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Chem 1061 Principles of Chemistry I Andy Aspaas, Instructor

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About this class Prereq. to CHEM 1062 More detail in underlying concepts than 1020 Expected that you already have basic familiarity with chemistry concepts Syllabus

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Chapter 1 Should be review, but still read carefully Why study chemistry? –Practical applications –Intellectual development –Crossover to other fields

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The scientific method Problem or question: what does the scientist wish to accomplish in research? Law: a fundamental consistent observation Experimentation: controlled observations, aim is to get consistent results that can be replicated and explained Hypothesis: a tentative explanation Theory: a tested explanation

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Matter and mass Matter is anything that occupies space Mass: a measure of the amount of matter in a sample –Only measured with a balance Weight: the gravitational force a sample exerts –Measured with a scale Mass is always conserved in a chemical transformation (reaction)

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Physical states of matter Solid: rigid material; made of closely packed, relatively motionless particles Liquid: mostly incompressible fluid; particles are in close proximity but move freely –(fluids flow easily and change shape to occupy the space of their vessel) Gas: compressible fluid; particles are distant and fast-moving

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Changes Physical change: form of the matter is changed, but not its chemical identity –Phase changes: melting, freezing, boiling, evaporation, condensation, sublimation –Dissolving, like Kool-Aid into water Chemical change: matter is changed into a different kind of matter

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Types of matter Substances: cannot be separated into other types of matter by physical processes –Elements: cannot be decomposed into simpler substances by chemical changes –Compounds: composed of 2 or more chemically combined elements

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Types of matter Mixtures: can be separated by physical means into other substances –Heterogeneous: has parts with different observable physical distinctions –Homogeneous: uniform in properties throughout, a.k.a. solutions

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Physical measurements Measurement: comparison of physical quantity with a fixed standard (a unit of measurement) Precision: how similar a number of measurements are Accuracy: how close a single measurement is to its true value

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Significant figures The number of meaningful digits in a measurement, indicating precision of measurement –All nonzero digits are significant –Zeroes at the beginning of a number are not –Terminal zeroes are if there’s a decimal point –Terminal zeroes with no decimal point usually are not –Exact quantities have an infinite number of sig figs

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Using sig figs x or ÷: answer has same number of sig figs as measurement with least number of sig figs + or -: answer has same number of decimal places as measurement with least number of decimal places Exact numbers do not change number of sig figs or decimal places in answer Round answer to give correct sig figs or decimal places, do not round intermediate calcs

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SI units Metric system uses SI units of measurement –Easy conversions between units –All decimal based –Power of 10 prefix

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SI base units Units from which all other SI units can be derived MeasurementUnitSymbol Lengthmeterm Masskilogramkg Timeseconds TemperaturekelvinK Quantitymolemol Electrical currentampereA Luminositycandelacd

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SI prefixes MultiplePrefixSymbol 10 6 megaM 10 3 kilok 10 -1 decid 10 -2 centic 10 -3 millim 10 -6 microµ 10 -9 nanon 10 -12 picop

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Length, mass, and time Meter (m) is SI base unit of length –Nanometer (nm, 10 -9 m) and Å (angstrom, 10 -10 m) commonly used by chemists Kilogram (kg): mass, about 2.2 pounds Second (s): time - many of the fastest chemical reactions are finished in picoseconds (10 -12 s)

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Temperature A measurement of the consequences of heat energy on a substance Not a measurement of heat itself Celsius scale: most common scientific use Kelvin (K): SI base unit, lowest theoretical possible temperature is zero (absolute zero) Fahrenheit scale: used in US Conversions

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Derived units Obtained by mathematically combining 2 or more other SI units Volume: length cubed. SI unit is m 3 –Traditional unit is liter (L), equal to 1 dm 3 –1 milliliter (mL) = 1 cm 3 Density: mass per unit volume. –Common unit is g/cm 3 –Important characteristic of a substance

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Dimensional analysis Systematic method of unit conversion Start with a single known quantity Multiply by conversion factors –An equality written as a division of two units –Arrange conversion factors so units cancel –Unit of answer is apparent through calculation –Estimate to make sure answer is reasonable

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Chapter One: CHEMICAL FOUNDATIONS. Copyright © Houghton Mifflin Company. All rights reserved.Chapter 1 | Slide 2 Chemistry: An Overview A main challenge.

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