CHEM 1030, Week 1/2 Notes
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This 5 page Class Notes was uploaded by Alyssa Anderson on Tuesday January 19, 2016. The Class Notes belongs to CHEM 1030 at a university taught by Dr. Streit in Spring 2016. Since its upload, it has received 44 views.
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Date Created: 01/19/16
Alyssa Anderson Chemistry Notes Weeks 1 and 2 1/13/16-1/22/16 Chemistry- the study of matter and changes that matter undergoes Matter- anything that has mass and occupies space Scientiﬁc Method- a procedure/set of guidelines to organize and publish efﬁciently 1. Gather data through observations and experiments 2. Identify patterns and trends in collected data and note any initial thoughts 3. Summarize ﬁndings with a law- a concise statement that makes a relation between phenomena 4. Formulate a hypothesis by observing the cause and effect relationship 5. With time, the hypothesis may evolve into a theory which can predict future occurrences Chemists classify matter as either a substance or a mixture of substances 1. Substance- a form of matter that has a deﬁnite composition and distinct properties - example: salt (NaCl), iron, water (H20), mercury, carbon dioxide (CO2) - substances differ from each other in composition and may be identiﬁed by appearance, taste, smell, etc. 2. Mixture- physical combination of 2 or more substances a. homogenous- uniform throughout solution - example: seawater, apple juice, cake b. heterogenous- not uniform throughout solution - example: trail mix, chicken noodle soup, shells in sand Classiﬁcation of Matter 1. Solid- particles are held close together in an ordered position and DO NOT conform to the container it is placed in 2. Liquid- particles are held relatively close together but do not have an organized pattern and DO conform to the the container it is placed in 3. Gas- particles are far apart and have no set pattern but DO conform to the container it is placed in 4. In principle, all substances can exist in the solid, liquid, or gaseous stage 5. We can convert a substance by changing its identity 6. Mixtures can be separated by physical means into its component without changing the identities of the components - example: magnet to separate sand and iron (iron is magnetic) - example: boil water to separate salt and water (water has a much lower boiling point) - example: boil water to separate water and alcohol (different boiling points) Properties of Matter 1. Quantitative- properties measured/expressed with a number/unit (QUANTITY) 2. Qualitative- properties measured without measurements but rather are based on observations using the senses (taste, color, smell, etc.) (QUALITY) 3. Physical Property- one that can be observed or measured without changing the identity of the substance - example: color, melting point, boiling point 4. Chemical Property- one that a substance exhibits as it interact with another substance - example: ﬂammability, corrosiveness, rust 5. Physical Change- change where the state of matter changes but the identity of the matter does not change - example: changes of state (melting, boiling, freezing, condensing) 6. Chemical Change- change in the composition so that the original composition no longer exists - example: digestion, combustion, oxidation 7. Extensive Property- depends on the amount of matter present - example: mass, volume, aka additive properties 8. Intensive Property- does NOT depend on the amount present - example: temperature, density 9. Physical Process- mixtures are separated but the identities do not change 10. Chemical Process- a process of changing mixtures/chemicals Scientiﬁc Measurement 1. Properties that can be measured are called quantitative 2. A measured quantity must always include a unit 3. Systems A. English- foot, gallon, pound, Fahrenheit B. Metric- meter, liter, kilogram C. International System of Units (SI units)- universally used by scientists 1. Meter 2. Kilogram 3. Kelvin 4. Second 5. Ampere 6. Mole 7. Candela 4. Mass (g or kg or amu) A. a CONSTANT measure of amount of matter in an object/sample B. Gravity varies from location to location constant so weight = mass x gravity C. the mass of an atom is 1 amu= 1.6605378 x 10^-24 g 5. Temperature (Celsius or Kelvin or Fahrenheit) A. Celcius- for water, freezing point is O*C, boiling point 100*C B. Kelvin (*SI UNIT*)- “absolute” scale because 0 K is the absolute lowest C. K = *C + 273.15 OR C* = K - 273.15 D. Fahrenheit- for water, freezing point is 32*F and boiling point is 212*F E. *F = (9/5)(*C) + 32*F OR *C = (5/9)(*F - 32) 6. Volume (meter^3 or Liter) A. V = (length)^3 B. 1 dm^3 = 1 L C. 1 cm^3 = 1 mL 7. Density (kg/m^3) A. d = mass/volume so d = mass/length^3 B. solid = g/cm^3 C. liquid = g/mL D. gas = g/L E. example: if d1>d2 then m1<m2 OR v1>v2 Uncertainty in Measurements 1. Exact Numbers- deﬁned values or counted numbers - example: 1 kg = 1000 g - example: 1 dozen = 12 items - example: 28 students in a class 2. Inexact Numbers- measured by anything but counting such as length, volume, mass A. It must be reported to indicate uncertainty by using signiﬁcant digits B. The last digit reported is called the uncertain digit C. example: if we have an item against a ruler and we think it’s about 2.5 inches long, we know it’s for sure 2 inches but not sure about the .5, so we put 2.5 +/- 0.1 inch, and with a more accurate ruler we could put 2.45 inches +/- 0.01 inch D. Guidelines of Signiﬁcant Figures 1. Any nonzero numbers ARE signiﬁcant 2. Zeros between nonzero numbers ARE signiﬁcant 3. Zeros to the LEFT of the ﬁrst nonzero digit are NOT signiﬁcant 4. Zeros to the RIGHT of the nonzero digits in decimals ARE signiﬁcant 5. Zeros to the RIGHT of the last nonzero digit in a number without a decimal MAY OR MAY NOT be signiﬁcant - example: 100 could have 1 2 or 3 signiﬁcant ﬁgures Calculations with Measured Numbers 1. Addition/Subtraction- line up the decimals and take the answer with the smaller amount of digits (rounding may be necessary) 2. Multiplication/Division- preform the action then take the fewer amount of digits from the original numbers given (rounding may be necessary) 3. NOTE: Be sure not to include exact numbers, such as the counted number - example: when ﬁnding the mass of each of 2 pennies, knowing together they equal 15 grams, 2 is not included in the measurement of signiﬁcant ﬁgures. Therefor, since together they had 15 grams and that is 2 signiﬁcant ﬁgures, your answer will have 2 signiﬁcant ﬁgures 4. Rounding A. Leave rounding for the LAST step. DO NOT ROUND AFTER EACH STEP B. If the last digit is greater than 5, round UP (ex: 318.175 = 318.18) C. If the last digit is less than 5, round DOWN (ex: 318.174 = 318.17) 5. NOTE: Be aware of powers of 10. Make sure that you are calculating variables with the same power, then proceed 6. NOTE: Signiﬁcant ﬁgures matter even when scientiﬁc notation changes 7. NOTE: Be sure to calculate the right mass or volume before proceeding to ﬁnd density or weight 8. Accuracy- how close the measurement is to the TRUE value 9. Precision- how close a series of measurements are to one another Using Units and Solving Problems 1. Conversion Factor- fraction in which same quantity is expressed one way in the numerator and another in the denominator - example: 1 inch = 2.54 cm aka 1 in/2.54 cm OR 2.54 cm/1 inch 2. Dimensional Analysis- use of conversion factors in problem solving A. Also known as the factor-label method B. example: convert 12 inches to meters (NOTE: only use signiﬁcant ﬁgures of the thing you are converting (so 2 s.f. because 12 inches has 2 s.f.); 12 inches x 2.54 cm/1 inch x 1m/100 cm = 0.3048 m = 0.30 m
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