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Chapter 1: Matter and Measurements

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## Chapter 1: Matter and Measurements

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**What is Chemistry?**Biology vs. Chemistry vs. Physics**What is Chemistry?**Biology Physics Chemistry The study of living organisms The study of forces & motion The study of matter and its reactions and properties**What is Chemistry?**Chemistry is the study of CHANGES in the “stuff” around us. (We formally define “stuff” as matter!)**What is Chemistry?**Think, pair, share: What are all the chemicals you use in your daily life?**Review: Scientific Methods**1. Hypothesis • Suggested solution to a problem 2. Experiment • A controlled method of testing a hypothesis 3. Data • Organized observations a. Data is always reproducible.**Review: Scientific Methods**4. Scientific Law • Statement which summarizes results of many observations and experiment a. Scientific laws explain WHAT is observed. • Example of a scientific law: 5. Scientific Theory • Explanation that supports a hypothesis and which has been supported with repeated testing b. Scientific theories explain WHY something is observed. • Example of a scientific theory:**Review: Scientific Methods**6. Steps of the Scientific Method—Review a. b. c. d. e. f.**What is Chemistry?**Biology Physics Chemistry The study of living organisms The study of forces & motion The study of matter and its reactions and properties**What is Chemistry?**Chemistry is the study of CHANGES in the “stuff” around us. (We formally define “stuff” as matter!)**Elements**• Type of matter that cannot be broken down into simpler, stable substances and is made of only one type of atom**Compounds**• A pure substance that contains two or more elements whose atoms are chemically bonded**Compounds**• Fixed compositions • A given compound contains the same elements in the same percent by mass**Compounds**• The properties of a compound are VERY DIFFERENT from the properties of the elements they contain Ex.) Sodium Chloride (NaCl) vs. Sodium & Chlorine Sodium: http://www.youtube.com/watch?v=RAFcZo8dTcU http://www.youtube.com/watch?v=92Mfric7JUc**Mixtures**• A blend of two or more kinds of matter, each of which retains its own identity and properties • Homogeneous • Heterogeneous**Homogeneous Mixtures**• Composition is the same throughout the mixture • Examples: salt water, soda water, brass • A.k.a. a solution • Solute in a solvent (salt dissolved in water)**Heterogeneous Mixtures**• Non-uniform; composition varies throughout the mixture**Separating Mixtures**• Filtration**Separating Mixtures**• Distillation**Separating Mixtures**• Chromatography**Scientific Measurements**Chemistry is a quantitative science. • This means that experiments and calculations almost always involve measured values. Scientific measurements are expressed in the SI (metric) system. • This is a decimal-based system in which all of the units of a particular quantity are related to each other by factors of ten.**SI System**• Definition: modernized version of metric system; uses decimals • All units derived from base units; larger and smaller quantities use prefixes with base unit • Must memorize prefixes from nano- (10-9) to tera- (1012)**Prefixes (see handout & Ebook)**• You will need to memorize all of the prefixes (factors, names and abbreviations from 109 (giga-) to 10-9 (nano)! • One example of a memory device:**INSTRUMENTS & UNITS**Use SI units — based on the metric system Length Mass Time Temperature Meter, m Kilogram, kg Seconds, s Celsius degrees, ˚C Kelvins, K**Length**The standard unit of length in the metric system is the METER which is a little larger than a YARD. USING THE PREFIXES WITH LENGTH: cm – often used in lab km – Gm –**Length**Base unit: METER Conversions: 1 km=1000 m 1 cm = 10 -2 m 1 Gm = 106 m**O—H distance =**9.58 x 10-11 m 9.58 x 10-9 cm 0.0958 nm Units of Length • 1 kilometer (km) = 1000 meters (m) • 10-2 meter (m) = 1 centimeter (cm) • 102 meter (m) = 1 hectometer (Hm) • 1 nanometer (nm) = 1.0 x 10-9 meter**Volume**THE COMMON UNITS OF VOLUME IN CHEMISTRY ARE: the liter (milliliter) and cubic centimeter (cm3) THE COMMON INSTRUMENTS FOR MEASURING VOLUME IN CHEMISTRY ARE: graduated cylinder & buret Note that 1 cm3 = 1 mL (We will use this exact conversion factor throughout the year, so you will need to memorize it!)**Mass**• THE COMMON UNIT OF MASS IN CHEMISTRY IS : the gram (g) — often used in lab • Mass IS A MEASURE OF THE AMOUNT OF MATTER IN AN OBJECT; • Weight IS A MEASURE OF THE GRAVITATIONAL FORCE ACTING ON THE OBJECT. CHEMISTS OFTEN USE THESE TERMS INTERCHANGEABLY. 1000 g= 1 kg 1 Mg = 10 6 g**Temperature Scales**Fahrenheit Celsius Kelvin Anders Celsius 1701-1744 Lord Kelvin (William Thomson) 1824-1907 TEMPERATURE IS THE FACTOR THAT DETERMINES the direction of heat flow.**Temperature Scales**212 ˚F 100 ˚C 373 K 100 K 180˚F 100˚C 32 ˚F 0 ˚C 273 K Fahrenheit Celsius Kelvin Boiling point of water Freezing point of water Notice that 1 kelvin degree = 1 degree Celsius**Temperature Scales**100 oF 38 oC 311 K oF oC K**SI System**English Units (inches, feet, degrees F, etc.) are NEVER used to take measurements in the lab!**Calculations Using Temperature**Fahrenheit/Celsius T (F) = 1.8 t (˚C) + 32**Calculations Using Temperature**• Some calculations are in kelvins (especially important for Ch 5!!) • T (K) = t (˚C) + 273.15 (273) • Body temp = 37 ˚C + 273 = 310 K • Liquid nitrogen = -196 ˚C + 273 = 77 K**Problem**Example 1L.1 A baby has a temperature of 39.8oC. Express this temperature in oF and K.**SI System: Base Units**ESTABLISHMENT OF THE INTERNATIONAL SYSTEM OF UNITS (SI)— SI UNITS AS ESTABLISHED BY THE SI: LENGTH – meter (m) VOLUME – cubic meter (m3) MASS – kilogram (kg) TEMPERATURE – Kelvin (K)**Time**Base unit: SECOND (sec) • Conversions: only non-decimal base unit 60 sec = 1 min 60 min = 1 hr**Precision and Accuracy in Measurements**Precision vs. Accuracy Definitions: Precision—how close answers are to each other (reproducibility) Accuracy—how close answer is to accepted (true) value (agreement to accepted value)**Precision and Accuracy in Measurements**Percent Error - a way to calculate accuracy in the lab Equation: % Error = | Accepted Value – Exp. Value | x 100 Accepted Value**Precision and Accuracy in Measurements**Ex1.9 A student reports the density of a pure substance to be 2.83 g/mL. The accepted value is 2.70 g/mL. What is the percent error for the student’s results? Equation: % Error = | Accepted Value – Exp. Value | x 100 Accepted Value**Scientific Notation**Exponential (Scientific) Notation—See Worksheet**Significant Figures: Why are they Important?**Numbers in math: no units, abstract, no context, can read calculator output exactly for answer. vs. Numbers in chemistry: measurements – include units. SIG FIGS WILL BE IMPORTANT THROUGHOUT THIS COURSE!**Graduated Cylinder Example**http://learningchemistryeasily.blogspot.com/2013/07/precision-of-measurement-and.html**What are significant figures?(aka sig figs)**• Significant figures are all the digits in a measurement that are known with certainty plus a last digit that must be estimated. • With experimental values your answer can have too few or too many sig figs, depending on how you round.**How Rounding Influences Sig Figs**• 1.024 x 1.2 = 1.2288Too many numerals(sig figs) Too precise • 1.024 x 1.2 = 1Too few numerals(sig figs) Not precise enough**Why This Concept is Important**• We will be adding, subtracting, multiplying and dividing numbers throughout this course. • You MUST learn how many sig figs to report each answer in or the answer is meaningless. • You must report answers on lab reports & tests/quizzes with the correct number of sig figs (+/- 1) or else you will lose points!!