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Chem Week 1 Notes

by: Kyla Tovar

Chem Week 1 Notes CHEM 111-003

Kyla Tovar
GPA 3.6
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About this Document

These notes cover material discussed in the book as well as each of the lectures this week.
General Chemistry I
MacFarland, Kerry Jane
Class Notes
atoms, Chemistry




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This 9 page Class Notes was uploaded by Kyla Tovar on Saturday August 27, 2016. The Class Notes belongs to CHEM 111-003 at Colorado State University taught by MacFarland, Kerry Jane in Fall 2016. Since its upload, it has received 23 views. For similar materials see General Chemistry I in Chemistry at Colorado State University.


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Date Created: 08/27/16
Chemistry Reading 1.1, 1.3-1.6 1.1 States of Matter:  All things physically real is matter  Matter: anything that has mass and occupies space  Chemistry: the study of the composition, structure, and properties of matter in an object  Matter exists in three phases or physical states: solid, liquid, or gas  Atomic theory: all forms of matter are composed of microscopic particles  Molecules: a collection of atoms chemically bonded together  Chemical bonds: a force that holds two atoms in a molecule of a compound together  Solid: free packed, vibrate a little  Liquid: randomly ordered, close  Gas: separated and vibrate  Matter can be transformed from one physical state to another as its temperature is raised or lowered  Sublimation: transformation of a solid directly into a vapor (gas)  Deposition: transformation of a vapor (gas) directly into a solid 1.3 Classes of Matter  The two principal classes of matter are pure substances and mixture  Pure substance: matter that cannot be separated into simpler matter by a physical process  Pure substances cannot be separated into simpler forms of matter by any physical process  Physical process: a transformation of a sample of matter, such as a change in its physical state, that does not alter the chemical identity of any substance in the example  Pure substances are two groups: elements an d compounds o Elements: a pure substance that cannot be separated into simpler substances by any chemical process  Only a few found in nature in their elemental states o Compound: a pure substance that is composed of two or more elements linked together in fixed proportions and that can be broken down into those elements by chemical process  Chemical formula: a notation for representing the elemental composition of a pure substance using the symbols of the elements; subscripts indicate the relative number of atoms of each element in the substance  Low of constant composition: the principle that all samples of a particular compound always contains the same elements combined in the same proportions  Mixtures: a combination of pure substances in variable proportions in which the individual substances retain their chemical identities and can be separated from one another by a physical process o Classified as homogeneous or heterogeneous  Homogeneous mixture: a mixture in which the components are distributed uniformly throughout and have no visible boundaries or regions  Also called solutions o Often liquids but may also by solids of gases  Heterogeneous mixture: a mixture in which the components are not distributed uniformly, so that the mixture contains distinct regions of different composition  Induct immiscible liquids  Combinations of liquids that are incapable of mixing with, or dissolving in each other  Nearly all forms of matter we encounter are mixtures  Mixtures can be separate  Distillation: a process using evaporation and condensation to separate a mixture of substances with different volatilities o volatility: a measure of how readily a substance vaporizes  inversely proportional to the strength of the interactions between its particles  filtration: a process for separating solid particles from a liquid or gaseous sample by passing the sample through a porous material that retains the solid particles  can be soaked in acetone to dissolve some of the compounds present  chromatography: a process involving stationary and mobile phases for separating a mixture of substances based on their different affinities for the two types of phases 1.4 Properties of matter  Intensive property: a property that is independent of the amount of substance present o Distinctive color, malleable, ductile, melts  Extensive property: a property that varies with the amount of substance present o Length, width, mass, and volumes  Properties of Substances o Physical property: a property of a substance that van be observed without changing the substance into another substance  Ex: density which is the ratio of the mass (m) of an object to its volume  D=m/v o Chemical properties: a property of a substance that can be observed only by reacting the substance chemically to form another substance  Does it react? How fast? 1.5 Atomic Theory: The Scientific Method of Action  Ancient Greeks believed in atoms but couldn’t test it  Laws and Theories o Scientific method: an approach to acquiring knowledge based on an observation of phenomena, development of testable hypothesis, development of a testable hypothesis, and additional experiments that test the validity of the hypothesis o One of the earliest descriptions published in 1620 by Francis Bacon o Scientific Law: verbal or in the form of an equation, is concise and generally applicable o Law of definite proportion: the principle that compounds always contain the same proportions of their component elements o Joseph Louis Proust research of the composition of compounds led him to conduce that compounds always contain the same proportions of their component elements  Law of definite proportions o Scientific theory: a general explanation of widely observed phenomena that has been extensively tested o Scientific laws/ theories describe natural phenomena and relationships o Scientific theories explain why phenomena happen or why relationships are true o Hypothesis: a tentative and testable explanation for an observation or a series of observations o A scientific theory explaining Proust’s law of definite proportion was developed by John Dalton o Law of multiple proportions: the principle that, when two masses of one element react with a given mass of another element to form two different compounds, the two masses of the first element have a ratio of two small whole numbers o Dalton concluded element combine this way because they are composed of atoms 1.6 Molecular View  Some pure elements also exist as molecules  Diatomic molecules: 2 atoms O2  Molecular formula: a chemical formula that shows how many atoms of each element are In one molecule of a pure substance  Structural formula a representation of a molecule that uses short lines between the symbols of elements to show chemical bonds between atoms  Condensed structural formulas: symbols of elements appear In a pattern that shows how the atoms are arranged relative to one another  Ball-and-stick molecules provide three dimensional views of molecules  Space filling models now the atoms are arranged and its overall three- dimensional shape 1.1, 1.3-1.6 (Lecture Notes) 1) Physical States of Matter  Matter: has mass and occupies space o Consists if atoms or molecules o Molecules: 2 or more atoms joined together  3 Physical States of Matter o solid (s): definite volume and shape, close together, locked into position, vibration o liquid (l): definite volume but not a definite shape close together but moving around o gas (g): neither definite volume or shape, far apart and vibrate and run around  you can compress gas  Phase Transitions o Changing phases require absorbing or releasing energy o Solid to liquid is called melting o Liquid to solid is called freezing o Liquid to gas is called vaporization o Gas to liquid is called condensation o Solid to gas is called sublimation o Gas to solid is called depositions 2) Classes of Matter  Pure substance o Element or compound o Has constant composition o Consists of one type of atom or molecule o Elements: have 1 type of atom o Compound: 2 or more elements in fixed proportions  Mixture o 2 or more substances in variable proportions o uniform homogeneous  ex: tea, air, seawater o non-uniform heterogeneous  ex: chocolate chip cookie, salad dressing o can be separated by physical process o changing a subscript makes a different compound 3) Chemical and Physical Properties/Processes  chemical process: observe by changing the composition o ex: burning, reacting, iron rusting  Physical process: observe without changing composition o Ex: boiling (point), density, vaporization 1.7: Coast: A Framework for Solving Problems  COAST o Collect and Organize, Analyze, Solve, and Think about the answer  Collect and Organize o Sorting through the info given o Identify and relevant info o Identify the key concepts o Define the key terms o Sort through the info o Assemble any supplemental info  Analyze o Relate it to the answer o Work backwards  Solve o Directly from your analysis of the problem  Think about it  Solve problems in a logical way 1.7 Making Measurements and Expressing the Results  Meter: the standard unit of length, names after the Greek Metron, which means measure  1960 used French metric system of units  joule: the SI unit of energy, equivalent to 1 kg  Precision and Accuracy o There is a limit to how well we can know the results o These terms used to describe how well a measured quantity is known o Precision: the extent to which repeated measurements of the same variable agree o Accuracy: agreement between an experimental value and the true value  Significant Figures o Significant figures: all the certain digits in a measured value plus one estimated digit, the greater the number of significant figures, the greater the certainty with which the value is known  Significant Figures in Calculations o Weak-link principle: the chain is only as strong as its weakest link o The weak link principle for significant figures also applies to calculations requiring addition or subtraction o Measurements always have some degree of uncertainty, which limits the number of significant figures we can use to report any measurement 1.9 Unit Conversions and Dimensional Analysis  Conversion factor: fraction in which the numerator is equivalent to the denominator but is expressed in different units, making the fraction equivalent to 1 1.10 Temperature Scales  Absolute zero (O K): zero point on Kelvin temperature scale; theoretically the lowest temperature possible  Kelvin (K): the SI unit of temperature  T(K)= T(C) +273.15  T(C): 5/9 (T(F) - 32)) Lecture Notes  Molecular view o Molecular formulas  Ex: C3H6O o Structural Formulas o Condensed Structural Formulas o Ball and Stick Models  Can see bonds o Space-filling Models  More realistic, however you can’t see bonds 1) SI Units and Prefixes a. Need to know i. 1 nm=10^-9 m ii. 10^-6 m iii. 1 mm= 10^-3 m iv. 1 cm= 10^-2 m v. 1 km= 10^3 m vi. 1 cm^3= 1mL b. these base units can be any form of measurement i. will have same prefixes in every measurement c. Don’t worry about conversions, they will be given to you 2) Uncertainty in Measurements a. Precision and accuracy i. precision is how close repeated measurements are to each other ii. accuracy is how close the experimental measurement is to the real value b. if you don’t know the answer it cannot be 100% accurate c. Significant Figures i. All significant except 0 whose job is to hold the decimal place ii. Think about if someone has measured them or estimated iii. 2 rules of sig figs in calculations 1. multiply/divide a. same number of sig figs as the measurement with the fewest sig figs 2. Addition/subtraction a. Same number of decimal places as the measurement with the fewest decimal places b. Always write units in an equation 3. When there are multiple steps, keep track of sig figs as you go along 4. Exact numbers have an unlimited number of sig figs a. Exact numbers are i. Counter ii. Same exact conversion 3) Problem Solving and Unit Conversions a. If units raised to power the base must be raised as well


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