# 1 Introduction The Scientific Method (1 of 20) 1 Introduction Observations and Measurements Qualitative, Quantitative, Inferences (2 of 20)

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1 The Scientific Method (1 of 20) This is an attempt to state how scientists do science. It is necessarily artificial. Here are MY five steps: Make observations the leaves on my plant are turning yellow State a Problem to be solved how can I get my plants healthy (non-yellow) Form a hypothesis maybe they need more water Conduct a controlled experiment water plants TWICE a week instead of once a week Evaluate results if it works, good... if not, new hypothesis (sunlight?) Observations and Measurements Qualitative, Quantitative, Inferences (2 of 20) Step 1 of the Scientific Method is Make Observations. These can be of general physical properties (color, smell, hardness, etc.) which are called qualitative observations. These can be measurements which are called quantitative observations. There are also statements that we commonly make based on observations. This beaker contains water is an example. You infer (probably correctly) it is water because it is a clear, colorless liquid that came from the tap. The observations are that it is clear, it is colorless, it is a liquid, and it came from the tap. Significant Digits I What do they mean? (3 of 20) Consider: cm In a measurement or a calculation, it is important to know which digits of the reported number are significant. That means if the same measurement were repeated again and again, some of the numbers would be consistent and some would simply be artifacts. All of the digits that you are absolutely certain of plus one more that is a judgment are significant. If all the digits are significant above, everyone who measures the object will determine that it is cm, but some will say 94 cm while others might say 95 cm. a b c Significant Digits II Some examples with rulers. (4 of 20) 1 2 (A composite ruler) a- No one should argue that the measurement is between 0.3 and 0.4. Is it exactly halfway between (.35 cm) or a little to the left (.34 cm)? The last digit is the judgment of the person making the measurement. The measurement has 2 significant digits. b- The same ruler so the measurement still goes to the hundredths place 1.00 cm (3 significant digits). c- A ruler with fewer marks reads 1.6 cm (2 sig digits).

2 Significant Digits III Rules for Recognizing Sig. Digits (5 of 20) In a number written with the correct number of sig. digits... All non-zero digits are significant. 523 grams (3) 0 s in the MIDDLE of a number are ALWAYS significant meters (4) L (4) 0 s in the FRONT of a number are NEVER significant kg (2) m (3) 0 s at the END of a number are SOMETIMES significant. Decimal point is PRESENT, 0 s ARE significant Liters (4) grams (3) Decimal point is ABSENT, 0 s are NOT significant 2000 Liters (1) 550 m (2) NOTE: textbook values are assumed to have all sig. digits Scientific Notation Useful for showing Significant Digits (6 of 20) Scientific notation uses a number between 1 and 9.99 x 10 to some power. It s use stems from the use of slide rules. Know how to put numbers into scientific notation: 5392 = x = 3.28 x = = 5.5 x 10 2 Some 0 s in numbers are placeholders and are not a significant part of the measurement so they disappear when written in sci. notation. Ex: above. In scientific notation, only the three sig. digits (3.28) are written. Scientific Notation can be used to show more sig. digits. Values like 550 ( 2 sig. digits) can be written 5.50 x 10 2 (3) When you perform a calculation using measurements, often the calculator gives you an incorrect number of significant digits. Here are the rules to follow to report your answers: Significant Digits IV Significant Digits in Calculations (7 of 20) x and : The answer has the same # of sig. digits as the number in the problem with the least number of sig. digits. example: 3.7 cm x 8.1 cm = cm 2 (2 sig. digits) + and : The last sig. digit in the answer is the largest uncertain digit in the values used in the problem. example: 3.7 cm cm = 11.8 cm (3 sig. digits) Know how to ilustrate why these rules work. Accuracy vs. Precision (8 of 20) Accuracy refers to how close a measurement is to some accepted or true value (a standard). Ex: an experimental value of the density of Al is 2.69 g/ml. The accepted value is 2.70 g/ml. Your value is accurate to within 0.37% % error is used to express accuracy. Precision refers to the reliability, repeatability, or consistency of a measurement. Ex: A value of 2.69 g/ml means that if you repeat the measurement, you will get values that agree to the tenths place (2.68, 2.70, 2.71, etc.) ± and sig. digits are used to express precision.

3 Metric System (9 of 20) We generally use three types of measurements: volume Liters (ml) length meters (km, cm and mm) mass grams (kg and mg) We commonly use the prefixes: centi- 1 / 100 th milli- 1 / 1000 th kilo Occasionally you will encounter micro(µ), nano, pico, mega, and giga. You should know where to find these in chapter 1. Know that 2.54 cm = 1 inch and 2.20 lb = 1 kg % and ppm (10 of 20) Percentage is a mathematical tool to help compare values. Two fractions, 3/17 and 5/31 are difficult to compare: If we set up ratios so we can have a common denominator: 3 17 = x 100 = = x 100 = so we can see that 3 17 > There are parts per 100 (Latin: parts per centum) or percent (17.65 %) the % is a ppm (parts per million) is the same idea, (use 1,000,000 3 instead of 100) 17 = x = 176,470 ppm Unit Analysis Converting between English and Metric Units (11 of 20) Consider the metric/english math fact: 2.54 cm = 1 inch This can be used as the conversion factor : 2.54 cm 1 inch or 1 inch 2.54 cm You can convert 25.5 inches to cm in the following way: Given: 25.5 in 2.54 cm Desired:? cm 25.5 in x = cm 64.8 cm 1 in This is the required way to show your work. You have two jobs in this class, to be able to perform the conversions and to be able to prove that you know why the answer is correct. Temperature Scales (12 of 20) The important idea is that temperature is really a measure of something, the average motion (kinetic energy, KE) of the molecules. Does 0 C really mean 0 KE? nope... it simply means the freezing point of water, a convenient standard. We have to cool things down to C before we reach 0 KE. This is called 0 Kelvin (0 K, note: NO symbol.) For phenomena that are proportional to the KE of the particles (pressure of a gas, etc.) you must use temperatures in K. K = C C = K 273

4 Mass vs. Weight Theory, Measuring, Conversions (13 of 20) mass is the amount of something... weight is how much gravity is pulling on the mass. (Weight will be proportional to the mass at a given spot.) Mass is what we REALLY want to use... measured in grams. You use a balance to measure mass... you compare your object with objects of known mass. Weight is measured with a scale (like your bathroom scale or the scale at the grocery store). If there is no gravity, it doesn t work. Note: electronic balances are really scales! You convert mass / weight using: 1 kg lbs or lbs 1 kg Potential Energy (PE) and Kinetic Energy (KE) (14 of 20) You can calculate the KE of an object: KE = 1 2 mv2 m = mass, v = velocity [Note units: 1 J = 1 kg m 2 s 2 ] Temperature is a measure of the average kinetic energy. PE = the potential to do work which is due to an object s position in a field. For example, if I hold a book 0.5 m above a student s head it can do some damage m above her/his head, more work can be done. Important ideas: Objects tend to change from high PE to low PE (downhill). High PE is less stable than low PE. Mass, Volume, and Density Intensive vs. Extensive Properties (15 of 20) Extensive properties depend on the amount of substance. We measure these properties frequently... (mass & volume... mostly). Intensive properties are independent of the size of the sample. These are useful for identifying substances... (melting point, boiling point, density, etc.) mass It is interesting that an intensive property, density = volume is the ratio of two extensive properties... the size of the sample sort of cancels out. Be able to do density problems (3 variables) and know the usefulness of specific gravity. Calorimetry (16 of 20) Heat is the total KE while temperature is the average KE. A way to measure heat is to measure the temperature change of a substance... often water. It takes 1 calorie of heat energy (or J) to heat 1 gram of H 2 O by 1 C. The specific heat of water = 1 cal g C = J g C heat = specific heat x mass H 2 O x T You can heat other substances as well, you just need to know their specific heats. Notice that this is simply heating or cooling a substance, not changing its phase.

5 Physical and Chemical Properties Physical and Chemical Changes (17 of 20) Equations to symbolize changes: reactants products Physical Properties can be measured from a sample of the substance alone... (density, MP, BP, color, etc.) Chemical Properties are measured when a sample is mixed with another chemical (reaction with acid, how does it burn in O 2 ) Physical Changes imply that no new substances are being formed (melting, boiling, dissolving, etc.) Chemical Changes imply the substance is forming new substances. This change is accompanied by heat, light, gas formation, color changes, etc. Matter Energy Pure Substances Mixtures Pure Substances, Elements, & Compounds Homogeneous & Heterogeneous Mixtures (18 of 20) Elements Compounds Homogeneous Heterogeneous This chart should help you sort out these similar terms. Be able to use chemical symbols to represent elements and compounds. For example... CuSO 4 5H 2 O, a hydrate, contains 21 atoms & 4 elements. Memorize the 7 elements that exist in diatomic molecules: HONClBrIF or BrINClHOF or H and the 6 that make a 7 starting with element #7 Separating Mixtures by Filtration, Distillation, and Chromatography (19 of 20) Mixtures are substances the are NOT chemically combined... so if you want to separate them, you need to exploit differences in their PHYSICAL properties. Filtration: some components of the mixture dissolve and some do not. The filtrate is what passes through the filter. Distillation: some components vaporize at different temperatures or one component may not vaporize at all (e.g.: salt+water) complete separation may not be possible. Chromatography: differences in solubility vs. adhesion to the substrate. Substaratemay be filter paper (paper chromatography), or other substances, GLC, TLC, HPLC, column, etc. Early Laws: the Law of Definite Composition & the Law of Simple Multiple Proportions (20 of 20) Definite Composition: samples of the same substance from various sources (e.g. water) can be broken down to give the same % s of elements. Calculation: percent composition Multiple Proportions: samples of 2 substances made of the same 2 elements... (e.g. CO 2 & CO or H 2 O and H 2 O 2 or CH 4 and C 3 H 8 ) if you break down each to give equal masses of one element, the masses of the other element will be in a simple, whole-number ratio. Calculation: proportions to get equal amounts of one element and then simple ratios.

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