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CHEM 101 Chapter 1

by: Lyna Nguyen

CHEM 101 Chapter 1 Chem 101

Marketplace > Texas A&M University > Chemistry > Chem 101 > CHEM 101 Chapter 1
Lyna Nguyen
Texas A&M

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Chapter 1 Notes textbook + lecture
General Chemistry 1
Dr. Daniel Collins
Class Notes
Chemistry, Chem
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This 9 page Class Notes was uploaded by Lyna Nguyen on Friday January 29, 2016. The Class Notes belongs to Chem 101 at Texas A&M University taught by Dr. Daniel Collins in Fall 2015. Since its upload, it has received 21 views. For similar materials see General Chemistry 1 in Chemistry at Texas A&M University.

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Date Created: 01/29/16
09/01/15­09/08/15 Chemistry 101 Chapter 1  What is chemistry?  Definition: study of matter and changes it undergoes using energy  Is a central science  Physics: chemistry in motion  Bio: chemistry in a cell Engineering: application of sciences   Why is it hard? Terminology   Words, numbers, equations Skeptic: looks at all sides and chooses the best   Organic: contains carbon and hydrogen  All chemists: eat, drive a car, cook  Chemical intuition  Biggest challenge: getting standard units, terms, and numbers  Chemistry and its method  Parts of Scientific Method  Observation  Hypothesis  Data Collection  Record what you do; proof  Summary of data  Laws and theories  Hypothesis, Laws and Theories  Hypothesis: tentative explanation or prediction w/ knowledge  Requires both qualitative and quantitative in experiments  Quantitative: numerical data  Supports claims  Ex: mass, temp  Qualitative: non-numerical; observations  Ex: color, physical appearance  Law: a concise verbal or mathematical statement of a behavior or a relation that is the same in all conditions  Narrow statement 1 09/01/15­09/08/15  Theory: well-tested, unifying principle that explains a body of facts and the laws based on them  Binds observations/data together  Capable of suggesting new hypotheses  Based on carefully determined and reproducible evidence  “Umbrella” statement  Can change w/ new facts  Systematic Approach  Scientific method  Observation -> representation -> interpretation -> observation … (repeating cycles)  Goals of Science  What makes a good scientist?  Patience  Knowledge  Hard work  Luck  Goals:  Prediction and control  Can be dangerous  Understanding and Explaining  Examples  Dilemmas and integrity in science  Results can be inconclusive  Guidelines (4 rules):  Results should be reproducible  Research should be reviewed (detailed)  Conclusions should be reasonable and unbiased  Credits should be given where it is due  Sustainability and Green Chemistry  Began to take root 20+ years ago  New ways of doing things w/ lower pollution levels  Classifying matter  Matter  Anything that occupies space and has mass  Governed by 2 things: compositions and energy  States of Matter and Kinetic-Molecular Theory  State: gas, liquid, solid  At low temps, all matter is solid  As temp increases, liquids becomes gases  Kinetic-Molecular Theory of Matter 2 09/01/15­09/08/15  All matter consists of small particles  Solid: packed closely, regular in array, vibrate in avg. positions  Liquid: random arrangement, fluid, not confined to position  Gas: parts are far part, moves around, volume=container, fluid  There are net forces between particles in all states  Gas: small  Solid/liquid: large  Determines properties of matter  Higher temp = faster movement  Energy of motion (kinetic) acts to overcome the forces of attraction between parts  Energy causes change in state  Matter at the Macroscopic and Particulate Levels  Macroscopic: world of experiments and observations observed by the naked eye  Submicroscopic/particulate: only seen by microscope (atoms and molecules) Matte r Substance Mixtures s Heterogeneo Homogenous us  Pure Substances  Solutions: mixture of liquid (water) and dissolved substances  Has unique properties which it can be recognized  Melting/boiling point  Cannot be separated into 2+ species by any physical means  Physical means: used to separate a mixture into its pure components  Ex: distillation, using a magnet  Mixtures: Heterogeneous and homogenous 3 09/01/15­09/08/15  Mixture: 2 or more pure substances that can be separated by physical techniques  Retains specific identities  Heterogeneous: uneven texture, not uniform throughout  Ex: cement  Homogenous: 2+ substances in same phase  Ex: milk, soda, solder  Often called solutions  Uniform throughout  Elements: only 1 type of atom; cannot be separated by chemical means  118 elements  82 occur naturally  Ex: Au, Al, Pb, O, C, S  36 elements created by scientists  Atom: smallest particle of an element that retains the characteristics chemical properties  Compounds: pure substances composed of 2+ elements held together by chemical bonds at fixed proportions  Referred as chemical compounds  Can be changed by chemical means  Ions: electrically charged atoms or groups of atoms  Molecules: smallest discrete unit that retain the composition and chemical characteristics  Chemical Formula: represents compound  Ex: H 2 -> 2 Hydrogen atoms w/ 1 Oxygen  Physical Properties: properties that can be observed and measured w/o changing the composition of a substance  Ex: height, weight, color, conductivity  Density: ratio of mass to volume  Temperature: affects the numerical values of its properties  Chemical Properties: indicates whether and sometimes how readily a material undergoes a chemical change w/ another material  Ex: combustion, fire, food digestion  Ex Questions:  Sugar dissolving in water  Physical  Methane burning in air  Chemical  Extensive and Intensive Properties  Extensive: depends on amount of substance present; size dependent  Ex: mass and volume, energy, length 4 09/01/15­09/08/15  Intensive: do NOT depend on amount  Ex: density, temp, color  Useful in identifying material  Physical and Chemical Changes  Physical changes: changes in physical properties  Identity of substance is sustained  No new chemicals  Ex: melting, dissolving  Chemical changes: 1 or more substances (reactants) are transformed into 1 or more other substances (products)  Alter composition; “reaction”  Chemical equation: representation of chemical reaction  Left: reactants  Right: products  Ex: hydrogen burns air to form water  Energy: Some basic principles  Defined as the capacity to do work  Kinetic energy: motion  Potential energy: results from an object’s position or state and included height, springs, and chemical bonds  Ex: chemical, position  Types of energy  Radiant: comes from the sun/primary source  Thermal: associated w/ motion of atoms/molecules  Chemical: stored w/in bonds of chemical substances  Nuclear: stored w/in collection of neurons and protons in the atom  Can be converted  Conserved most of the time  Ex: chemical -> kinetic  Law of Conservation of Energy  Correct title: First Law of Thermodynamics  Thermodynamics: movement of energy  Defined as “energy can neither be created nor destroyed  Total energy of the universe is constant  Math Review: Measurements, quantitative  Measurements and importance:  Length: ruler (m)  Mass: balance (g, kg)  Volume: beaker, flask, bottle (oL,oL)  Temperature: thermometer ( F, C, K)  Critical Parts: number and unit 5 09/01/15­09/08/15  SI units: created by the General Conference of Weights and measurements; standard units  2 major parts: base unit and prefix  Base unit: tells us what it is  Prefix: tells us size of it  Mass vs. Weight  Mass: measure of what is there  SI: kilograms 3  1kg=1000g=1x10 g=2.21lbs  Weight: force gravity exerts on an object  Weight=c*mass  Volume 3  SI: cubic meter (m )  1 cm =(1x10 ) =1x10 m -6 3 3  1 cm =1mL=1cc  Density 3 3 3  Density = kg/m -> 1g/cm (solids) =1g/mL OR 1000kg/m (volume/liquids) =m/v  1g/L = .001 g/mL (gases)  Most universal  Intensive  Energy  SI: Joules 2 2  Mechanical energy: 1J=1 kg*m /s  Calories: old energy unit  1 calorie (cal) = 4.184 joules (J)  1 dietary cal = 1000 cals  Temperature  Celsius: o  Freezing point: 0 C  Boiling point: 100 Co o  Room temp: 20 C  Body temp: 37 C o  Kelvin  Absolute zero: lowest temperature o o  0 C = 273.15K = 32 F  100 C = 373.15K = 212 F o  Conversion  K= C+273.15  F=(9/5)x C+32 6 09/01/15­09/08/15  Making measurements  Precision: of measurement indicates how well several determinations of the same quantity agree  How close data is to each other  Multiple and reproducible results  Accuracy: agree of a measurement w/ the accepted value of the quantity  Correct value; “true value”  Expressed as percent error  Experimental Error  Error in measurement = experimentally determined value – accepted value  Percent error=(Error in measurement/accepted value)*100%  Standard deviation  Defined: equal to the square root of the sum of the square of the deviations for each measurement from the average, divided by one less than the number of measurements  Steps:  Average calculated  Difference b/w measurement and average  Determinate errors: caused by faulty instruments or human errors  Ex: incorrect record keeping  Indeterminate (random errors): uncertainties in a measurement  Mathematics in chemistry  Exponential or scientific notation  Fixed notation: full number (324)  Way of presenting very large/small numbers n  N x 10  N: digit term  Between 1-10; +/-  N: power  Integer; +/-  Significant figure rules  Any digit that is not 0 is significant  0 b/w 2 other sigfigs are significant  0 place holders are not significant  0 to the right of a nonzero number and to the right of a decimal are significant  In numbers less than 0, only zeros after nonzero sigfigs are significant  Significant Figures in Calculations  Ambiguous: prefer 1 sigfig; use scientific notation 7 09/01/15­09/08/15  Exact numbers: numbers from definition or numbers of objects are considered to have an infinite number of sigfigs  Ex: numbers of trials, data  Sigfig calculation rules  Rule 1: fewest decimal places when adding/subtracting  Rule 2: fewest sigfigs when multiplying/dividing  Rule 3: When rounded off, the last digit to be retained is increased by one inly if the following digit is 5 or greater  Truncation: chop off  Round off at end of calculations  Problem Solving by Dimensional Analysis  Dimensional analysis (factor-label): general problem-solving approach that uses dimensions/units  Steps  1) Determine which unit conversion factors are needed  Conversion factor: equivalence of a measurement in 2 different units  Number in original [new/old]=new unit  2) Carry units  3) Desired unit/target unit should be the only thing left  Given*conversion factor=desired factor  Graphs and Graphing  Standard straight line  Y=mx+b  DO NOT CONNECT DOTS  Problem Solving and Chemical Arithmetic  Step 1: state the problem  Step 2: what do you know?  Step 3: strategy (sketch, pictures, roadmap)  Step 4: solution (execute plan) REMEMBER UNITS  Step 5: think about answer  Step 6: check your understanding  Problem Solving Tips  Answer the question  Read the entire problem  Equations, tables, and unit analysis  Know how to get data  Time management  Is your answer reasonable?  Check calculator mistakes 8 09/01/15­09/08/15 9


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