Fundamentals for Chemistry
Fundamentals for Chemistry CH 100
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Ch 100 Fundamentals for Chemistry Chapter 1 Introduction Lecture Notes What is Chemistry Chemistry is often described as the central solence Chemistry is the study of matter Matter is the stuff that makes up the universe ie anything that has mass and occupies space The fundamental questions of Chemistry are 1 How can matter be described 2 How does one type of matter interact with other types of matter 3 How does matter transform into other forms of matter Major Developments in Chemistry 400 BC Democritus proposed the concept of the atom 300 BC Aristotle developed 15 comprehensive model of matter 700 AD Chinese alchemists invent gunpowder 1661 Robert Boyle proposed the concept of elements 177090 Lavoisier proposed the concept of compounds amp the Law of Mass Conservation 1774 Priestly isolates oxygen 1797 Proust proposed the Law of Definite Proportions 1803 Dalton reintroduces the concept of the atom and establishes Dalton s Laws 1869 Mendeleev creates the 1 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 GeIIMann proposes the theory of QCD proposing the existence of quarks Fonlaua commum college 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 Henri Becquerel 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 BoseEinstein condensate in the laboratory Scientific Method 1 OBSERVATION Recognize a problem Make observation Formulate a question 2 EXPLANATION Make an educated guess a hypothesis Predict the consequences of the hypothesis 3 VALIDATION Perform experiments to test the predictions Does experimental data support or dispute hypothesis Formulate the simplest rule that organizes the 3 main ingredients develop a theory 4 Observations Laws analysis I explanation gt Hypothesis I 4 Experiment 1analysis Theory model Hypothesis I I Theory I Tentative Explanation Explanation of the of General Principles Certain Facts of Certain Phenomena Provides a Considerable Evidence Basis for or Facts Further Experimentation Support lt Simple Statement No Exceptions of Natural Phenomena Under the Given Conditions Bottom Line The Scientific Attitude All hypotheses must be testable ie there must be a way to prove them wrong Scientific Matter is made up of tiny particles called atoms NonScientific There are tiny particles of matter in the universe that will never be detected The Particulate Nature of Matter Matter is the tangible substance of nature anything with mass that occupies space At the most fundamental level matter is discrete or particulate in nature The smallest most basic units of matter are called atoms All matter is thus comprised of individual atoms or specific combinations of atoms called molecules Molecules can be broken apart into their constituent atoms but atoms cannot be further broken apart and still retain the properties of matter Matter can exist in one or more physical states or phases Fontmi community can States of Matter Solid E igy Liquid E igy Gas State Shape Volume Compress Flow Solid Keeps Keeps No No Shape Volume Liquid Takes Keeps No Yes Shape of Volume Container Gas Takes Takes Yes Yes Shape of Volume of Container Container Solid 53 LiquidEltne yGas womaua commun39 Cottage 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 eg 02 Compound composed of more than one type of atom Consists of molecules onstant Compositio Pure Substance Homogeneous F Mixture 391 Variable Compositio Impure or mixture Homogeneous uniform throughout appears to be one thing Pure substances Solutions single phase homogeneous mixtures Suspensions multiphase homogeneous mixtures Heterogeneous nonuniform contains regions with different properties than other regions lammi community can Separation of Matter 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 ie chemical change For a pure substance there is nothing to separate its only 1 substance to begin with Mixtures be separated by physical means and also by chemical methods as well There are 2 general methods of separation 1 Physical separation based on physical properties 1 Filtration 2 Distillation 3 Centrifugation 2 Chemical separation based on chemical properties lammi community can Ch 100 Fundamentals for Chemistry Ch 9 Calculations from Chemical Reactions Lecture Notes Sections 91 to 95 Fonlaua commun39 college Chemical Equations What do they tell us A properly written chemical equation will provide the following information 1 All reactants amp products involved in the reaction 2 The physical state of all reactants amp products 3 The presence of any catalysts involved in the chemical reaction 4 The relative quantity of all reactants amp products a Molecule to molecule ratios b Mole to mole ratios c Even mass to mass ratios can be determined with use of molar mass values lammi community can Information Given by the Chemical Equation A balanced chemical equation provides the relationship between the relative numbers of reacting molecules and product molecules Example The formation of carbon dioxide from carbon monoxide and oxygen gas 2 CO 02 gt 2 002 In this chemical equation it is indicated that 2 CO molecules react with every 1 O2 molecule to produce 2 002 molecules Alternative interpretation there is a 21 numerical ratio of CO to 02 for this completed reaction 2 001 02 2 002 5 r cart3W mind Interpretation of the Chemical Equation Since the information given in a balanced chemical reaction is relative 2 C0 02 gt 2 CO2 the following are alternative interpretations of the chemical equation a 200 CO molecules react with 100 O2 molecules to produce 200 C02 molecules b 2 billion CO molecules react with 1 billion 02 molecules to produce 20 billion CO2 molecules 0 2 moles CO molecules react with 1 mole O2 molecules to produce 2 moles CO2 molecules 01 12 moles CO molecules react with 6 moles O2 molecules to produce 12 moles CO2 molecules Note The coefficients in the balanced chemical equation also shows the molecules and mole ratio of the reactants and products Since moles can be converted to masses we can determine the mass ratio of the reactants and products as well Mole and Mass Ratios in Chemical Equations For the following chemical equation 2 C0 02 gt 2 CO2 The following mole relations are implied 2 moles CO 1 mole O2 2 moles CO2 Note the molar masses of the compounds in this reaction a 1 mole of CO 2801 g b 1 mole OZ 3200 g c 1 mole COZ 4401 g The mass ratio of the compounds in this reaction can be determined using the molar mass values 22a01 g CO13200g o2 24401 g 002 The mass ratio of the compounds in this reaction are 5602 9 CO 3200 g O2 8802 9 CO2 Example Determine the Number of Moles of Carbon Monoxide required to react with 32 moles Oxygen and the moles of Carbon Dioxide produced 1 Write the balanced equation 2 CO O2 gt 2 002 2 Use the coefficients to find the mole relationship 2 moles CO 1 molO2 2 moles 002 3 Use dimensional analysis to obtain the of moles a The mol of CO 32 moles 02 x mo 64 moles oo 1 mole 02 b The mol of 002 32 moles 02 x M 64 moles C02 1 mole O2 Ch 100 Fundamentals for Chemistry Ch 5 Early Atomic Theory amp Structure Section 1 1 Complete the following a The first observer of protons b First scientist to experimentally show the existence of electrons c This scientist proposed the nuclear atom hypothesis 1 This scientist discovered neutrons N Distinguish between the mass number and the atomic number for a given element LA Distinguish between the mass number and the atomic mass for a given element 4 Write out the simple representation for the formation of the following ions a Na b Mg2 c CI39 d SO42 5 What are isotopes 6 In what ways are isotopes of the same element similar amp different 7 Write out the representation for the following isotopes a Hydrogen mass number 2 b Uranium mass number 236 c Calcium with 25 neutrons Ch 100 Fundamentals for Chemistry Ch 5 Early Atomic Theory amp Structure Section 2 Using the periodic table as a guide fill out the following table Name Mass of the must Neutrons common Protons Electrons Ch 100 Fundamentals for Chemistry 3 Ch 5 Early Atomic Theory amp Structure The Per10d1c Table of Elements 1 18 1 2 1 H He 308 2 13 14 15 16 17 4003 3 4 5 6 7 8 9 10 2 Li Be B C N F Ne 6941 9012 1081 1201 1401 1600 1900 2018 11 12 13 14 15 16 17 18 3 Na Mg Al Si P S Cl Ar 2299 2431 3 4 5 6 7 8 9 2698 2809 3097 3206 3545 3995 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 3910 4008 4496 4788 5094 5200 5494 5585 5893 5870 6355 6538 6972 7259 7492 7896 7990 8380 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 8547 8762 8891 9122 9291 9594 98 1011 1029 1064 1079 1124 1148 1187 1218 1276 1269 1313 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 5 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg T Pb Bi Po At Rn 1329 1373 1389 1785 1809 1839 1862 1902 1922 1951 1970 2006 2044 2072 2090 209 210 222 87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 7 Fr Ra Ac Rf Db Sg Bh Hs Mt Unn Uuu ub Uu Uuh Uuo 223 2260 2270 261 262 263 262 265 266 269 272 277 289 289 293 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Lanthanides Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 1401 1409 1442 145 1504 1520 1573 1589 1625 1649 1673 1689 1730 1750 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Actinides Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 2320 231 2380 2370 244 243 247 247 251 252 257 258 259 262 Ch 100 Fundamentals for Chemistry Chapter 5 Worksheet Naming Compounds Naming elements part 1 Using the periodic table or table of elements write the chemical formula for each element and determine whether each element is a metal metalloid or nonmetal a hydrogen H2 nonmetal e chlorine Clz nonmetal b helium He nonmetal f silicon Si metalloid c calcium Ca metal g sodium Na metal d gold Au metal h sulfur S nonmetal Naming elements part 2 Using a Table of Elements write the chemical name for each element and determine whether each element is a metal metalloid or nonmetal a Pb lead metal e Hg mercury metal b K potassium metal f As arsenic metalloid c Ag silver metal g Ne neon nonmetal d Pt platinum metal h Fe iron metal Identifying Ions using the Periodic Table Identify the type of ion cation or anion and charge 2 1 etc that is formed by each of the following elements a H cation H or anion H39 d Ca cation Ca2 b CI anion CI39 e N cation A3 c Ne does not form an ion f O anion OZ39 Naming the Ions Name the ion formed by the following elements use the textbook if necessary a Na sodium ion d Mgmagnesium ion b I iodide ion e N aluminum ion c He does not form an ion f Hg mercuryI or mercury II Ch 100 Fundamentals for Chemistry Chapter 5 Worksheet Naming Compounds Naming Simple Compounds part 1 Write the chemical name systemic for each substance a KCI potassium chloride d N203 dinitrogen trioxide b MgClz magnesium chloride e AgZS silver sulfide c CCI4 carbon tetrachloride f PbOz eadIV oxide Naming Simple Compounds part 2 Write the chemical formula for each substance a oxygen Oz d ironIII chloride FeCI3 b lithium sulfide LiZS e copperI oxide CuZO c nitrogen dioxide N02 f tricarbon octahydride C3H8 Naming Polyatomic Ions Part 1 Name the following ions a CN39 cyanide d HCOg39 hydrogen carbonate b NH4 ammonium e 504239 sulfate c NOZ39 nitrite f OH39 hydroxide Naming Polyatomic Ions Part 2 Write the chemical formula for the following ions a carbonate CO3239 d sulfide 5239 b peroxide 02239 e nitrate N0339 c acetate C2H3OZ39 f phosphate PO4339 Ch 100 Fundamentals for Chemistry Chapter 5 Worksheet Naming Compounds Naming Simple Compounds part 3 Write the chemical name systemic for each substance a NH4C ammonium chloride d Fe2SO43 ironIII sulfate b NaNO3 sodium nitrate e NH4ZSO4 ammonium sulfate c Ca3PO42 calcium phosphate f BaOH2 barium hydroxide Naming Simple Compounds part 4 Write the chemical name systemic for each substance a potassium nitrate KN03 d calcium bicarbonate CaHC032 b ammonium hydroxide NH4OH e copperII acetate CuC2H3022 c sodium hypochlorite NaCIO f potassium cyanide KCN Naming Acids part 1 Write the chemical name systemic for each acid a HCI hydrochloric acid d H2503 sulfurous acid b HN03 nitric acid e HCIO hypochlorous acid c H3PO4 phosphoric acid f HZS hydrosulfuric acid Naming Acids part 2 Write the chemical name systemic for each acid a hydrofluoric acid HF d hydrocyanic acidHCN b nitrous acid HNOZ e acetic acid HC2H302 c phosphorous acid H3PO4 f carbonic acid HZCO3 Ch 100 Fundamentals for Chemistry Chapter 5 Worksheet Naming Compounds lammi community can Ch 100 Fundamentals for Chemistry Chapter 2 Measurements amp Calculations Lecture Notes Fonliua Deluge Types of Observations Qualitative Descriptivesubjective 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 Notes on Measurements Described with a value number amp 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 Example The temperature is 85 F outside today Application of Scientific Notation Writing numbers in Scientific Notation 1 Locate the Decimal Point 2 Move the decimal point to the Liqm of the nonzero digit in the largest place The new number is now between 1 and 10 3 Multiplythe new number by 10quot where n is the number of places you moved the decimal point 4 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 lfthe decimal point was not moved n is 0 Writing Scientific Notation numbers in Conventional form 1 Determine the sign of nof 10 If n is the decimal point will move to the right If n is the decimal point will move to the left 2 Determine the value of the exponent of 10 Tells the number of places to move the decimal point 3 Move the decimal point and rewrite the number womaua commum Cottage 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 other countries Base units are defined but seem arbitrary eg 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 eg there are 1000 meters in 1 kilometer 3 SI L Systeme Internationale a subset set of metric units Used by scientists and most science textbooks Not always the most practical unit system for lab work Measurements amp 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 When a measurement has a specific metric unit Le 25 cm it can be expressed using different metric units without changing its meaning Example 25 cm is the same as 025 m or even 250 mm The choice of measurement unit is somewhat arbitrary what is important is the observation it represents Fonlaua commun39 college Measurement Uncertainty amp Significant Figures 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 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 othenNise the uncertainty in the last digit is 1 Examples 1 The measurement 252 cm uncertainty 01 cm 2 The measurement 2520 cm uncertainty 001 cm 3 The measurement 25200 cm uncertainty 0001 cm 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 Rules for Rounding Off If the digit to be removed is 1 less than 5 the preceding digit stays the same 2 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 placeholding zeros if necessary to keep value the same Fonlaua commum Cottage Exact Numbers Exact Numbers are numbers that are assumed to have unlimited number of significant figures are considered to be known with absolute certainty You do not need to consider or count signi cant gures for exact numbers The following are considered exact numbers for CH100 1 Counting numbers such as The number of sides on a square The number of apples on a desktop 2 Defined numbers such as those used for conversion factors such as 100cm1m12in1ft1in254cm 1kg1000g1LB160z 1000mL1L1gal4qts 1 minute 60 seconds 3 Numbers or constants defined in equations such as y 3x 15 both the 3 and the 15 are exact numbers lammi community can Converting between Unit Systems Converting units from one unit system to another especially within the Metric system can appear daunting at first glance However with a little guidance and a lot of practice you can develop the necessary skill set to master this process To begin here is a simple mnemonic to guide you through the unit conversion process 1 Eliminate 2 Replace 3 Relate All unit conversions regardless of how complex they appear involve these 3 simple steps In the following sections you will be stepped through the unit conversion process using these 3 words as a guide Fonlaua commun39 college Example Unit Conversion 1 Convert 250 m to cm 2 Convert 126 g to kg Pomona Community Calloge Metric Prefixes Table 22 The Commonly Used Prefixes in the Metric System Power of to for Prefix Symbol Meaning Scientific Notation mega M 1000000 10quot kilo k 1000 10quot deci d 01 10quot1 centi C 001 10 2 1111111 m 0001 10 3 micro p 0000001 10 name it 0000000001 10 9 mm Szwm Temperature Scales The 2 traditional temperature scales Fahrenheit and Celsius were originally defined in terms of the physical states of water at sea level L N Celsius Scale C Fahrenheit Scale F For water freezing point 32 F boiling point 212 F For water freezing point 0 C boiling point 100 C 1 Celsius temperature unit is largerthan 1 Fahrenheit unit The SI unit for temperature is a variant of the Celsius scale 3 Kelvin Scale K For water freezing point 273 K boiling point 373 K The Kelvin temperature unit is the same size as the Celsius unit Fonland 0 Community College Temperature of ice water and boiling water a b Unit Conversion amp Temperature Scales Unit conversion involving temperature is tricky since the zero value for each scale is different and thus requires accounting for this offset between the various scales At 000 the Kelvin scale has a 27315 unit head start and the Fahrenheit scale has a 32 unit head start 1 The temperature span between the freezing and boiling points of water reveal the relation between the temperature scale increments 1000C 100K 1800F 2 However the zero points are different as evident for the freezing point for water OOC 27315K 320F 3 The relations between the temperature scales 180 F Cl ch htTT 32 F a esrus o a ren er OF OC1000C b Celsiusto Kelvin TK To 100K 27315K C 100 C Mass Mass is the quantity of matter in a substance Mass is measured in units of grams Mass does not reflect how much volume something has The kilogram kg unit is the preferred unit of mass in the SI system a 1 kilogram is equal to the mass of a platinumiridium cylinder kept in a vault at Sevres France b 1 kg has the weight equivalent on Earth of 2205 lb wN k Conservation of Mass The total quantity of mass is never created nor destroyed during a chemical process womaua commum college Distinguishing Mass vs Weight The terms mass and weight are commonly used interchangeably but they are fundamentally different The following are some important differences between mass and weight Mass is a fundamental 1 Weight is the effect or property of matter it is the fOFPE Of graVitY 0 an amount of stuff in an object Obi 3 mass I Mass represents an object s 239 We39ght deper ds on locat39on inertia tendency to resist 8 39F Ca399raquot39t3 change in motion 339 We39ght 393 n Ot a I fundamental property of Mass IS the same everywhere matter m the un39Verse 4 SI units of weight are SI Units of mass are kilograms newtons N k9 USCS units are pounds lb N 9 P 01 Volume Volume is the 3dimensional space that an object occupies Volume Units The SI unit for volume is the cubic meter or m3 meters xmeters xmeters The more common metric unit of volume is the Liter L 1 In the laboratory the milliliter mL is often more convenient 1 mL 10393 L height I height I width lt length Note mass and volume are not the same thing try not to confuse them Two objects with the same volume eg a pillow amp a sack of potatoes can have different masses and vice versa Density Density is a property of matter representing the mass per unit volume For equal volumes a denser object has greater mass For equal masses a denser object has smaller volume Commonly used units 1 Solids gcm3 Note 1 cm3 1 mL 2 Liquids gmL Densityzm 3 Gases gL Volume Useful Notes on Density Volume of a solid can be determined by water displacement Density of matter in various states solids gt liquids gtgtgt gases exception water In a heterogeneous mixture the denser matter will tend to sink to the bottom Ponlnna 39 00m mun RV eonoqe Table 28 Densities of Various Common Substances at 20 C Substance Physical State Density gcm3 oxygen gas 000133 hydrogen gas 0000084 ethanol liquid 0785 benzene liquid 0880 water liquid 1000 magnesium solid 174 salt sodium Chloride solid 216 aluminum solid 270 iron solid 787 copper solid 896 silver solid 105 lead solid 1134 mercury liquid 136 gold solid 1932 At 1 atmosphere pressure Fu l nd com mun nv college Manipulating the Density Equation Densityz Volume Volume 2 LS Densnty Mass Density gtlt Volume Ch 100 Fundamentals for Chemistry Chapter 1 Scientific Method Worksheet Scientific Method a Describe the primary steps of the scientific method 1 Observationa State the problem 2 Expanationa Formulate a hypothesis 3 Experimentationa Test the hypothesis b Who is credited with establishing the scientific method It depends on who you ask 1 Sir Francis Bacon English philosopher a 2 Galileo Galilei Italian quotnatural philosopher 3 Robert Boye Irish chemist c What is required for a hypothesis to be scientific It must be testable i e there must be a testable way to prove it wrong d Give an example of both a scientific and nonscientific hypothesisstatement e Describe how qualitative and quantitative observations are different Give an example of each Quantitative observations are based on a numerical value and a unit reference scale Qualitative observations are not based on a numerical value and objective reference scale They are based on subjective qualities Being Scientific For each observationproblem formulate a scientific hypothesis to explain it Describe how you might test your hypothesis a Traffic is terrible and you are late for class What is the fastest way to get to school b You just finished a game of tennis and your elbow is sore What caused your elbow to be sore c The moon is bright yellow Why is the moon so yellow lammi community can Ch 100 Fundamentals for Chemistry Chapter 4 Properties of Matter Lecture Notes womaua commun39 college Physical amp Chemical Properties Physical Properties are the characteristics of matter that can be changed without changing its composition These characteristics are directly observable or measurable Types of Physical Properties Extrinsic Physical Properties are unique to objects ie size shape mass etc Intrinsic Physical Properties are unique to substances ie density conductivity color etc L P Chemical Properties are the characteristics of a substance that determine the tendency of the matter to transform in composition as a result of the interaction with other substances the influence of energy or both These are characteristics that describe the behavior of matter Physncal amp Chemical Changes Physical Changes are changes that do not result in a change the fundamental composition of the substance Typical Examples 1 Physical State Changes boiling melting condensing etc 2 Shape Size or Texture Changes Chemical Changes involve a change in the fundamental composition of the matter Notes on Chemical Change 1 Production of a new substances 2 Referred to as chemical reactions 3 The basic representation Reactants gt Products Note Both physical and chemical changes will likely produce an alteration of appearance the key is to discern the type of change that has occurred Fonlaua communmr College Energy Energy is loosely described as the capacity of something to do work or alter the physical or chemical state of an object or system Common Forms of Energy mechanical chemical thermal electrical radiant nuclear The Si unit of energy is the Joule J Other commonly used units are Calories cal and Kilowatthours kWhr Types of energy 1 Potential stored energy 2 Kinetic energy associated with motion and vibration 3 Heat energy that flows from high to low temperature Principle of Energy Conservation energy is never created nor destroyed but it does change from one type to another Distinguishing Heat Energy amp Temperature Temperature is How hot or cold something is an extrinsic physical property it represents a particular thermal state 2 Related to the avera e kinetic ener of the substance not thetotal ener but the average energy gy gy 3 Measured in units of Degrees Fahrenheit 0F Degrees Celsius OC Kelvin K Heat is 1 Energy that flows from hot objects to cold objects Heat isnot a physical propert 2 Energy absorbed or released by an object resulting in its temperature 0 an e 3 Measured in units of Joules J Calories Cal Kilowatt Hours kW hr Bottom Line Heat energy absorbed or released is measured by changes in temperature but do not confuse heat energy for temperature Fonlaua commun39 college Temperature Scales The 2 traditional temperature scales Fahrenheit and Celsius were originally defined in terms of the physical states of water at sea level Fahrenheit Scale F For water freezing point 32 F boiling point 212 F Celsius Scale C For water freezing point 0 C boiling point 100 C 1 Celsius temperature unit is largerthan 1 Fahrenheit unit l The SI unit for temperature is a variant of the Celsius scale 3 Kelvin Scale K For water freezing point 273 K boiling point 373 K The Kelvin temperature unit is the same size as the Celsius unit Fonland amp Community col Iaqe Temperature of ice water and boiling water a b Heat Energy Heat is energy that flows due to a temperature difference Heat energy flows from higher temperature to lower temperature Heat is transferred due to collisions between atomsmolecules of different kinetic energy When produced by friction heat is mechanical energy that is irretrievably removed from a system Processes involving Heat 1 Exothermic A process that releases heat energy O Example burning paper is an exothermic process because energy is produced as heat the temperature risesl 2 Endothermic A process that absorbs energy O Example melting ice to form liquid water is an endothermic process because heat energy must be absorbed to change the physical state in this case the temperature does not changel Heat cont When something absorbs or loses heat energy 1 of 2 things can occur 1 Its temperature will change eg hot coffee will cool down 2 Its physical state will change eg ice will melt For the former case above the heat energy absorbed or lost by an object is proportional to 1 The mass of the object m 2 The change in temperature the object undergoes AT 3 The specific heat capacity S a physical property unique to the substance To calculate heat gained Q Q smAT Portland commum college Specific Heat Capacity 3 Specific heat capacity reflects how absorbed heat energy relates to the corresponding increase in temperature for a given amount of mass ie energy per unit mass per unit temperature change or Q S m AT Specific Heat Capacity is commonly measured in units of 1 Jg C SI 2 calg C metric amp more useful in the lab Specific Heat Capacity is a unique intrinsic physical property of matter Typically 1 Metals have low specific heat capacity 2 Nonmetals have higher specific heat capacity than metals 3 Water has an unusually large specific heat capacity a cm que Fn l Ella I Commun y Table 32 The Specific Heat Capacities of Some Common Substances Substance Specific Heat Capacity Jg C water 0 liquid water 3 ice water g steam aluminum 5 iron 5 mercury 1 carbon 5 silver 5 gold 5 4184 203 20 0473 014 071 024 013 The symbols 5 l and g indicate the solid liquid and gaseous states respectively P l lld amp co quotin Co Inmu liege Table of Specific Heat for Various Substances Substance JgK calgK JmolK Aluminum 0900 0215 243 Iron 0473 0113 264 Copper 0385 00921 245 Brass 0380 0092 Gold 0131 00312 256 Lead 0128 00305 264 Silver 0233 00558 249 Tungsten 0134 00321 248 Zinc 0387 00925 252 Mercury 0140 0033 283 Alcohol ethyl 2138 0511 111 i Water i i 4184 i i 1000 i i 752 Ice 10 C 2059 0492 369 Granite 790 019 Glass 84 020 lammi community can Ch 100 Fundamentals for Chemistry Ch 7 Quantitative Composition of Compounds Lecture Notes Sections 71 to 73 womaua commun39 college The Mole The mole is a counting unit analogous to the dozen unit A large unit used to describe large quantities such as number of atoms 1 mole 6022 X 1023 units 6022 X 1023 is known as Avogadro s number NA Relationship between the mole amp the Periodic Table The atomic mass is the quantity in grams of 1 mole of that element The units of atomic mass are gramsmole Mass is used by chemists as a way of counting number of atomsmolecules of a substance Mole calculations lammi community can Got mole problems Call Avogadro at 6021023 What do you get if you have Avogadro39s number of donkeys Answer molasses a mole of asses m m and mum m eo college Molar Mass Molar mass is the mass in grams of 1 mole of a substance Molar mass refers to both atoms amp molecules Elements atoms Examples 1 mole of Na has a mass of 2299 g 1 mole of Cl has a mass of 3545 1 mole of Cl2 has a mass of 7090 g Compounds molecules Examples 1 mole of NaCl has a mass of 5844 9 Mass of Na 2299 g Mass of CI 3545 g 1 mole of CO2 has a mass of 4401 9 Mass of C 1201 g 2 x Mass of O 1600 g lammi community can Mole Calculations 1 To convert from atoms or molecules to moles divide the of atoms or molecules by Avogadro s Example How many moles are 10x1024 atoms 1 mol 10gtlt 1024at0ms 6022x1023 17 mol 2 To convert from moles to atoms or molecules multiply the of atoms or molecules by Avogadro s Example How many molecules are in 25 moles 6022gtlt1023 25 mol 1 mol J15gtlt 1024 molecules womaua communw can MoleMass Calculations 1 To convert from moles to grams multiply the of moles by atomic mass Example How many grams in 25 moles of carbon 1201 g 1 25 mol 30 or 3x10 1 mol g 2 To convert from grams to moles divide the mass in grams by atomic mass Example How many moles are in 25 g of lithium 1 mol 25 036 I 36gtlt101 g6941 9 mo or Percent Composition Percent composition is the percentage of each element in a compound by mass Percent composition can be determined from either 1 the formula of the compound 2 the experimental mass analysis of the compound Composition jxlomo whole Note The percentages may not always total to 100 due to rounding Percent Composition Calculations To determine Composition from the chemical formula 1 Determine the molar mass of compound 2 Multiply the molar mass of the element of interest by the number of atoms per molecule then 3 Divide this value by the molar mass of the compound Composition of A atoms of Aatomc mass of A X100 molar mass of compound Example The Composition of sodium in table salt 1 The molar mass of NaCI is 5844 gmol 2 There is 1 atom of Na in each NaCI molecule 3 The atomic mass of Na is 2299 X 100 3933 00 Composition of Na j Pomn4 Com mullRy Collag Percent Composition Calculations Perform the following Composition calculations 1 The composition of carbon in carbon monoxide 2 The composition of oxygen in water 3 The composition of chlorine in sodium hypochlorite Pamm Com munity Collage Amadeo Avogadro 1 743 1 794 Italian lawyer turned chemist Major contributions included 1 Established difference between atoms amp molecules Oxygen amp nitrogen exist as molecules 02 amp N2 2 Reconciled the work of Dalton amp GuyLussac 3 Establishing Avogadro s Principle equal volumes of all gases at the same temperature and pressure contain the same number of molecules Note Avogadro did not determine Avogadro s number nor the mole these concepts came later 1 Avogadro is honored because the molar volume of all gases should be the same 2 Much of Avogadro s work was acknowledged after he died by Stanislao Cannizarro lammi community can Ch 100 Fundamentals for Chemistry Chapter 6 Nomenclature of Inorganic Compounds Fonlaua commun39 college Types of Compounds When compounds are formed they are held together by the association of electrons This association is called a chemical bond There are 3 general types of chemical bonds 1 Ionic 2 Covalent or molecular 3 Polar covalent Simple compounds are classified and thus named according to the type of chemical bonds that hold together its atoms Note many compounds have more than one type of chemical bond present but we will focus on only simple compounds Types of Compounds cont For practical purposes will separate all simple compounds into 2 general categories Ionic Compounds a Made up of ions both positive and negative charge b Must have no net charge ie combined charge of zero c Depend on the attraction between positive and negative charges of the ions I Usually a metal is present as a cation and a nonmetal is present as an anion 2 NonIonic aka Molecular or Covalent Compounds a Made up of atoms that share their outer electrons b Electric charge plays no direct role in their formation c There are usually no metals are present in these compounds womaua commum college Naming Compounds The easiest way usually to identify an ionic compound is to ask whether or not there is a metal present in the chemical formula or the name Is a metal present Yes gt it is an Ionic Compound eg CaCI2 No gt it is a NonIonic Compound eg CCI4 or an Acid Notes 1 Ionic compounds do not use the Greek prefixes and are named according to the identity of the ions present 2 NonIonic compounds require the use of Greek prefixes to indicate the number of each element present in one molecule Naming Simple Compounds A simple or binary compound is a compound made of only 2 types of elements When the first element is a metal The first element metal keeps its full name The nonmetal goes by its root with the suffix idequot added to the end Example NaCI is sodium chloride When there are no metals present Same as above except Greek prefixes must be used to identify the number of each element present in the compound Example 002 is carbon goxide Fonlaua commum college Determining Chemical Formula of an Ionic Compound To determine the chemical formula of an ionic compound from its chemical 1 Identify the ions present both cations and anions from the name Example potassium sulfide Cation potassium Anion sulfide 2 Determine the ionic charge of the ions Example from above potassium ion K sulfide ion 8239 3 Determine the number of each ion needed to obtain a neutral compound Example from above a 2 K ions are needed for every 8239 3 Combine the chemical sysmbols of the ions to get the final chemical formula Example from above a K28 is the formula for potassium sulfide lammi community can Ionic Charges amp the Periodic Table The position of an element in the Periodic Table is a useful indicator of the type of ion an element is capable of forming 1 5 90 NSDPquot Group 1 metals form 1 cations Na sodium ion Group 2 metals form 2 cations Ca2 calcium ion Group 13 metals form 3 cations Al3 aluminum ion Group 312 Metas plus Sn Pb amp Bi can form more than one type of cation Roman numerals are used to indicate the charge of the cation Example Fest is called ironlll FeCl3 is called ironlll chloride Notable Exceptions an Group 15 nonmetals form 3 anions eg N339 nitride ion Group 16 nonmetals form 2 anions eg 0239 oxide ion Group 17 nonmetals form 1 anions eg Cl39 chloride ion Group 18 elements do not form ions M 4 EVBVV 1 8393 Fonliua Deluge Greek Prefixes for Compound Names Mono 6 Hexa Di 7 Hepta Tri 8 Octa Tetra 9 Nona Penta 10 Deca CCI4 is carbon tetrachloride Notes C3H8 is gearbon octahydride Prefixes are used when the compound does not have a metal present or when H is the first element in the formula Prefixes must be used for every element present in the compound Mono is not used for the first element in a compound name eg carbon dioxide Ionic Compounds containing Polyatomic ions Some ionic compounds are made up of polyatomic ions Polyatomio ions are usually ions formed from nonionic molecules eg The sulfate ion 804239 is essentially a molecular compound containing 8 and O with 2 additional electrons When you encounter polyatomic ions in compounds do not freak out Become familiar with the common polyatomic ions on the handout Example The nitrate ion NO339 Fortunately the naming of ionic compounds containing polyatomic ions is similar to that for ionic compounds Fonlaua communw can Acids From the Latin term for sour Acids are sourto the taste Acids are substances that donate or release hydrogen cations H usually when dissolved in water The chemical formula for acids usually begins with H Example hydrochloric acid HCI HClaq gt H Cl39 aq Bases Taste bitter Note it is not advised to taste strong bases Usually metal containing hydroxides Substances that accept hydrogen cations H when dissolved in water Example potassium hydroxide KOH KOHaq H gt Kaq H20 I Naming Acids Lets separate acids into 2 types 1 Acids that contain oxygen 2 Acids that do not contain oxygen Naming acids containing oxygen 1 For acids containing atequot anions a Use root ofthe anion for sulfate 804239 use sulfur b Add ic suffix then end with acid Example H2804 is sulfuric acid 2 For acids with itequot anions a Use root ofthe anion for sulfite 803239 use sulfur b Add ous suffix then end with acid Example H2803 is sulfurous acid Fonlaua commum college Naming Acids cont Naming acids not containing oxygen 1 Add hydroquot prefix to beginning 2 Use root of the anion ie Cl39 use chlor 3 Add icquot suffix then end with acid Examgle HCI is hydrochloric acid Name the following acids HF HNO2 HCN HSPO4 Antoine Lavoisier 17431794 Considered by many to be the Father of Modern Chemist Major contributions included Demonstrated that water cannot be transmuted to earth Established the Law of Conservation of Mass lquot 0 Developed a method of producing better gunpowder Observed that oxygen and hydrogen combined to produce water dew Invented a system of chemical nomenclature still used in pan today Wrote the 13 modern chemical textboo 5 01 o lammi community can Ch 100 Fundamentals for Chemistry Chapter 8 Chemical Equations Lecture Notes Fonlaua communmr college Chemical Equations Intro Chemical equations are used to symbolically describe chemical reactions 2 In a chemical equation or reaction for that matter the substances that undergo chemical changes are called the reactants The resulting substances formed are called the products The standard representation of a chemical equation Reactants gt Products Example The production of water 2H2 9 102 9 gt gHzo 9 00 The underlined numbers are called coefficients a The number of each molecule for each reactant amp product in the chemical reaction b They are always whole numbers Chemical Equations cont Balanced chemical equations indicate the identity of each reactant amp product involved in the reaction 2 phase of each reactant and product involved in the reaction ie solid 5 liquid I or gas 9 relative quantity of each reactant and product involved in the reaction the coefficients relative molar quantity of each reactant and product involved in the reaction the coefficients 9 P womaua commun39 college Balancing Chemical Equations According to the Law of Mass Conservation amp John Dalton matter is never created nor destroyed during chemical reactions All of the atoms in the reactants of a chemical reaction must be accounted for in the products The Basic Process of Balancing Chemical Equations 1 Identify all reactants amp products in the reaction amp write out their formulas this is the unbalanced chemical equation 2 Count the number of each atom for each compound for each reactant amp product these values must be the same for both reactants amp products when the reaction is balanced 4 Starting with the most complicated molecule systematically adjust the coefficients to balance of the atoms on each side of the reaction balance one atom at a time 5 Repeat until all atoms are balanced for the reaction 6 Now you should have a balanced chemical equation Balancing Chemical Equations example When sodium metal is added to water a violent reaction takes place producing aqueous sodium hydroxide and releasing hydrogen gas 1 Write out the unbalanced chemical reaction 2 Now balance the chemical reaction Fonlaua commun39 college Balancing Chemical Reactions Hint When a polyatomic ions appears on both the reactant amp product side of the reaction unchanged treat the whole ion as a unit when balancing the reaction Example AgNO3aq CaC12aq gt AgC1s CaNO32aq Note the nitrate ion NO339 gets swapped between the Ag and the Ca2 ions in this reaction 2 80 N0339 can be treated as a whole unit when balancing this reaction Balance it 0 Common Classifications for Chemical Reactions 1 Corgbgtation or Synthesis reactions in which reactants combine to make one pro u 2 Decomposition reactions in which one reactantbreaks down into smaller products 3 Sin le Displacement reactions where a part of one reactant is displaced and com ined with another reactant 22ns 2HCaq a ZnCl2aq H2g 4 Double Displacement reactions where a part of two reactants is displaced and exchanged AgN03aq NaClaq a AgCls NaN03aq Examples a Acidbase neutralization b Formation of insoluble products Precipitation reactions 0 Metal oxide acid d Gas formation 5 OxidationReduction Reactions reactions involving thetransfer or rearrangement of electrons womaua commum college Combination amp Decomposition Reactions Reactions in which chemicals combine to make one productare called Combination or Synthesis Reactions a Metal I Nonmetal reactions can be classified as Combination Reactions 2 Nas Cl2g gt 2 NaCls b Reactions between Metals or Nonmetals with 02 can be classified as Combination Reactions N2g 029 gt 2 NOg Note these two types of Combination Reactions are also subclasses of OxidationReduction Reactions 2 Reactions in which one reactantbreaks down into smaller molecules are called Decomposition Reactions a Decomposition reactions are generally initiated by the addition of energy via electric current or eat b Decomposition reactions are the opposite of Combination Reactions 2 NaClI gt 2 Nal CI2g Single Displacement Reactions Single displacement reactions involve one part of a reactant being transferred to another The basic pattern of the single displacement reaction XY A gt X AY Example 1 Metal Acid gt Salt Hydrogen Zns 2 HCaq gt ZnCI2aq H2g Example 2 Metal Water gt Hydrogen Metal Oxide or metal hydroxide 3 Fes 4 H20l gt 4 H2g Fe203s Example 3 Metal Salt gt Metal Salt 2 Als Fe203s gt 2 Fes Al203s Example 4 Halogen Halide Salt gt Halogen Halide Salt CI2g 2 NaBrs gt Br2g 2 NaCs Fonlaua Commu 39w college Double Displacement Reactions Double Displacement Reactions involve the double exchange of a component such as ions between two reactants The basic form of double displacement reactions is XY AB gt XB AY where X Y A and B are the components of the reactants Example 1 Acid Base Neutralization HZSO4aq CaOH2aq gt CaSO4aq 2 HZOU or 2HOH Example 2 Metal Oxide Acid CaOs 2HCaq gt CaCI2aq H20 or HOH Example 3 Formation of an Insoluble Precipitate Precipitation KCaq AgNO3aq gt KNO3aq AgCIs Example 4 Formation of a Gas HCI aq ZnS s gt ZnCI2aq H28 9 Solubility amp PreCIpItatIon Reactions When 2 solutions are combined and result in the formation of an insoluble product a The product will not dissolve in the solvent b The product will form am Solubility is an intrinsic physical property and a measure of how well a substance solute will dissolve in another substance solvent a Solubility is temperature dependent b Solid solubility increases with increased temperature ie you can dissolve more sugar in hot water than in cold water 0 Gas solubility increases with decreased temperature ie you can dissolve more CO2 in cold water than hot water 3 A solute is soluble if any of it will dissolve in a solvent Eg NaCl is soluble in water 4 A solute is insoluble if no appreciable amount of it will dissolve in solvent Eg AgCl is insoluble in water Precipitation formation of an insoluble solid is one indication that a chemical change has occurred N 01 Fonliua Deluge General Rules for Solubility 1 Most compounds that contain NO339 ions are soluble 2 Most compounds that contain Na K or NH4 ions are soluble 3 Most compounds that contain Cl39 ions are soluble except AgCl PbClz and H920I2 4 Most compounds that contain 804239 ions are soluble except BaSO4 PbSO4 and CaSO4 5 Most compounds that contain OH39 ions are slightly soluble will precipitate except NaOH KOH are soluble and BaOH2 CaOH2 are moderately soluble 6 Most compounds that contain 8239 003239 or PO4339 ions are slightly soluble will precipitate OxidationReduction Reactions Reactions that involve transfer or rearrangement of electrons are called oxidationreduction reactions Examples of oxidationreduction reactions 1 Metal Nonmetal 2Nas Cl2g gt 2NaCls a The metal loses an electrons and becomes a cation oxidation gt metal gets oxidized Na gt Na e39 b The nonmetal gains an electrons and becomes an anion reduction gt nonmetal gets reduced Cl e39 gt Cl39 c In this reaction electrons are transferred from the metal to the nonmetal 2 02 as a reactant or product CH4s 02g gtC02g HZOg a In this reaction it is not obvious that electron transfer has taken place In this case oxidation states are altered b Often this type of reaction involves the release of large amounts of energy even combustion Rates of Chemical Reactions How quickly a chemical reaction occurs is indicated by its reaction rate 1 How quickly the concentration of products increases 2 How quickly the concentration of reactants decreases The Factors that influence reaction rates 1 Reactants must be in contact Reactions occur due to collisions Without contact between reactants there can be no reaction 2 Concentration of reactants The more reactant molecules packed into a given space the more likely a collision amp reaction will occur 3 Temperature the average KE of each reactant affects how much energy will be transferred between reactants during a molecular collision Molecules must transfer enough KE to breakthe existing bonds Femmi community can 1 Energy is requi energy 2 1 reactants 2 release N29 The Role of Energy in Chemical Reactions Energy transformations always accompany chemical reactions red to break bonds energy absorbed or activation Energy is released when bonds are formed Note The amount of energy required to break a chemical bond equals the energy released when that type of bond is formed this is called the Bond Energy For a chemical reaction to occur Energy must be absorbed in order to break chemical bonds in the Energy is released as new bonds are formed in the products Endothermic reactions absorb more energy than they 029 393 kJ gt 2NOg m reactions release more energy than they absorb H2g Cl2g gt 2NOg 185 kJ Fonliua Deluge Energy in Chemical Reactions Exothermic Reactions PotentialA Activation Energy Energy EA Reactants Energy Released Q Products Endothermic Reactions Potential Energy Activation Energy EA Products Alysetbed Q Reactants Energy in Reactions cont Example Sodium Water Reaction Internal Energy Low Activation Energy EA 2Nas 2H200 Large amount of Energy Released 0 2Na0Haq H2g Fonlaua communw College Catalysts 1 Catalysts are substances that speed up chemical reactions a Allow reactions to occur that might not othenNise take place due to low temperature for example b Lower activation energy for a chemical reaction Participation of catalysts in a chemical reaction a They may undergo a chemical change as a reactant but they are always recycled as a product so there is no net change in the catalyst molecule Catalysts are indicated in a chemical reaction by placing the chemical formula overunder the reaction arrow N 0 catalyst Reactants gt Products Example The breakdown of hydrogen peroxide catame ZHzoz aCI gt H20 0 029 lammi community can Catalysts amp Energy in Reactions Catalysts lower Activation Energy Activation Energy Potentia without catalyst Energy Activation Energy Reactams wtth catalyst Products Ch 100 Fundamentals for Chemistry Chapter 6 Nomenclature of Inorganic Compounds Naming elements part 1 Using the periodic table or table of elements write the chemical formula for each element and determine whether each element is a metal metalloid or nonmetal a hydrogen e chlorine b helium f silicon c calcium 9 sodium d gold h sulfur Naming elements part 2 Using a Table of Elements write the chemical name for each element and determine whether each element is a metal metalloid or nonmetal a Pb e Hg b K f As c Ag 9 Ne d Pt h Fe Identifying Ions using the Periodic Table Identify the type of ion cation or anion and charge 2 1 etc that is formed by each of the following elements a H d Ca b Cl e AI c Ne f O Naming the Ions Name the ion formed by the following elements use the periodic table if necessary a Na d Mg b I e AI c He f Hg Ch 100 Fundamentals for Chemistry Chapter 6 Nomenclature of Inorganic Compounds Naming Simple Compounds part 1 Write the chemical name systemic for each substance a N203 b MgClz e Agzs c CCI4 f PbOz Naming Simple Compounds part 2 Write the chemical formula for each substance a oxygen d ironIII chloride b lithium sulfide e copperI oxide c nitrogen dioxide f tricarbon octahydride Naming Polyatomic Ions Part 1 Name the following ions a CN39 d Hcog39 b NH4 e sof c Noz39 f OH39 Naming Polyatomic Ions Part 2 Write the chemical formula for the following ions a carbonate d sulfide b peroxide e nitrate c acetate f phosphate Ch 100 Fundamentals for Chemistry Chapter 6 Nomenclature of Inorganic Compounds Naming Simple Compounds part 3 Write the chemical name systemic for each substance a Fe2SO43 b NaNO3 e NH4st4 c Ca3PO42 f BaOH2 Naming Simple Compounds part 4 Write the chemical name systemic for each substance a potassium nitrate d calcium bicarbonate b ammonium hydroxide e copperII acetate c sodium hypochlorite f potassium cyanide Naming Acids part 1 Write the chemical name systemic for each acid a HCI d H2503 b HNo3 e HCIO C H3PO4 st Naming Acids part 2 Write the chemical name systemic for each acid a hydrofluoric acid d hydrocyanic acid b nitrous acid e acetic acid c phosphorous acid f carbonic acid Ch 100 Fundamentals for Chemistry Chapter 6 Nomenclature of Inorganic Compounds Reference Chart Useful Polyatomic Ions Ion Name NH4 ammonium N02quot nitrite N03quot nitrate 803239 sulfite SO42quot sulfate HSO439 hydrogen sulfate bisulfate OHquot hydroxide CN39 cyanide PO43quot phosphate HPO4239 hydrogen phosphate biphosphate H2PO439 dihydrogen phosphate C032quot carbonate HCOg39 hydrogen carbonate bicarbonate CIO39 hypochlorite CIOZ39 chlorite Clogquot chlorate CIO439 perchlorate C2H302 acetate 02239 peroxide Ch100 Fundamentals for Chemistry Common Polyatomic Ions Ion NH4 N02quot N03quot 503239 504239 Hsor OHquot CN39 PO43quot HPO4239 H2PO439 c03239 Hcog CIO39 00239 CIO339 CIO439 C2H3OZ39 MnO439 Cr207239 CrO4239 02239 Name ammonium nitrite nitrate sulfite sulfate hydrogen sulfate bisulfate hydroxide cyanide phosphate hydrogen phosphate biphosphate dihydrogen phosphate carbonate hydrogen carbonate bicarbonate hypochlorite chlorite chlorate perchlorate acetate permanganate dichromate chromate peroxide Pomanu amp Community College Ch 100 Fundamentals for Chemistry Chapter 5 Early Atomic Theory amp Structure Lecture Notes Ponland a Community Col lege Early Model of Matter Aristotle 384322 BC Introduced observation as an important step in understanding the natural world According to his model of nature all forms of matter are mixtures of one of 4 basic elements 1 Earth 3 Air 2 Water 4 Fire All matter has one or more of 4 basic qualities 1 Cold 3 Hot 2 Moist 4 Dry According to Aristotle Any substance could be transformed into any other substance by altering the relative proportion of these elements and qualities ie lead to gold Dalton39s Atomic Theory 1 Each element consists of individual particles called atoms Atoms can neither be created nor destroyed 3 All atoms of a given element are identical 4 Atoms combined chemically in definite 39 wholenumber ratios to form compounds 5 Atoms of different elements have different masses The Modern Atomic Model According to our modern model of the matter the atom has 2 primary regions of interest 1 Nucleus Contains protons amp neutrons called nucleons collectively Establishes most of the atoms mass l l 7 Small dense r ion at thecenter oi the atom The radius ot the nucleus if m l lerntorneter 2 The Electron Cloud 7 Contains electrons Mass oil electron 9 mg XlU k r Establishes the effective volume of the atom The radius oi the electron cloud if 0 rn l Angstrom r Determines the chemical properties of the atom tht ra h h occur of each atom The electron properties ot the atorn will detthe the types ot interaction that will ta e piace wonany cumum Structure of the Atom Electron cloud What holds the atom together Electromagnetic interaction aka electric force holds the electrons to the nucleus The negative charge of the electrons are attracted to the positive charge of the nucleus Strong interaction aka strong force holds the nucleons together within the nucleus The positive charge of the protons repel each other All nucleons protons and neutrons possess a STRONG attraction to each other that overcomes the protons mutual repulsion lammi community can Electric Charge Electric charge is a fundamental property of matter We don t really know what electric charge is but we do know that there are 2 kinds Positive charge Negative charge Opposite charge polarity is attractive attracts Same charge polarity is repulsive repels and repels The magnitude of electric charge q is the same for protons and electrons The charge of a proton qproton or electron qelectron is the smallest amount that occurs in nature it is called the quantum of charge 1 qpm on 1602 x 103919 Coulombs or 1 2 qelec run 1602 x 103919 Coulombs or 1 Fonlaua commun39 college Ions Atoms or molecules that have gained or lost one or more electrons ons that have lost electrons are called cations ons that have gained extra electrons are called anions Ionic compounds have both cations and anions so that their net charge is zero lammi community can Ions cont Ions are electrically charged atoms and thus carry electric charge The electric charge of an ion is due to the imbalance of electrons and protons When an atom has lost one or more of its electrons it carries a positive charge 1 for each electron that is lost When an atom has gained one or more of its electrons it carries a positive charge for each excess electron that is gained When an atommolecule is an ion its charge must be specified Sodium ion Nat Chloride ion Cl39 Hydroxide ion OH39 Notes on Electric Charge Opposite charges attract Like charges repel Fonlaua commum college Atomic Bookkeeping Atomic number Z The number of protons in an atom or ion The number that defines the identity of the atom Mass number A The number of protons amp neutrons in a specific atom or isotope The number that represents the mass of an atom To determine number of neutrons in an atom of neutrons Mass Atomic Or of neutrons A Z Mass vs Atomic Mass Isotopes are the equivalent of sibling members of an element 1 Unique atoms of the same element with different mass numbers ie they have different numbers of neutrons Unique isotopes are identified by their mass number Isoto 39e notation p Mai Atomlc Symbol Atomlc Example carbon12 1753 amp carbon14 14C 6 Atomic mass 1 The Isoto The unit of Atomic Mass is the Dalton or amu lton onetwellth mass of one ZC atom 1 661x1027kq Note There 6 protons amp 6 neutrons in a l2C atom but the mass of a l2C atom Is actually Sll hlly less than the combined mass of all of the nucleons individua Where is this lost mass It s rel V V eased as energy when the nucleons comblne blnd to form the nucleus of the atom average total mass of an element39s various naturally occurring pes N a a Q i proton l proton 1 proto 0 neutron 1 neutron 2 neutrons Itprotium deuterium tritium l Hydrogen isotopes 26 protons 26 protons 30 neutrons 29 neutrons l l Iron isotopes LEWIS STRUCTURES AND VSEPR Rules for Drawing Lewis Structures in priority order Calculate the total number of valence electrons from each atom in the molecule a positive ions deduct one electron for each positive charge b negative ions add one electron for each negative charge The number of valence electrons for atoms in groups I through VIII is equal to the Group Number The number of valence electrons for transition metal atoms is equal to the oxidation state of the atom Arrange the atoms appropriately a the central atom is usually the least electronegative b the molecule is usually symmetrical c the molecular formula is usually drawn indicating the order of atoms ie CH3CNO and CH3NCO have a different order of the carbon and nitrogen only d if you cannot obtain a reasonable Lewis structure considering the number of bonds formed for each type of atom as in the table above then merely try another arrangement of atoms Add lone pairs of electrons to satisfy the octet rule for each atom Atoms from Period 2 C N O F and higher take eight electrons Be care u o a Hydrogen H an atom from Period 1 takes two electrons only instead of the eight b The electron deficient atoms Be only takes a total of 4 electrons and B 6 electrons instead of the usual eight Total the electrons in your structure counting bonds sticks joining atoms as two electrons and lone pairs as two electrons a if the total electrons in your structure equals the total valence electrons then you have a good structure b if the total electrons exceeds the number of valence electrons then add a double bond 1 bond for each pair of electrons in excess of the valence Note Resonance occurs when there is more than one possible position to place the double bond 1 bond c if the total electrons is less than the number of valence electrons add electron pairs to the central atom until you have achieved the total number of valence electrons The central atom is the least electronegative and so is the most likely candidate for the extra electrons Atoms from Period 1 can never take more than two electrons H Atoms from Period II can never take more than eight electrons C N O F Thus the central atom will be from Period 3 or higher when you add electron pairs to it Hydrogen will only ever form one bond and for this course fluorine will only ever form one bond Formal Charge When you count the electrons around each atom in a Lewis structure by assigning one electron from each bond stick to that atom and assigning both electrons from a lone pair the total electrons should equal the valence electrons of that atom When this total differs from the number of valence electrons then the atom has a formal charge Because we are localizing electrons in a Lewis structure when in reality all the electrons are smeared over the Whole molecule not uniformly smeared some regions will have more electron that another this artificial localization of charge results However it does give us a reasonable picture of the electron distribution in a molecule a starting point for more sophisticated models Thus the formal charge on an atom in a molecule is given by Formal charge number of valence electrons number of electrons in lone pairs number of bonds Notes a an atom will only have a formal charge if the number of bonds it forms is more or less than its desired number b the sum of the formal charge in a molecule must equal the charge on that molecule including zero charge c the preference for carrying negative charge is O gt N gt d the best structure has the lowest or least separation of formal charge The Lewis structure with the least amount of formal charge is the more reasonable picture of the electron distribution in a molecule or the stable structure The formal charge in a Lewis structure may be minimized in a situation where an atom with a negative formal charge is bonded to an atom with a positive formal charge A pair of electrons on the atom with the negative formal charge is moved to form a double bond between the two atoms of opposite formal charge The effect of this relocation is to move an electron from the negatively charged to the positively charged atom thus eliminating the two charges The total number of electrons in the molecule remains unchanged the only difference is that two electrons have been relocated from a lone pair to a bond pair See the example below with ClOz e e Cgt g O Z1 O Resonance Resonance occurs when there is a choice of placement of a double bond 1 bond in developing a Lewis structure Each placement will give rise to a different Lewis structure and each of these Lewis structures will contribute to our picture of the real molecule The sum of these Lewis structures is called a resonance hybrid Each Lewis structure may contribute to the resonance hybrid to a different extent Resonance structures for a molecule differ only in the positions of their electron pairs The resonance hybrid is more stable has a lower energy than any of the contributing structures Less stable structures can sometimes be ignored altogether depending on their contribution Comparitive stabilities can often be evaluated on the basis of formal charge using the following rules a the most stable structure has the least formal charge b the preference for a negative charge is oxygen gt nitrogen gt carbon c structures in which adjacent atoms have the formal charges of the same sign are especially unstable Electronegativity Electronegativity is the ability of an atom to attract electrons to itself A Table of electronegativity values follows this section on Lewis structures Polar bonds generally have higher bond energies the energy necessary to separate two atoms than nonpolar bonds composed of the parent atoms HF has a higher bond energy than either H2 or F2 This is attributed to an ionic component in the HF bond Fluorine attracts the electrons to itself obtaining a partial negative charge and hyrogen being deficient in electrons obtains a partial positive charge Pauling devised a scale for electronegativity based on this difference in bond energies The electronegativity of uorine is defined as being 400 The difference in electronegativity AEN between two atoms is given as l AEN XA 7 XB 0102DAB 7 DAA gtlt DBB A where DAB is the bond dissociation energy between atoms A and B and X A is the electronegativity of A and XE is the electronegativity of B Calculate the electronegativity of hydrogen given the following data Molecule I HF I H2 2 I IRondDi nciatinnPnergvkJmol391 568 436 158 y 1 AEN XF ixH 0102568ii J436 x158 j Z 010230553 A 178 mol mol mol mol but the electronegativity of fluorine is defined as 400 and so the electronegativity of hydrogen is XH X1 7 AEN 4007178 222 Bond Polarities and Electronegativity Differences The greater the electronegativity difference between two atoms the closer the electron density is located to the more electronegative atom and the greater the degree of ionic character in the bond between them in other words the less electronegative atom develops a partial positive charge and the more electronegative atom a partial negative charge A covalent bond with ionic character is called a polar bond and the polarity of a molecule is assessed by experimentally measuring the dipole moment The dipole moment IJ is defined as the magnitude q multiplied by the distance d between their centres The ST unit is the Cm but a much smaller unit the debye is commonly used for molecules One debye CD 333 x 103930 Cm The dipole moment correlates directly with the electronegativity difference as shown in the table Thus we can use electronegativity differences to predict the polarity of molecules Bonds are classified into nonpolar polar and ionic Electronegativity differences can be used to make this classification but they are not infallible Valence Shell Electron Pair Repulsion Theory VSEPR The shape of a molecule can be predicted by assuming the regions of electron density around an atom will be as far apart as possible This assumption when applied to the valence electrons of a molecule is called VSEPR For now we will consider the central atom but later we will apply the theory to any atom in a molecule with more than one bond The regions of electron density when applied to the valence electrons are the lone pairs and bond pairs multiple bonds count as one bond pair Rules a draw the Lewis structure b count the number of lone pairs and bond pairs on the central atom c obtain the general shape from the table below d alter the shape for repulsion between lone pairs Number of VSEPR electron between VSEPR Pairs 2 180 linear 3 120 4 1095 tetrahedral 5 120 amp 90 6 90 octahedral Examples of all the general VSEPR shapes With bonding and nonbonding electron pairs HH H C C 0 A32 AB3 AB4 ABS Linear Triangular Tetrahedral Triangular bipyramidal CdBr2 Ba CH4 SnBr 1 PFs o o o o A132 E Angular or bent SnCl2 AB3 E Triangular pyramidal NH PCl3 Distorted tetrahedral seesaw cn5 O AB2 152 A8352 A84 E2 Angular Tshaped Square or bent planar H20 SCI2 XeF4 ABZ E3 Linear XeF2 Predicting the Polarities 0f Molecules The polarity of a molecule with more than one bond will depend on the individual dipole moments of each bond and the geometry of the molecule Molecular polarity is then the vector sum of the bond dipole moments It was shown above Bond Polarities and Electronegativity Differences that bond dipoles correlate closely with bond electronegativity differences and thus we are able to predict the polarity of a molecule from the vector sum of the electronegativity differences The electronegativity difference for a bond is treated as a vector with the magnitude of the vector representing the electronegativity difference and direction of the vector representing the orientation of the bond in the molecule The head arrow of the vector points towards the most electronegative atom in the bond and thus the resultant vector points toward the electron dense side of the molecule The examples below show BF3 where all the vectors cancel to give a nonpolar molecule and CHF3 where the resultant vector indicates a very polar molecule 035 143 178 The dipole moment for BF3 is 0 D and for CHF3 165 D AENFB 398 204 194 AENFC 398 255 143 AENCH 255 7 220 035 H l F The next topic deals with the magnitude and direction of the vector we assign to a lone pair of electrons on a molecule The direction of the vector is straightforward as the lone pair is considered a region of electron density in VSEPR We can estimate the magnitude if we compare the molecules NHg NC13 and N39Fg In each of these molecules we have three substituents bonded to a nitrogen and the nitrogen has a lone pair The object of this exercise is to estimate the magnitude of the electronegativity vector for the lone pair on nitrogen The shapes of each of these molecules is tetrahedral with the three hydrogen electronegativity vectors or chlorine or uorine forming an inverted umbrella with the lone pair above see a in the diagrams The three substituent vectors will form a resultant that is equal in magnitude to one of the substituent vectors but aligned directly underneath the lone pair vector see b and c in the diagrams In the case of the hydrogen vectors this resultant vector will increase the magnitude of the lone pair vector For a quantitative example of how to find the resultant of three pyramidal vectors in a tetrahedral molecule see 39Calculation of the resultant of the electronegativity vectors for the molecule ClF03 following AENN H 304 7 220 084 AENClN 316 7 304 012 AENGN 398 7 304 094 a b resultant m gt 4 l gt M gt If we assign the variable x to the magnitude of the lone pair vector and we assume that the dipole moment correlates directly with the resultant electronegativity vector for the molecule we can write the equation magnitude of lone pair vector x magnitude of resultant of the 3 substituent vectors dipole moment in debyes Diagram c illustrates the two electronegativity vectors that are being summed to give the dipole moment So the lone pair electronegativity vector on nitrogen can be taken as the average of these three values which is 062 But AEN le 7 XN 062 and since XN 304 thus XII 366 on nitrogen This is a reasonable value to apply to lone pairs on most atoms when using electronegativity vectors to assess polarities of molecules Calculation of the resultant of the electronegativity vectors for the molecule ClFO3 The resultant electronegativity vector for the molecule will pass along the uorinechlorine bond and through the centre of the pyramid formed by the three oxygens and the chlorine Thus the y axis will be in the direction of this resultant vector and the x axis will be through the chlorine and at 90 to the y axis in the same plane as the left side 0 Since the FClO angle is 1095 the angle between the OCl bond and the x axis is 195 1095 90 and the angle between the OCl bond and the y axis is 705 90 195 See the diagrams below AENOCl 344 7 316 028 AENFCl 398 7 316 082 Each OCl electronegativity vector will contribute an amount h to the resultant vector Cos705 AENO 7 Cl h 028Cos705 009346 The contribution for the three OCl vectors to the resultant is 3h 0280 This is the same as one OCl vector which must be the case if you think about it for the molecule C104 the resultant of three OCl vectors must be equal and opposite to the fourth as the molecule is nonpolar The resultant vector is then AENFCl 7 3h 082 7 028 054 in the direction of uorine F F F y axis y axis I 10950 1 xaxis C xaxis xaxls C o 1950 195 h O o O 7050 0 04 O O O O Electronegativity Values 1 2 H He 2 20 3 4 5 6 7 8 9 10 L1 Be B C N O F Ne 098 157 204 255 304 344 398 11 12 13 14 15 16 17 18 Na Mg A1 Si S C1 Ar 093 131 161 190 219 258 316 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc T1 V Cr M1 Fe Co Ni Cu Zn Ga Ge As Se Br Kr 082 100 136 154 163 166 155 183 188 191 190 165 181 201 218 255 296 29 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Zr Nb M0 Tc Ru Rh Pd Ag Cd 1n Sn Sb Te 1 Xe 082 095 122 133 16 216 19 22 228 220 193 169 178 196 205 21 266 26 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba La Hf Ta W Re Os 1r Pt Au Hg T1 Pb B1 Po At Rn 079 089 110 13 15 236 19 22 220 228 254 200 204 233 202 20 22 87 88 89 Fr Ra Ac 0 7 0 9 11
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