout Example: sugar mixed with iron filings 10 Classification of Matter 11  Classify the following Aluminum foil Baking soda Milk Air Copper wire 12 Aluminum foil: substance, element Baking soda: substance, compound Milk: mixture, homogeneous Air: mixture, homogeneous Copper wire: substance, element 13  1.3 Properties of Matter Quantitative: expressed using numbers Qualitative: expressed using properties Physical properties: can be observed and measured without changing the substance – Examples: color, melting point, states of matter Physical changes: the identity of the substance stays the same – Examples: changes of state (melting, freezing) 14  States of Matter – Solid particles close together in orderly fashion little freedom of motion a solid sample does not conform to the shape of its container – Liquid particles close together but not held rigidly in position particles are free to move past one another a liquid sample conforms to the shape of the part of the container it fills 15  – Gas particles randomly spread apart particles have complete freedom of movement a gas sample assumes both shape and volume of container. States of matter can be inter-converted without changing chemical composition solid ® liquid ® gas (add heat) gas ® liquid ® solid (remove heat) 16 States of Matter 17  Chemical properties: must be determined by the chemical changes that are observed – Examples: flammability, acidity, corrosiveness, reactivity Chemical changes: after a chemical change, the original substance no longer exists – Examples: combustion, digestion 18  Extensive property: depends on amount of matter – Examples: mass, length Intensive property: does not depend on amount – Examples: density, temperature, color 19  1.4 Scientific Measurement Used to measure quantitative properties of matter SI base units 20 SI Prefixes 21  Mass: measure of the amount of matter – (weight refers to gravitational pull) Temperature: – Celsius Represented by °C Based on freezing point of water as 0°C and boiling point of water as 100°C – Kelvin Represented by K (no degree sign) The absolute scale Units of Celsius and Kelvin are equal in magnitude – Fahrenheit (the English system) (°F) 22 Equations for Temperature Conversions 5 o C = ( F - 32) ´ o 9 K = C + 273.15 o 9 o F = ´ C + 32 o 5 23  Temperature Conversions A clock on a local bank reported a temperature reading of 28oC. What is this temperature on the Kelvin scale? K = C + 273.15 o K = 28 C + 273.15 = 301 K o 24  Practice Convert the temperature reading on the local bank (28°C) into the corresponding Fahrenheit temperature. 9 o F = ´ C + 32 o 5 9 o F = ´ 28 C + 32 = 82 F o o 5 25  Volume: meter cubed (m3) – Derived unit – The unit liter (L) is more commonly used in the laboratory setting. It is equal to a decimeter cubed (dm3). 26  Density: Ratio of mass to volume – Formula: m d= V – d = density (g/mL) – m = mass (g) – V = volume (mL or cm3) (*gas densities are usually expressed in g/L) 27  Practice The density of a piece of copper wire is 8.96 g/cm3. Calculate the volume in cm3 of a piece of copper with a mass of 4.28 g. m d= V m 4.28 g V= = = 0.478 cm 3 d 8.96 g 3 cm 28  1.5 Uncertainty in Measurement Exact: numbers with defined values – Examples: counting numbers, conversion factors based on definitions Inexact: numbers obtained by any method other than counting – Examples: measured values in the laboratory 29  Significant figures – Used to express the uncertainty of inexact numbers obtained by measurement – The last digit in a measured value is an uncertain digit - an estimate 30 Guidelines for significant figures – Any non-zero digit is significant – Zeros between non-zero digits are significant – Zeros to the left of the first non-zero digit are not significant – Zeros to the right of the last non-zero digit are significant if decimal is present – Zeros to the right of the last non-zero digit are not significant if decimal is not present 31  Practice Determine the number of significant figures in each of the following. 345.5 cm 4 significant figures 0.0058 g 2 significant figures 1205 m 4 significant figures 250 mL 2 significant figures 250.00 mL 5 significant figures 32 Calculations with measured numbers – Addition and subtraction Answer cannot have more digits to the right of the decimal than any of original numbers Example: 102.50 two digits after decimal point + 0.231 three digits after decimal point 102.731 round to 102.73 33  Multiplication and division – Final answer contains the smallest number of significant figures – Example: 1.4 x 8.011 = 11.2154 round to 11 (Limited by 1.4 to two significant figures in answer) 34  Exact numbers – Do not limit answer because exact numbers have an infinite number of significant figures – Example: A penny minted after 1982 has a mass of 2.5 g. If we have three such pennies, the total mass is 3 x 2.5 g = 7.5 g – In this case, 3 is an exact number and does not limit the number of significant figures in the result. 35  Multiple step calculations – It is best to retain at least one extra digit until the end of the calculation to minimize rounding error. Rounding rules – If the number is less than 5 round “down”. – If the number is 5 or greater round “up”. 36  Practice 105.5 L + 10.65 L = 116.2 L Calculator answer: 116.15 L Round to: 116.2 L Answer to the tenth position 1.0267 cm x 2.508 cm x 12.599 cm = 32.44 cm3 Calculator answer: 32.4419664 cm3 Round to: 32.44 cm3 round to the smallest number of significant figures 37  Accuracy and precision – Two ways to gauge the quality of a set of measured numbers – Accuracy: how close a measurement is to the true or accepted value – Precision: how closely measurements of the same thing are to one another 38 both accurate and precise not accurate but precise neither accurate nor precise 39 Describe accuracy and precision for each set Student A Student B Student C 0.335 g 0.357 g 0.369 g 0.331 g 0.375 g 0.373 g 0.333 g 0.338 g 0.371 g Average: 0.333 g 0.357 g 0.371 g True mass is 0.370 grams 40 Student A’s results are precise but not accurate. Student B’s results are neither precise nor accurate. Student C’s results are both precise and accurate. 41  1.6 Using Units and Solving Problems Conversion factor: a fraction in which the same quantity is expressed one way in the numerator and another way in the denominator – Example: by definition, 1 inch = 2.54 cm 1 in 2.54 cm 2.54 cm 1 in 42  Dimensional analysis: a problem solving method employing conversion factors to change one measure to another often called the “factor-label method” – Example: Convert 12.00 inches to meters Conversion factors needed: 2.54 cm = 1 in and 100 cm = 1 meter 2.54 cm 1m 12.00 in ´ ´ = 0.3048 m 1 in 100 cm *Note that neither conversion factor limited the number of significant figures in the result because they both consist of exact numbers. 43  Notes on Problem Solving Read carefully; find information given and what is asked for Find appropriate equations, constants, conversion factors Check for sign, units and significant figures Check for reasonable answer 44  Practice The Food and Drug Administration (FDA) recommends that dietary sodium intake be no more than 2400 mg per day. What is this mass in pounds (lb), if 1 lb = 453.6 g? 1g 1 lb -3 2400 mg ´ ´ = 5.3 ´ 10 lb 1000 mg 453.6 g 45  Key Points Classifying matter SI conversions Density Temperature conversions Physical vs chemical properties and changes Dimensional analysis 46