CH 221 Reading Notes: Chapter 1
CH 221 Reading Notes: Chapter 1 CHEM 221
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This 10 page Class Notes was uploaded by Adeline Fecker on Tuesday September 27, 2016. The Class Notes belongs to CHEM 221 at University of Oregon taught by Deborah Exton in Fall 2016. Since its upload, it has received 55 views. For similar materials see Chemistry 221 in Chemistry at University of Oregon.
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Date Created: 09/27/16
September 28th, 2016 CH 221 Reading Notes: Chapter 1 Keys to the study of Chemistry 1.1 Some fundamental definitions -Chemistry: the study of matter and its properties, the changes that matter undergoes, and the energy associated with those changes -Composition: the types and amounts of simpler substances that make up matter -States of Matter: 3 physical forms Solid (fixed, rigid shape, does not conform) Liquid (varying shape, conforms to container, has upper surface) Gas (varying shape, fills entire container, no surface) 1 Rectangular Randomly arranged Randomly arranged 3D shape Move fluidly Move freely Close together Close together Spread out Clear Disorganized Disorganized organization -Properties of Matter: the characteristics that give each substance its unique identity Physical properties- characteristics a substance shows by itself without changing into or interacting with another substance Ex. Ice (solid)-------Water (liquid) Chemical properties- characteristics a substance shows as it changes into or interacts with another substance(s) Ex. Water + electric current = hydrogen + oxygen 1 http://www.sparknotes.com/testprep/books/sat2/chemistry/chapter5.rhtml 1 -Physical change: No change in composition A physical change occurs with a substance alters its physical properties not its composition -Chemical change/Chemical reaction: A change in composition A substance is converted into a different substance *can be affected by temperature and pressure *a physical change caused by heating can generally be reversed by cooling Sample Problem 1.1 Visualizing change on the atomic scale Distinguish between physical and chemical changes by looking at a picture of atoms Sample problem 1.2 Distinguish between physical and chemical change Examples of naturally occurring chemical and physical changes The Central Theme in Chemistry We study observable changes in matter to understand their unobservable causes -Macroscopic-scale: properties and behavior we can see -Atomic-scale: properties and behavior we can’t see -Energy: The ability to do work Potential energy- The energy due to the position of the object relative to other objects Kinetic energy- The energy due to the motion of an object Potential energy + kinetic energy = total energy the object possesses Central concepts 1. When energy is converted from one form to the other, it is consumed not destroyed 2. Situations of lower energy are favored over situations of high energy Example 1: Gravitational system 1. As you lift the weight the potential energy increases because it is working against gravity. 2. When you drop the weight the potential energy is converted to kinetic energy. 2 3. Once the ball hits the ground it is at a lower potential energy (more stable) 2 Example 2: Springs 1. Stretching the spring increases potential energy because you are exerting energy onto the spring 2. When you release the spring the potential energy is converted to kinetic energy 3. The spring is more stable relaxed than stretched 3 Example 3: Oppositely charged 4 particles (Electrostatic force) 1. Opposite charges attract each other therefore potential energy increases when you force opposite particles away from each other 2. Potential energy converts into kinetic energy when opposite particles are pulled together. 3. Energy is more stable when opposite particles are closer SUMMARY: ❖ Chemists study the composition and properties of matter and how they change ❖ There are 3 states: solids, liquids, gases ❖ Each substance has unique physical and chemical properties and can undergo physical and chemical changes ❖ Changes in matter are accompanied by changes in energy ❖ Potential energy is due to position, kinetic energy is due to motion Chemical Arts and the Origins of Modern Chemistry 1.2 Prechemical Traditions 2https://www.youtube.com/watch?v=C1w_hL6mag 3 Potentialenergysprings.jpg 4 519fb817ce395fff0a000000.png 3 1. The alchemical tradition - 1st century AD - Dominated thinking for over 1500 years - Greek idea “matter strives for perfection” - Obsessed with converting metal - Invented distillation, percolation, and extraction and devised apparatus still used today - Encouraged observation and experimentation 2. The medical tradition - 13th centruy AD - Roots, herbs, and other plant matter used as medicine - Body is a chemical system and imbalance can be restored by treatment with drugs - Alliance between chemistry and medicine forms 3. The technological tradition - Contributed to people’s experience with the materials - Renaissance books described chemical processes - Introduced quantitative measurements The Phlogiston Fiasco and the Impact of Lavoisier -Combustion: the process of burning with the phlogiston theory -Phlogiston theory: an undetectable substance released called Phlogiston is released when material burns Critics: Why is air needed for combustion? Why does charcol stop burning in a closed vessel? Supporters: Air attracts phlogiston out of charcoal and burning stops when air is saturated with phlogiston Critics: How can the loss of phlogiston cause a gain in mass? Supporters: Phlogiston has negative mass Issue was resolved by Antoine Lavoisier (1743-1794) who did several experiments: 1. Heating mercury calx: total mass equaled starting mass of the calx 2. Heating mercury with gas reformed calx: total mass remained constant 3. A burning candle placed in the remaining air was extinguished - Named the gas oxygen and the metal calx were named metal oxides - Oxygen is a component of air and combines with a substance when it burns - In closed containers, the combustible substance stops burning when it has binded with all the oxygen - Metal calx weighs more than metal because it’s mass includes the mass of the oxygen - Successful because it was quantitative, and reproducible measurements SUMMARY: ❖ Alchemy, medicine, and technology focused on mystical explanations or practical experience but they contributed to the apparatus and methods that are still important 4 ❖ Lavoisier overthrew phlogiston theory by showing that oxygen is a component of the air and required for combustion 1.3 The Scientific Approach: Developing a Model - Observations: Facts our ideas must explain - Quantitative observations (data) can reveal trends - Summarized in mathematical terms as a natural law - Hypothesis: a proposal to explain an observation - Must be testable by experiment - If the results do not support it the hypothesis must be discarded - Experiment: A set of procedural steps that tests a hypothesis - Model/Theories: distinguishes scientific thinking from speculation, formed from experiments 1.4 Measurement and Chemical Problem Solving General Features of SI Units -Fundamental units/base units: identified with a physical quantity -Derived units: a combination of the seven base units (ex. m/s) 5 5 http://www.flinnsci.com/store/Scripts/prodView.asp?idproduct=16644 5 -Mass: the quantity of matter an object contains (kg) Mass is constant because an object’s quantity of matter cannot change Weight is variable because it depends on the local gravitational field Units and Conversion Factors in Calculations All measured quantities consist of a number and a unit Make sure to include units in calculations -Conversion factor: ratios used to express a quantity in different units, equal to 1 Multiply by the inverse to cancel out the unwanted unit (desired unit/undesired unit) 10cm/1m is equal to 1 therefore the quantity remains the same Systematic Approach to Solving Chemistry Problems 1. Problem: State all the information you need to solve the problem 2. Plan: think about the problem before acting a. Clarify known and unknown b. Think of necessary steps c. Plan procedure 3. Solution: Show calculations in the same order as the plan 4. Check: Verify your answer makes sense, correct units 5. Comment: Provide an application, alternative approach, or overview 6. Follow up Problems: Present similar problems with similar concepts 7. Some similar problems: Other practice problems Sample Problem 1.3 Converting units of length, volume, mass, raised to a power Conversion factors Conversion to SI -Density: mass/volume Isolate these factors by treating density as a conversion factor Density is a characteristic physical property so always has a specific value Sample Problem 1.7 Calculating Density from Mass and Volume Converting mass Converting lengths 6 Calculating density Temperature Scales -Temperature: a measure of how hot or cold an object is relative to another -Heat: Energy that flows from the object with the higher temperature to the object with the lower temperature. -Thermometer: A tube containing a fluid that expands when heated In hot substance: Heat flows from substance into the fluid In cold substance: Heat flows from the fluid from substance 3 Temperature Scales: Celsius, Kelvin (SI unit), Fahrenheit -Freezing point of water: 0 ० C, 273.15 ०K -Boiling point of water: 100 ० C, 373.15 ०K F to C= 9/5 C to K = +273.15 -Absolute Zero: 0 ०K, -273.13 ० C (therefore all Kelvin is positive) Sample Problem 1.8 Converting Units of Temperature Fahrenheit to Celsius Celsius to Kelvin Extensive and Intensive Properties -Intensive properties: independent of the amount of substance -Extensive properties: dependent on the amount of substance SUMMARY: ❖ SI unit system consists of seven base units ❖ The mass of an object (the quantity of matter in it) is constant ❖ A measured quantity consists of a number and a unit ❖ Follow the problem solving approach ❖ Density is a constant ❖ Temperature is not the same as heat ❖ Extensive properties are mass volume and energy ❖ Intensive properties are density and temperature 1.5 Uncertainty in Measurement: Significant Figures -Uncertainty: We can never measure an exact quantity, every measurement has uncertainty 7 The measuring device we choose depends on how much uncertainty is acceptable Example: Mass on a scale is measured in 0.1kg increments therefore the uncertainty is ±0.1kg 1. The smaller the uncertainty the more exact the measurement 2. Always estimate the rightmost digit of the measurement 3. Uncertainty can be expressed using ± -Significant Figures: Digits recorded both uncertain and certain 1. The greater number of Sig figs the greater certainty of a measurement Determining which Digits are Significant 1. Make sure the measurement has a decimal point 2. Move left to right until you reach the first nonzero digit 3. Count that digit and every digit to its right as significant. Example: 1.1300--------- 5 significant figures 6500.----------4 significant figures 6500---------- 2 significant figures 5.30 x 10₃-----3 significant figures Sample Problem: Determining the number of significant figures Significant Figures: Calculations and Rounding off -Rounding off: When there are too many significant figures 1. The least certain measurement sets the limit on certainty for the entire calculations and answer 2. If the digit removed is greater than 5, preceding number increases by 1 3. If the digit removed is less than 5, the preceding number remains the same 4. If the digit removed is 5 a. If preceding is odd, it increases by 1 b. If preceding is even, it remains the same 5. Round off the final answer only RULES: Arithmetic Operations Multiplication and Division: - Number of sigfigs in answer are same as measurement with the fewest sigfigs - Use rounding to reduce number of significant figures Example: 9.2cm x 6.8cm x 0.3744cm = 23.4225cm₃= 23cm₃ (2 sigfigs) Addition and Subtraction: - Number of decimal places in answer matches measurement with fewest decimals Example: 83.5mL + 23.28mL = 106.78 = 106.8 (1 decimal place) 8 Significant Figures in the Lab - Measuring device determines number of significant figures obtainable - More significant figures is means certainty - Choose more exact measuring tools to avoid heavily rounding off the final answer -Exact Numbers: No uncertainty associated with them Example: There are exactly 60 minutes in 1 hour *Exact numbers do not limit the number of sigfigs* Precision, Accuracy, and Instrument Calibration -Precision: reproducibility, how close measurements in a series are to each other (low random error) -Accuracy: how close are measurements to the actual value (low systematic error and random error) -Error: 1. Systematic Error: Part of the experimental system, often caused by faulty device or mistakes in the reading (values are all higher or all lower) 2. Random Error: Always occurs but depends on instrument precision and observer skill (values are all higher and all lower) -Calibration: Comparing device with known standard Example: setting scale to zero before taking mass SUMMARY: ❖ All measurements have some uncertainty which expressed by sigfigs ❖ Answer contains as many significant figures as least certain measurement ❖ Excess digits are rounded off in the final answer ❖ Lab device depends on certainty needed ❖ Precision is how close values are to each other, accuracy is how close values are to actual ❖ Systematic errors and random errors affect precision and accuracy ❖ Systematic errors can be compensated for by calibration Review Define features of the 3 states of matter 1.1 Distinguish between chemical and physical properties and changes 1.1, 1.2 Understand conversion of potential to kinetic energy 1.1 List the steps in the Scientific Model (observation, hypothesis, etc…) 1.3 Distinguish between extensive and intensive properties (mass, 1.4 density, etc...) Calculate number of significant figures in calculations 1.5 9 Distinguish between accuracy and precision 1.5 Use conversion factors 1.3 Find density from mass and volume, convert temperature (F to C to 1.4 K) 10
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