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Ch 100: Fundamentals for Chemistry

Ch 100: Fundamentals for Chemistry. Chapter 1: Introduction Lecture Notes. What is Chemistry?. Chemistry is considered to be the central science Chemistry is the study of matter Matter is the “stuff” that makes up the universe The fundamental questions of Chemistry are:

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Ch 100: Fundamentals for Chemistry

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  1. Ch 100: Fundamentals for Chemistry Chapter 1: Introduction Lecture Notes

  2. What is Chemistry? • Chemistry is considered to be the central science • Chemistry is the study of matter • Matter is the “stuff” that makes up the universe • The fundamental questions of Chemistry are: • How can matter be described? • How does one type of matter interact with other types of matter? • How does matter transform into other forms of matter?

  3. Scientific Method 1. Recognize a problem • Make observation • Ask a question 2. Make an educated guess - a hypothesis • Predict the consequences of the hypothesis 3. Perform experiments to test the predictions • Does experiment support or dispute hypothesis? 4. Formulate the simplest rule that organizes the 3 main ingredients - develop a theory

  4. The Scientific Attitude • All hypotheses must be testable (i.e. there must be a way to prove them wrong!!) • Scientific: “Matter is made up of tiny particles called atoms” • Non-Scientific: “There are tiny particles of matter in the universe that will never be detected”

  5. Major Developments in Chemistry I ~400 BC: Democritus proposed the concept of the “atom” ~300 BC: Aristotle developed 1st comprehensive model of matter ~700 AD: Chinese alchemists invent gunpowder 1661: Robert Boyle proposed the concept of elements 1770-1790: Lavoisier proposed the concept of compounds & the Law of Mass Conservation 1774: Priestly isolates oxygen 1797: Proust proposed the Law of Definite Proportions 1803: Dalton re-introduces the concept of the atom and establishes Dalton’s Laws 1869: Mendeleev creates the 1st Periodic Table 1910: Rutherford proposes the “nuclear” model of the atom 1915: Bohr proposes a “planetary” model of the hydrogen atom 1920: Schroedinger publishes his wave equation for hydrogen 1969: Murray Gell-Mann proposes the theory of QCD (proposing the existence of quarks)

  6. Major Developments in Chemistry II • Discovery of subatomic particles: • 1886: Proton (first observed by Eugene Goldstein) • 1897: Electron (JJ Thompson) • 1920: Proton (named by Ernest Rutherford) • 1932: Neutron (James Chadwick) • Other Important Discoveries: • 1896: Antoine HenriBecquerel discovers radioactivity • 1911: H. Kamerlingh Onnes discovers superconductivity in low temperature mercury • 1947: William Shockley and colleagues invent the first transistor • 1996: Cornell, Wieman, and Ketterle observe the 5th state of matter (the Bose-Einstein condensate) in the laboratory

  7. Ch 100: Fundamentals for Chemistry Chapter 2: Measurements & Calculations Lecture Notes

  8. Types of Observations • Qualitative • Descriptive/subjective in nature • Detail qualities such as color, taste, etc. • Example: “It is really warm outside today” • Quantitative • Described by a number and a unit (an accepted reference scale) • Also known as measurements • Example: “The temperature is 85oF outside today”

  9. Measurements • Described with a value (number) & a unit (reference scale) • Both the value and unit are of equal importance!! • The value indicates a measurement’s size (based on its unit) • The unit indicates a measurement’s relationship to other physical quantities

  10. Scientific Notation • Technique Used to Express Very Large or Very Small Numbers • Based on Powers of 10 • To Compare Numbers Written in Scientific Notation • First Compare Exponents of 10 (order of magnitude) • Then Compare Numbers

  11. Writing Numbers in Scientific Notation • Locate the Decimal Point • Move the decimal point to the right of the non-zero digit in the largest place • The new number is now between 1 and 10 • Multiply the new number by 10n • where n is the number of places you moved the decimal point • Determine the sign on the exponent, n • If the decimal point was moved left, n is + • If the decimal point was moved right, n is – • If the decimal point was not moved, n is 0

  12. Writing Numbers in Standard Form • Determine the sign of n of 10n • If n is + the decimal point will move to the right • If n is – the decimal point will move to the left • Determine the value of the exponent of 10 • Tells the number of places to move the decimal point • Move the decimal point and rewrite the number

  13. Measurement Systems There are 3 standard unit systems we will focus on: 1. United States Customary System (USCS) • formerly the British system of measurement • Used in US, Albania, and a couple others • Base units are defined but seem arbitrary (e.g. there are 12 inches in 1 foot) 2. Metric • Used by most countries • Developed in France during Napoleon’s reign • Units are related by powers of 10 (e.g. there are 1000 meters in 1 kilometer) 3. SI (L’Systeme Internationale) • a special set of metric units • Used by scientists and most science textbooks • Not always the most practical unit system for lab work

  14. Related Units in the Metric System • All units in the metric system are related to the fundamental unit by a power of 10 • The power of 10 is indicated by a prefix • The prefixes are always the same, regardless of the fundamental unit

  15. Units & Measurement • When a measurement has a specific unit (i.e. 25 cm) it can can be expressed using different units without changing its meaning • Example: • 25 cm is the same as 0.25 m or even 250 mm • The choice of unit is somewhat arbitrary, what is important is the observation it represents

  16. Measurement & Uncertainty • A measurement always has some amount of uncertainty • Uncertainty comes from limitations of the techniques used for comparison • To understand how reliable a measurement is, we need to understand the limitations of the measurement

  17. Measurements & Significant Figures • To indicate the uncertainty of a single measurement scientists use a system called significant figures • The last digit written in a measurement is the number that is considered to be uncertain • Unless stated otherwise, the uncertainty in the last digit is ±1

  18. Rules for Counting Significant Figures • Nonzero integers are always significant • Zeros • Leading zeros never count as significant figures • Captive zeros are always significant • Trailing zeros are significant if the number has a decimal point • Exact numbers have an unlimited number of significant figures

  19. Rules for Rounding Off • If the digit to be removed • is less than 5, the preceding digit stays the same • is equal to or greater than 5, the preceding digit is increased by 1 • In a series of calculations, carry the extra digits to the final result and then round off • Don’t forget to add place-holding zeros if necessary to keep value the same!!

  20. Exact Numbers • Exact Numbers are numbers known with certainty • Unlimited number of significant figures • They are either • counting numbers • number of sides on a square • or defined • 100 cm = 1 m, 12 in = 1 ft, 1 in = 2.54 cm • 1 kg = 1000 g, 1 LB = 16 oz • 1000 mL = 1 L; 1 gal = 4 qts. • 1 minute = 60 seconds

  21. Converting between Unit Systems To convert from one unit to another: • Identify the relationship between the units (e.g. 100 cm = 1 m) • Write out the starting measurement and multiply it by a quantity that will yield the desired value: 25 cm ( ) = _____ m • The number in the “( )” is called the “conversion factor”

  22. Metric Prefixes

  23. Mass is the amount of “stuff” in an object Mass is inertia Mass is the same everywhere in the universe SI Units of mass are kilograms (kg) Weight is the effect of gravity on an object’s mass Weight is a force Weight depends on location SI units of weight are newtons (N) USCS units are pounds (lb) Weight vs. Mass

  24. Volume • The 3-D space an object occupies • The SI unit is m3(meters x meters x meters) • The common metric unit is the Liter (L) • Mass and volume are not the same thing • Do not confuse mass & volume

  25. Density • Density is a property of matter representing the mass per unit volume • For equal volumes, denser object has larger mass • For equal masses, denser object has small volume • Solids = g/cm3 • 1 cm3 = 1 mL • Liquids = g/mL • Gases = g/L • Volume of a solid can be determined by water displacement • Density : solids > liquids >>> gases • In a heterogeneous mixture, denser object sinks

  26. Using Density in Calculations

  27. Ch 100: Fundamentals for Chemistry Chapter 3: Matter & Energy Lecture Notes

  28. Aristotle (384-322 BC) 1) Earth 1) Cold 3) Air 3) Hot 2) Water 2) Moist 4) Fire 4) Dry • Introduced observation as an important step in understanding the natural world • All types of matter are mixtures of one of 4 basic “elements”: • All matter has one or more of 4 basic “qualities”: • According to Aristotle: • Any substance could be transformed into another substance by altering the relative proportion of these qualities (i.e. lead to gold)

  29. Physical & Chemical Properties • Physical Properties are the characteristics of matter that can be changed without changing its composition • Characteristics that are directly observable • Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy • Characteristics that describe the behavior of matter

  30. Physical & Chemical Changes • Physical Changes are changes to matter that do not result in a change the fundamental components that make that substance • State Changes : boiling, melting, condensing • Chemical Changes involve a change in the fundamental components of the substance • Produce a new substance • Chemical reaction • Reactants  Products

  31. States of Matter +Energy +Energy Solid→ Liquid → Gas Solid← Liquid ← Gas +Energy +Energy

  32. Classification of Matter Matter can be classified as either Pure or Impure: • Pure • Element: composed of only one type of atom • Composed of either individual atoms or molecules (e.g. O2) • Compound: composed of more than one type of atom • Consists of molecules • Impure (or mixture) • Homogeneous: uniform throughout, appears to be one thing • pure substances • solutions (single phase homogeneous mixtures) • Suspensions (multi-phase homogeneous mixtures) • Heterogeneous: non-uniform, contains regions with different properties than other regions

  33. Separation of Mixtures • A pure substance cannot be broken down into its component substances by physical means only by a chemical process • The breakdown of a pure substance results in formation of new substances (i.e. chemical change) • For a pure substance there is nothing to separate (its only 1 substance to begin with) • Mixtures can be separated by physical means (and also by chemical methods, as well) • There are 2 general methods of separation • Physical separation • Chemical separation

  34. Methods of Separation • There are 2 ways of separating various substances: 1) Physical separation: separation of substances by their physical properties (such as size, solubility, etc.) • Mixtures can be separated by physical separation • There are several methods of separating mixtures • Filtration (solids from liquids) • Distillation (liquids from liquids) • Centrifugation (liquids from liquids) 2) Chemical separation: separation of substances by their chemical properties • Usages: • Compounds can be separated into their individual elements • Mixtures can be separated by chemical separation as well • There are several methods of chemical separation • Ion exchange (such as water purification systems) • Chemical affinity (using antibodies to isolate specific proteins) • Various Chemical reactions

  35. Energy • The capacity of something to do work • chemical, mechanical, thermal, electrical, radiant, sound, nuclear • The SI unit of energy is the Joule (J) • Other common units are • Calories (cal) • Kilowatt-hour (kW.hr) • Types of energy: • Potential • Kinetic • Heat • Energy cannot be created nor destroyed (but it does change from one type to another!)

  36. Heat & Temperature • Temperature is _____. • how hot or cold something is (a physical property) • related to the average (kinetic) energy of the substance (not the total energy) • Measured in units of • Degrees Fahrenheit (oF) • Degrees Celsius (oC) • Kelvin (K) • Heat is energy that _____. • flows from hot objects to cold objects • is absorbed/released by an object resulting in its change in temperature • Heat absorbed/released is measured by changes in temperature

  37. Temperature Scales • Fahrenheit Scale, °F • Water’s freezing point = 32°F, boiling point = 212°F • Celsius Scale, °C • Temperature unit larger than the Fahrenheit • Water’s freezing point = 0°C, boiling point = 100°C • Kelvin Scale, K • Temperature unit same size as Celsius • Water’s freezing point = 273 K, boiling point = 373 K

  38. Temperature of ice water and boiling water.

  39. Heat • Heat is the flow of energy due to a temperature difference • Heat flows from higher temperature to lower temperature • Heat is transferred due to “collisions” between atoms/molecules of different kinetic energy • When produced by friction, heat is mechanical energy that is irretrievably removed from a system • Processes involving Heat: • Exothermic = A process that releases heat energy. • Example: when a match is struck, it is an exothermic process because energy is produced as heat. • Endothermic = A process that absorbs energy. • Example: melting ice to form liquid water is an endothermic process.

  40. Heat (cont.) • The heat energy absorbed by an object is proportional to: • The mass of the object (m) • The change in temperature the object undergoes (DT) • Specific heat capacity (s) (a physical property unique to the substance) • To calculate heat (Q): Q = s . m .DT

  41. Specific Heat Capacity (s) • The amount of heat energy (in J or Cal) required to increase the temperature of 1 gram of a substance by 1oC (or 1K) • The Units of Specific Heat Capacity: • J/goC (SI) • cal/goC (metric & more useful in the lab) • Specific Heat Capacity is a unique physical property of different substances • Metals have low specific heat capacity • Non-metals have higher specific heat capacity • Water has an unusually large specific heat capacity s = Q/(mDT)

  42. Table of Specific Heat for various substances @ 20oC Substance c in J/gm K c in cal/gm K orBtu/lb F Molar CJ/mol K Aluminum 0.900 0.215 24.3 Bismuth 0.123 0.0294 25.7 Copper 0.386 0.0923 24.5 Brass 0.380 0.092 ... Gold 0.126 0.0301 25.6 Lead 0.128 0.0305 26.4 Silver 0.233 0.0558 24.9 Tungsten 0.134 0.0321 24.8 Zinc 0.387 0.0925 25.2 Mercury 0.140 0.033 28.3 Alcohol(ethyl) 2.4 0.58 111 Water 4.186 1.00 75.2 Ice (-10 C) 2.05 0.49 36.9 Granite .790 0.19 ... Glass .84 0.20 ...

  43. Ch 100: Fundamentals for Chemistry Chapter 4: Elements, Ions & Atoms Lecture Notes

  44. Dmitri Mendeleev (1834-1907) • Russian born chemist • Considered one of the greatest teachers of his time • Organized the known elements into the first “periodic table” • Elements organized by chemical properties (& by weight) -> called periodic properties • Predicted the existence of 3 new elements

  45. Chemical Symbols & Formulas • Each element has a unique chemical symbol • Examples of chemical symbols: • Hydrogen: H • Oxygen: O • Aluminum: Al • Each molecule has a chemical formula • The chemical formula indicates • the chemical symbol for each of the elements present • The # of atoms of each element present in the molecule • Examples of chemical formulas: • Elemental oxygen: O2(2 O atoms per molecule) • Water: H2O (2 H atoms & 1 O atom) • Aluminum sulfate: Al2(SO4)3(2 Al, 3 S & 12 O atoms)

  46. Dalton’s Atomic Theory • Each element consists of individual particles called atoms • Atoms can neither be created nor destroyed • All atoms of a given element are identical • Atoms combined chemically in definite whole-number ratios to form compounds • Atoms of different elements have different masses

  47. The Atom The atom has 2 primary regions of interest: 1) Nucleus • Contains protons & neutrons (called nucleons, collectively) • Establishes most of the atom’s mass • Mass of 1 neutron = 1.675 x10-27 kg • Mass of 1 proton = 1.673 x10-27 kg • Small, dense region at the center of the atom • The radius of the nucleus ~ 10-15 m (1 femtometer) 2) The Electron Cloud • Contains electrons • Mass of 1 electron = 9.109 x10-31 kg • Establishes the effective volume of the atom • The radius of the electron cloud ~ 10-10 m (1 Angstrom) • Determines the chemical properties of the atom • During chemical processes, interactions occur between the outermost electrons of each atom • The electron properties of the atom will define the type(s) of interaction that will take place

  48. Structure of the Atom

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