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- StudyBlue
- North-carolina
- North Carolina School of Science and Mathematics
- AP Chemistry I Ch401/a
- Allen
- 1_Notecards_-_Basics.pdf

José Luis S.

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1 • Introduction 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?) 1 • Introduction 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. 1 • Introduction Significant Digits I What do they mean? (3 of 20) Consider: 16.82394 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 16.8239 cm, but some will say …94 cm while others might say …95 cm. 1 • Introduction Significant Digits II Some examples with rulers. (4 of 20) 1 2 a b c (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). 1 • Introduction 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. 5082 meters (4) 0.002008 L (4) • 0’s in the FRONT of a number are NEVER significant. 0.0032 kg (2) 0.00000751 m (3) • 0’s at the END of a number are SOMETIMES significant. • Decimal point is PRESENT, 0’s ARE significant 2.000 Liters (4) 0.000500 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 1 • Introduction 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 = 5.392 x 103 0.000328 = 3.28 x 10–4 1.03 = 1.03 550 = 5.5 x 102 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: 0.000328 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 102 (3) 1 • Introduction Significant Digits IV Significant Digits in Calculations (7 of 20) 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: 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 = 29.97 ≈ 30. cm2 (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 + 8.1 cm = 11.8 cm (3 sig. digits) Know how to ilustrate why these rules work. 1 • Introduction 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. 1 • Introduction 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- 1/100th milli- 1/1000th kilo- 1000 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 1 • Introduction % 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 = 17.65 100 5 31 = x 100 = 16.13 100 so… we can see that 317 > 531. There are 17.65 parts per 100 (Latin: parts per centum) or 17.65 percent (17.65 %)… the % is a “1 0 0” ppm (parts per million) is the same idea, (use 1,000,000 instead of 100) 317 = x1 000 000 = 176,470 ppm 1 • Introduction 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 cm1 inch or 1 inch2.54 cm You can convert 25.5 inches to cm in the following way: Given: 25.5 in Desired: ? cm 25.5 in x 2.54 cm1 in = 64.77 cm ≈ 64.8 cm 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. 1 • Introduction 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 –273.15°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 + 273 °C = K – 273 1 • Introduction 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 kg2.205 lbs or 2.205 lbs1 kg 1 • Introduction Potential Energy (PE) and Kinetic Energy (KE) (14 of 20) You can calculate the KE of an object: KE = 12mv2 m = mass, v = velocity [Note units: 1 J = 1 kg·m2·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... 1.0 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. 1 • Introduction 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.) It is interesting that an intensive property, density = massvolume 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. 1 • Introduction 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 4.184 J) to heat 1 gram of H2O by 1 °C. The specific heat of water = 1 calg·°C = 4.184 Jg·°C heat = specific heat x mass H2O 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. 1 • Introduction 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 O2) 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. 1 • Introduction Pure Substances, Elements, & Compounds Homogeneous & Heterogeneous Mixtures (18 of 20) Pure Substances Matter Energy CompoundsElements Mixtures HeterogeneousHomogeneous This chart should help you sort out these similar terms. Be able to use chemical symbols to represent elements and compounds. For example... CuSO4•5H2O, 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” 1 • Introduction 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. 1 • Introduction 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. CO2 & CO or H2O and H2O2 or CH4 and C3H8) 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. Paul Groves APCh01.pdf Chapter 1

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