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UCSB / Chemistry / CHEM 07534 / What do you think the phrase cognitive elite actually means?

What do you think the phrase cognitive elite actually means?

What do you think the phrase cognitive elite actually means?


School: University of California Santa Barbara
Department: Chemistry
Course: Chemistry 109A
Professor: Paula bruice
Term: Fall 2014
Tags: Chemistry
Cost: 25
Name: Chemistry 101 Notes - ALL
Description: These are all of my notes for my Chemistry 101 class. Enjoy!
Uploaded: 03/07/2017
323 Pages 548 Views 1 Unlocks

Contents Notes Page Notes Page Why and What is Chemistry? 2 Gases 152 The Components of Matter 13 Thermochemistry 169 Measurement 25 Modern Atomic Theory 193 Dimensional Analysis 48 Modern Atomic Theory - Apps 212 Atomic Theory 1 55 Periodic Trends 228 Molecules and Ions 69 Chemical Bonding Basics 244 Stoichiometry 1 86 More Chemical Bonding 259 Quantitative Stoichiometry 99 Intermolecular Forces 282 Aqueous Reactions 120 Exam Tips and Final Review 296 Acids - Bases Reactions 134 Blank Practice Exams 301 REDOX reactions 143 Periodic Table 322

Why and What is Chemistry?

See the Course Website (http://www3.jjc.edu/staff/pmills) for specific test dates and other important  information Legend You will often find specific icons embedded within the notes. These respective  symbols alert the student to the following: Represents a key fact or other piece of information, such as the  definitions of an element and a compound. Represents a useful trick the student will likely find useful, such  as an 'EZ' way to convert between grams and moles for a  substance Alerts the student to an important relationship between micro and macro scale properties or phenomena with respect to the  material under discussion * Such material provides a link to interesting (briefly discussed)  supplemental material, often beyond the scope of the course  syllabus

“What’s my motivation?

Don't forget about the age old question of isu ois

Disclaimer: This document may only be downloaded, without charge, by students enrolled in Dr. Mills’  CHM 101and CHM 102 course(s) at Joliet Junior College. This document MAY NOT be resold, or in any  other way utilized for profit, by any third party interest. Cover Art: The ‘Fundamental Sun’ (Atom) STM image 1 Why Chemistry? “What’s my motivation?” Why are you sitting in this class? In other words, why chemistry? Task: Write down as many reasons as you can that explain why you are  taking this class: (We will also justify responses relating to 3rd party requirements during the session) Professional programs that benefit directly from a background in chemistry 1. Nursing and allied health (pre-pharm., pre-med., pre-dentistry) Example: Chlorothiazide (Diuril) is ordered b.i.d. for a infant weighing 6.5 kg. It is  supplied in elixir form 100 mg/tsp. The recommended dosage for Diuril is 25 mg/kg/day.  How many cc’s should the nurse give to the child for each dose? A. 6.15 cc. B. 8.13 cc. C. 4.06 cc. D. 0.81 cc.2 2. Engineering (mechanical, civil, chemical, electrical) Example: Your company decides to import child safety seats manufactured in Asia.  Unfortunately, the safety guidelines for the seats are quoted in ‘metric’ units. The label  reads: “Do not exceed a 150 N load” and you must use this information to determine the  maximum weight a child must not exceed in order to be protected during a collision at 55  mph. Can you do it? A child’s life, not to mention the financial future of your employer,  may depend on your ability to solve questions such as this. 3. Everyday / Real life situations Example: It is time to re-carpet your 12 ft x 24 ft. family room. You visit a few carpet  stores and select a brand that costs $ 20.50 per square meter. The sales person quotes you  a total price of $749 – is this price fair, or have you just been taken advantage of? ** We will return to and solve each of these three problems at some point during the course Discussion:  What do all three of the above examples have in common? Which professions (or professionals) utilize such skills most commonly? Hint, “I pretend to be  one on occasion” “We’re the first link”3 The “Cognitive Elite” Discussion: What do you think the phrase “cognitive elite” actually means? Data from ‘The Emergence of a Cognitive Elite’ (Chapters 1 and 2 of The  Bell Curve). ∙ People with IQ’s of > 120 (the top 10%) preferentially enter the 10 or so ‘High IQ  professions’ discussed above.  ∙ Developing good cognitive skills is essential to entering and being successful  within the ‘High IQ’ and related professions. We are the first link in the chain Example: Medical careers Increased Problem Solving AbilityTake home message: People with good cognitive / problem solving skills  preferentially find employment within fields of their choosing that are  financially rewarding and/or intellectually satisfying.  A question of some importance: How can one’s cognitive skills be  improved? Answer(s):  4 The Role of Chemistry as a Prerequisite Course Key facts and results: Fact: The problem solving skills routinely utilized in the ‘high IQ’ and  related professions (such as nursing, business management, accounting,  etc.) are introduced, learnt and mastered during physical science  courses. Result: Professional programs and subsequent employers insist that their  candidates have a background in one of the physical sciences – both for  specific (allied health, engineering) and general (your family room carpet)  reasons.  Fact: Study within any of the ‘high IQ fields’ will increase cognitive  skills, but only the physical sciences do so via the study of fundamental,  everyday phenomena so are of broad relevance and interest (we all  interact with and benefit from the manipulation of matter on a daily basis  after all).  Result: Chemistry (and physics) may be considered to be the ‘gatekeepers’  of cognitive learning – chemistry in particular introduces, develops and  subsequently equips students with cognitive skills necessary to succeed  in their chosen careers Take home message: While the direct relevance of chemistry to your chosen  course of study may at times seem tenuous, remember that the cognitive  skills developed during such programs of study are of significant importance  to your professional development and employability. In essence, this is why  you are here.5 How Chemistry is Perceived & Skills Needed to Succeed in  Chemistry How Chemistry is Perceived: Discussion: How did your friends and family respond when you told them  you were taking a chemistry course this semester?? [“Frank” slide] Study Skills Needed to Succeed in Chemistry: Fact: As discussed above, chemistry is all about the student developing  and learning to apply problem solving skills - your study habits should  reflect this. Do NOT fall in to the trap of believing you can learn chemistry  simply by memorizing the information from your text – you must practice  applying this information, not just be familiar with it.  Result: Successful chemistry students typically spend most of their  independent study time working assigned problems, not just reading about  them. To learn chemistry you must do chemistry is a truism worth  remembering. An analogy would be this: you read all the books out there on  the subject of golf, but don’t get round to swinging a club – what do you  think happens when you tee off for the first time? Fact: Chemistry relies on a cumulative method of learning, i.e. theories  learnt from week 1 onwards will be repeatedly applied all the way through  the course. Thus, it is important that the student does not let any ‘gaps’  in their knowledge develop. This fact exemplifies the differences in  philosophy between the sciences and arts, as art courses are often more  modular in nature. Example: I overhead a student tell another: “Yeah, I blew  off reading the first book in my English class, but read the second one and  got a ‘B’”. This method of study is not recommended in chemistry!6 Analogy: Building a tower Result: Successful chemistry students typically have exemplary  attendance records. In some cases they may not be the ‘best’ students, but  guarantee themselves a better grade than more capable students, who in turn  typically may miss as few as one or two lecture sessions (this is especially  true with regard to 3 hr. class sessions). Pictorial analogy of attendance vs cumulative knowledge ‘I missed a lab’ ‘I missed a lecture’ ‘I missed a couple of lectures’ Don’t ‘Swiss cheese’ or ‘torpedo’ your chances of passing the course  because of missed work!

Why are you sitting in this class?

Don't forget about the age old question of What makes up the fluid mosaic model?

Take home message: Simply by attending class regularly and completing  the HWK assignments you essentially guarantee yourself a passing  grade for the course, while, due to the nature of the material, deviating  from this approach may ensure the opposite7 What is chemistry? What do Chemists do? Task: In your own words describe what you consider chemistry to be, plus  make a list of what you think the job of a chemist is: What is chemistry? “Official” definition of what chemistry is:8 Key words: Matter: “Stuff” – anything with mass and volume. Can you think of  anything without mass or volume?   What are the basic building blocks of all matter, be it a diamond, a tree or  the air around us? ‘High Tech’ science (STM or AFM, top left) is often based on simple ideas  (gramophone, top right). Click logo for ‘flyby’. More recent atomic (STM) images9 Example: What is water made up from? How do you know? Summary: Atoms and molecules are MICROSCOPIC particles (they are very  small) A drop of water is a MACROSCOPIC particle (because you can see it,  hold it in your hand etc.)10 What do chemists do? In other words, what is the basic most, fundamental  goal of every chemical investigation? Hint: Think how chemists express  their findings.. “Official” definition of what chemists do:11 Chemistry in action: Explaining what happens on your  BBQ grill.    The burning of a charcoal brick on your backyard grill  (MACRO) explained in terms of a balanced chemical  equation (MICRO) ANY large (MACRO) scale chemical process can be  described using a MICRO scale chemical equation  featuring individual atoms and/or molecules

We also discuss several other topics like How is the body organized?

A cartoon representation of the reaction of the pertinent atoms and  molecules; along with the Chemists’ description – a balanced chemical  equation illustrating a single microscopic event.  Cartoon Balanced Equation This process is repeated many billions of times (MICRO) for the burning of  a charcoal briquette (MACRO)12 The Components of Matter Reading: Ch 1 sections 1 - 5 Homework: Chapter 1: 37, 39, 41, 43, 45, 47*, 49

* = ‘important’ homework question Review: What is matter? Recall: “Chemistry is the study of matter and its properties, the changes  matter undergoes and the energy associated with those changes” Recap: There are 3 stable states of matter – solid (s), liquid (l) and gas (g). 13 Specific macro- and microscopic physical properties define the three*  states of matter  State of Matter  Macroscopic Description (observation) Microscopic Description  (chemical model) Solid  



If you want to learn more check out How to convert parametric to rectangular?

The state matter is in depends on the strength of the forces  (chemical bonds) between the individual microscopic  particles within the matter

If you want to learn more check out orrorin tugenensis cranial capacity

Task: Rank the intermolecular forces present in steam, ice and water in  order of increasing strength. Use the included figures as a guide. Ranking14 Changing between the 3 states of matter Describe the relationship between the mpt. and bpt. of matter,  with regard to microscopic processes, occurring at these  specific temperatures

Example: The boiling of water to make steam ( H2O(l) →( H2O(g))15 Physical and Chemical Properties – what’s the difference? Analogy: We all posses ‘as is’ physical properties, or  characteristics, that define us. For example, Dr. Mills is 5’11”  and has green eyes.  As with people, each chemical also possesses a unique set of ‘as is’ physical  properties that define it. For example, water is a clear, colorless, tasteless  molecular material that has a fpt. of 0oC and a bpt. of 100 oC. Chemical Properties, in contrast, are a function of change (usually  associated with a chemical reaction). For example, Iron (Fe) reacts with  oxygen gas to form rust: 4 Fe (s) + 3 O2(g) → 2 Fe2O3 (s) Task: Identify the flowing as either chemical or physical properties Property Chemical or Physical Diamond is the hardest known  substance. Charcoal burns to make CO2 (g) The statue of liberty turned ‘green’ Copper is a good conductor of  electricity Sugar dissolves in water* Melting of ice*

Don't forget about the age old question of What is the goal of an institution?

Think up two more chemical properties of your own16 Elements and Compounds – the further classification of pure matter Task: State which of the following are elements, and which are compounds.  When done, try to come up with a definition of what elements and  compounds are. Material  Chemical Formula Element or  Compound? Water H2O (l)

Oxygen gas O2 (g)

Pure silver coin Ag (s)

Sugar crystals C6H12O6 (s)

Carbon dioxide gas CO2 (g)

Elements:  Compounds:

17 Compounds and elements can have either ‘giant’ or molecular structures:‘Giant’: Repeating lattice of particles – usually  strongly bound (high mpt.) solids.  Examples: sand (SiO2), diamond (C), table salt  (NaCl)

Molecular: a collection of independent molecular  units (molecules will be discussed in more detail  later). Usually (low mpt) liquids or gasses at room  temp.  Definition: Molecule – a small, independent  particle of matter made up from 2 or more atoms Examples: water (H2O), carbon dioxide (CO2),  Nitrogen gas (N2)

Think of molecules like cars on the expressway – each car  (molecule) is a separate, independent unit that contains a  number of passengers (atoms). The cars (molecules) are  free to move while the people (atoms) stay fixed inside.   ‘Giant’ materials are like people (atoms) ‘locked’ in place  at a very crowded concert, the DMV waiting room  etc……

18 Review: A microscopic scale view of several materials is presented below.  Label each using elemental or compound and molecular or ‘giant’ tags Water (H2O (l))  Silicon (Si (s)) Steam (H2O (g)) Sodium Chloride (NaCl)

Details: Ice is a solid (crystalline) form of water (a  molecular compound). How would you describe  the structure of ice? Can you think of other similar  examples? More Details: Allotropes of an Element  Example:  Carbon C(diamond) C(graphite) C6019 Pure Matter v Mixtures Recap: Pure matter is classed as either an ELEMENT or a  COMPOUND. Elements can have either Molecular or ‘giant’ structures. Examples: N2 (g) (Nitrogen gas, molecular), Pb(s) (metallic  lead, a ‘giant’ structure)  Compounds can also have either Molecular or ‘giant’  structures. Examples: H2O(l) (water, molecular), Fe2O3(s)  (‘rust’ (iron oxide), a ‘giant’ structure)  Recall: A molecule is an independent unit containing two or  more atoms. Remember the car / passenger analogy.  Molecules can exist as either elements or compounds

Mixtures ANY combination of different types of pure matter ‘placed  together’ is defined as a mixture (eg. air, milk, pepsi).  Mixtures are NOT pure materials. eg. Pure gold (Au) vs  ‘white’ gold (Au+ Ag), or water (H2O) vs pepsi (H2O +  sugar….)

Discussion: Air contains a number of different components – what are they?  How would you describe what air is made up from using words like element,  compound, gas, molecular etc.?20 Task: Assign generic labels that describe to microscopic scale matter shown  on the slide (e.g. ‘gaseous atomic element’ etc.)  Mixture Types As viewed from a macroscopic perspective, mixtures are classified  as either HOMOGENEOUS or HETEROGENEOUS

HOMOGENEOUS MIXTURES: Examples: HETEROGENEOUS MIXTURES: Examples:21 Discussion: Can you think of something that is both a homogeneous mixture and a solid? A Bronze statue of  Caesar Augustus Examples of Alloys: Classification of Matter Flowchart (Dr. Mills really likes this slide – why? Hint: Recall the fundamental job  of a chemist)22 Task: Use the ‘Classification of Matter’ flowchart (above) to classify the  following: 1. The compressed gasses in a deep sea diver’s gas bottle (He(g) and  O2(g)) 2. A ham and cheese omelet 3. An ice cube (made from pure water) 4. A ruby (Al2O3(s) with Cr3+ impurities) Extra Credit: Ask me about the separation of mixtures  assignment (based on background reading)23   “Mixtures, Elements and Compounds” The following questions were taken from your 1st practice  midterm:

State whether the following are classified as elements, compounds or mixtures**: Diamond: Carbon dioxide gas: Air: A cup of coffee: Water: Sand (SiO2): Oxygen gas: **include additional details for extra credit!24 Units of Measurement Reading: Ch. 1 sections 6 & 7 Homework: Chapter 1: 51, 55*, 57*, 59, 61, 65,  67, 69*, 71*, 75, 77*, 81, 83*

* = ‘important’ homework question Common Units Discussion: List some common units of measurement we use on a daily  basis. How did these units originate? Quantity measured Familiar Unit Mass Question: What are the ‘metric’ (S.I.) versions of the everyday units listed  above?    Quantity measured Fundamental S.I.  Unit (base unit) Symbol Notes: SI base units are used to determine derived S.I. units, as discussed  below. Some S.I. base units feature a decimal prefix – which one(s)?25 Derived S.I. Units Insert appropriate S.I. base units into an equation that  defines the respective derived S.I. unit desired. Example: Area = length x length = m x m = m2 ⇒ the derived S.I. unit for area is m2 Determine derived S.I. units for the following quantities  Quantity measured Math involving S.I. base units Derived S.I. unit Volume

Velocity (speed)




*These are harder examples. To solve them start by inserting appropriate S.I. base units  into an equation that defines the quantity sought.  Discussion: Why do scientists prefer the S.I. system?

26 Questions:  Is the S.I. unit of volume (m3) reasonable for everyday applications? Why? What unit of volume do chemists prefer? Why?

1 dm3= 1 L More detail on the chemist’s volume unit27 Scientific Notation and S.I. Prefixes Large Quantities Fact: Chemical problem solving most often involves using either very large  or very small numbers (e.g. counting the number of molecules in a drop of  water, or quoting the mass of the water drop in kilograms) Recall: How many individual H2O (l) molecules are there  in a drop of water. Write this amount as a regular number: Number H2O (l) molecules in 1 drop water = _____________________________ Problem: How do we represent and manipulate such quantities in an ‘easier’  way? Answer: Overview Example: Consider the statement “eight million people live in  London”. How can this quantity be best expressed numerically?

‘Everyday’: ‘Better’:28

Examples: Write the following quantities using regular numbers and powers  of 10 (scientific notation). Try to do this without a calculator at first, then see  the below tip for how to do this with your calculator’s SCI button

Just move the decimal point to the left until you get a  single digit with decimals. The power of ten is the number  of places the decimal point moved. Example:  3000 = 3 x 10(number decimal places to left moved) = 3 x 103 Quantity ‘Regular’ quantity ‘Power of ten’  quantity (SCI) One hundred miles

One thousand students

Five million people

Twenty million dollars

Five and a half billion people

29 TIP: Scientific notation (SCI) is different than the powers of 10 used in  engineering (ENG)*. When converting to SCI powers of 10 from a ‘real’  number press the SCI button on your calculator, or put it in SCI mode and  

press the = key.  

Wrap up: quote the number of H2O molecules in 1 drop water using SCI  notation: 1,000,000,000,000,000,000,000 molecules = ______________ molecules REMEMBER: In chemistry we ALWAYS use scientific notation (SCI)  for expressing large (>100) or small (<0.1) numbers. ________________________ Small Quantities Question: How can very small numbers be expressed in SCI notation? Just move the decimal point to the right until you get a  single digit with decimals. The negative power of ten is the  number of places the decimal point moved. Example:  0.00125 = 1.25 x 10- (number decimal places to right moved) = _________________ Examples: Convert the following regular numbered quantities to powers of  10 (scientific notation). Try to do this without a calculator at first, then check  with your calculator.30 ‘Regular number’  (quantity) ‘Power of ten’ (SCI)  quantity 0.00015 grams

0.125 %

0.0458 mL

Review: You now know how to convert large or small ‘regular’ numbers  into SCI notation either on paper or using your calculator. ________________________ Manipulating Large and Small Numbers Enter the following SCI numbers into your calculator using the EE or EXP 

keys to express numbers in SCI notation, then press the = (in FLO mode)  to obtain the ‘real’ number equivalent:  

‘Power of ten’ number (SCI) Regular number (quantity) (FLO) 5 x 10-1

1.5 x 103

3.56 10-3

Determine: 3 x 107÷6 x 103 = ___________________ What answer did you get? Where any problems were encountered?31

Making things even simpler – S.I. Prefixes Certain powers of 10 can be replaced by a symbol known as a decimal (or S.I.) prefix32 For decimal (S.I.) prefixes, just swap the appropriate “x  10n” part of the number for the equivalent prefix’s  symbol. Example:  1.25 x 10-3 g = 1.25 mg (milligrams) Task: Convert the following quantities to SCI notation and decimal prefix  notation: Quantity With SCI notation With Decimal Prefix 0.0015 grams

0.0000020 meters

30,000 dollars

12 million people

Discussion: Make a list of as many ‘everyday’ quantities as possible that use  decimal prefixes (or similar related expressions):33 Significant Figures and Rounding Off Question: What are significant figures?

Task: Measure the length of your pencil (or some other object) in cm using a  standard ruler. To how many sig. figs can you determine this value? Object Size measured (cm) Number sig. figs.

Let’s figure out the rules for sig. figs. What is: 1.002 to 3 sig. figs. 1 sig. fig. 569.74 to 3 sig. figs. 4 sig. figs. 1 sig. fig. 0.00017 to 1 sig. fig.34 Zeros before the first number are NOT counted as  significant Zeros after the first number ARE counted as significant Round UP if the number after the last significant digit is  > 5 Quote numbers in SCI notation if number sig. figs. <  digits before decimal point. Multiplication and Division (99% of your work is features one and/or  the other*) ‘You are only as good as your worst measurement’   The result of any multiplication or division has the same  number of sig. figs. as the measurement with the lowest  number of sig. figs.  Example: Lead has a density of 11.4 g/cm3. What volume (in cm3) does 2.1g  of lead have?35 Uncertainty in Measurement ‘Statistical’ Precision and Accuracy – group measurements Discussion: What is the difference between precision and accuracy for a  series of measurements? Can you be precise and at the same time  inaccurate? Think of a dartboard or other target when remembering  the differences between precision (‘grouping’) and  accuracy (‘hitting the bulls eye’)* Notes Notes Notes36 This picture was found at a private individual’s  internet site. The author states, among other  things, that he “would make a bad murderer.”  How true, he fired 10 shots from 21 feet, (the  target line was set at 75 feet) and hit the target 7  times.  How would you describe the spread of his shots  in terms of accuracy and precision? Formal definitions Accuracy: How close individual measurements are to the ‘true’ (or  ‘real’) value.  Accuracy is increased by repeating measurements (trails) and determining  their average or mean. (Average value → true value for large number of  trials) Precision: How closely individual measurements agree with one  another.  Standard deviation (or variance) is a relative measure of how precise a set of  measurements are with respect to their average. Precise measurements are not necessarily accurate – this is known as  systematic error: Example: Suppose a 2.0 mg grain of sand is accidentally dropped on to an  analytical balance. Unless the balance is recalibrated, it will weigh ‘too  heavy’ by 2.0 mg for each subsequent weighing. This is an example of a  systematic (poor accuracy) error. Note: Quantities determined from a series of measurements are most often  quoted as their mean ± standard deviation for that data set. For large data  sets (>20 measurements), this is equal to the true value ± error.37 Individual measurements For each individual measurement (mass, vol. etc.), its accuracy is typically  quoted as to as good as the last measured (significant) digit. This fact is  assumed for most physical measurements. Example: For a mass reading of 21.245g, the last ‘5’ is uncertain so the  accuracy of the reading is 21.245 ± 0.001 g.  Real Life: Dr. Mills owns a digital bathroom scale,  which quotes weights to the nearest 0.1 pound  (nice?). However, the scale often displays different weight values when a person steps off, then  immediately steps back on to, the balance (not  good), e.g. 212.4 lb, 212.7 lb, 212.5 lb.    How would you describe the performance of Dr.  Mills’ scale, in terms of any individual  measurement, with regard to accuracy and  precision?38 Temperature Background: There are three temperature scales in common use today. Can  you name them?   How were the end points of the two ‘metric’ scales defined? In other words,  what natural conditions define these respective temperature values? The Centigrade and Kelvin Scales The Centigrade scale compared to the state of H2O39 Converting between Degrees Celsius and Kelvin Task: By looking at the above figures, describe how the oC and K scales are  related. What do they have in common? What is different? 1. 2.

Simply add 273.15 to ANY temp. quoted in oC to obtain  the equivalent K value  OR Simply subtract 273.15 from ANY temp. quoted in K to  obtain the equivalent oC value

Examples: 1. What is 50oC in Kelvin? 2. What is 200 K in Celsius?40 Comparing the Fahrenheit, Kelvin and Celsius Temperature Scales Discussion: We saw that the end points for the oC scale  corresponded to specific ‘natural’ temperatures – the same  is true for the oF scale. What ‘natural’ temperatures do you  “You want to put what where?!..” Notes: think 0 oF and 100 oF (originally a centigrade scale)  correspond to in nature. How about 212 oF and 32 oF? Diagram: Fahrenheit, Celsius and Kelvin thermometers side by side. 41 Task: By looking at the slide provided, describe how the oC and oF scales are  related. What do they have in common? What is different? These two basic differences between the oC and oF scales  allow for equations relating them (conversion equations)  to be constructed:

For converting oC to oF: For converting oF to oC: Task: Convert the following Celsius temperatures to oF and K oF K Then ask me about the extra credit temperature….42 Density Review: How, in everyday words, is the property of density defined?   Density:

Where: ‘amount of matter’ = _______________ Discussion: What is the S.I. unit of density? Is this a convenient unit? ⇒ Density = __________________________________ Question: What are the two ‘convenient’ derived S.I. units of density used  by chemists? Ask me about the extra credit density….43 Density Math Recall: Density is defined by a simple equation, which has  three related forms: 1. 2. 3.

If you have problems with cross multiplication,  remember that simple ‘pyramids’ can also be used to  solve density and other 3 variable equations:

Example: 23.5 mL of a certain liquid weighs 35.062 g. What is the density  of the liquid? What mass will 20mL of this liquid have?44 Density of regular shaped objects Regularly shaped objects (cubes, ‘bricks’, spheres,  cylinders, cones….) have equations that define their  volume.

Task: Sketch the following 3-D shapes and list the equations that define their volume (see  your text book) Sketch of 3-D Shape Volume equation Cube V = ‘Brick’ V = Sphere V = Cylinder or disk V =

1. Find the volume of the object in question via the  equation that defines its volume (be sure to use cm for all  length dimensions). 2. Substitute the derived volume value in D = M/V to find  the object’s density (recall that mass is measured in  grams).

Recall: the radius of a circle equals half of it’s diameter (i.e. dia.= 2r)45 Example: Dice used in Las Vegas weigh 2.65 g and have  sides of length 1.2 cm. What is the density of a Las Vegas  dice? What relationship would you expect to exist between the  density (macro), and associated physical state (macro), of any  material with respect to spacing of its component particles  (micro)? Exceptions??

Densities of Some Common Materials   46 “Will it Float?” The David Letterman Show on CBS often features a segment  called ‘Will it Float’. Simply, Dave and Paul try to determine  if an object, such as a refrigerator or 100 ft of insulation  cable, will float when dropped into a large container of  water.  Question: What physical property of a material will determine ‘if it will  float’? What would be a more scientifically accurate (if less catchy) name  for the ‘Will it float’ segment on Dave’s show?

Discussion: “Battleships and dating  advice” Task: Using the table supplied above, sketch a picture of what would happen  if ~30 mL samples of ethanol, mercury and water, as well a lead ‘BB’ and a  gold ring were added to a volumetric cylinder.47 Dimensional Analysis (Conversion Factors)  Reading: Ch 1 section 8 Homework: Chapter 1: 89*, 91, 93, 95, 99, 101,  119

* = ‘important’ homework question Background We do simple conversions between different units on a daily basis. For  example: Question:    How many eggs are there in 1 dozen eggs? _________ A statement such as this can be written as an identity  1 dozen eggs = 12 eggs

Recall: Is the above identity a measured or exact relationship? How would  its use affect the number of significant figures used in any answer? Using Identities and Conversion Factors Question: How many eggs are there in 42 dozen eggs? Use the appropriate identity to create a conversion factor.  The conversion factor will transform the quantity into the  desired form.

Math:  42 dozen eggs x = _______________48 Conversion factors are simply identities written as  fractions. Each conversion factor has two ‘versions’

Task: Complete the following table by transforming the stated identity into  its two corresponding conversion factors. Also include at least three  additional conversion factors that you have encountered. Identity Conversion factors Exact? (Y/N) 1 in = 2.54* cm

1 kg = 2.205 lb

1 m = 100 cm

1 ft = 12 inches

49 Discussion:  How do you know which version of the conversion factor to use? Why? For none exact identities and conversion factors, how many sig. figs are  implied? Example: Use the information from above to determine how many cm there  are in 12.00 inches.  12.00 inches x = _______________ The unit belonging to the quantity and the denominator  of the conversion factor cancel to leave a final answer  with the desired unit

Generic Form: Extra Credit (3 pts., typed for next time): Dr. Mills gets very cranky with the  text as it says “1 in = 2.54 cm exactly”. What is the source of Dr. Mills’  confusion? Hint: Is the inch really an S.I. unit? Try searching for  ‘International inch’ to get started.50 Task: Complete the following conversions. See your text for appropriate  conversion identities. 5.51 cm to meters 23.0 ounces to pounds 50.0 nm to meters* 6.56 miles to km 45.7 inches to cm 220 pounds to kg # of drunk guys from 20  beers51 Conversion Factor ‘Chains’ Question: How do you approach a problem like “Convert 55 cm into feet” – where there is no available single conversion factor? Answer: 55 cm x x = _______________ Link as many conversion factors as necessary together in  order to crate a ‘chain’. Each ‘link’ in the chain converts  one unit to another and so on until the answer is reached

Task: Complete the following ‘chain’ conversions. See your text for  appropriate conversion identities. 4.00 ounces to grams 1.68 m to inches 5.8 km to feet52 Question of the week: How many atoms would have to be placed end to end  in order to reach Chicago. Assume an atom is 0.15 nm wide and Chicago is  40.0 miles away. Conversion factors: ‘Chain’ Math:53  “The Wire” The following questions were taken from your 1st practice  midterm:

Question 1a (20 points) A spool of copper (Cu) wire has a mass of 2.00 pounds and a  diameter of 50.0 μm. Determine the wire’s mass, volume and length in the units  specified below. Include any appropriate decimal prefixes in your final answers. Assume density copper (Cu) = 8.95 gcm-3 Mass of the wire in kg: (ANS = 0.907 kg) Volume of the wire in cm3: (ANS = 101 cm3) Length of the wire in meters:  (ANS = 5.14 x 104 m)54 Atomic Theory – Part 1 Reading: Ch 2 sections 1 – 6, 8 Homework: Chapter 2: 39, 47, 43, 49, 51*, 53,  55, 57, 71, 73, 77, 99, 103 (optional)

* = ‘important’ homework question The Atomic Theory (John Dalton, 1803) Dalton revisited the Ancient Greek Philosophers’  (Democritus et. al., 460 BC) ideas pertaining to how all matter is constructed from very small indivisible particles  (‘atomos’).   Dalton formulated a set of ideas (postulates), known as  “The Atomic Theory of Matter”, that would (~100 years)  later be shown to be correct Postulates of Dalton’s Atomic Theory of Matter 1. Matter is composed of extremely small particles called atoms 2. All atoms of a given element are identical, the atoms of each  element are different and have different chemical and physical  properties 3. Atoms are not changed into different types of atom(s) via  chemical reactions. Atoms can neither be created nor destroyed 4. Compounds are formed when atoms of more than one type are  combined. A compound always has the same relative number and  kind of atoms Notes on Dalton’s Atomic Theory55 Atomic Structure Discussion: In the introductory lectures we took a brief look at different  types of matter (i.e. elements, compounds and mixtures). We know these  materials are made from the smallest stable units of matter, atoms.  Atoms themselves are in turn made from smaller unstable particles – recall  as many of these fundamental ‘building blocks’ of matter as you can:  Question: How are all of these fundamental ‘building blocks’ of matter  related? Sketch a flow chart: Fermi Lab, located in West Chicago,  IL, is the world’s largest ‘atom  smasher*’. Fermi is where scientists  perform experiments in an attempt to  understand the origins of the universe56 Example: Water   In many ways we take atomic theory, as well as its eventual  confirmatory experimental results, for granted (i.e. we don’t think  about where it came from, we just assume and use it).    At the turn of the 19th century and for the next ~ 100 years this  work was a the cutting edge of scientific research and was  pursued by some of the world’s greatest scientific minds  

Task (complete outside of class): Complete the reading assignment for this  note packet. Make notes on the following topics and make reference to the  included illustrations in your discussions. This material, as well as other  similar assignments, will likely form the basis of any extra credit questions  appearing on midterm exams.  57 Cathode Rays and electrons (J.J. Thompson) Notes The Nuclear Atom (Rutherford)  Notes58 Atoms and Isotopes Review: Atoms are the smallest type of stable matter, they are typically  spherical and have diameters of ~ 0.18 – 0.60 nanometers.   Shown on the left is an STM image of a  silicon chip’s (Si (s)) surface. Note that it  has a repeating ‘giant’ structure. Question: Based on the scale, what is the  approximate width of a silicon atom in  nm? Answer: Ask me about the extra credit magnification…. The ‘Classical’ view of atomic structure Questions:    1. What is found at the center of an atom?  2. What two different types of subatomic particle are found inside the  nucleus? (subatomic means ‘smaller than’ atomic) 3. What ‘orbits’ the nucleus? 59 4. Sketch a generic diagram of an atom using the slide as a guide. Based  on the slide, how many times smaller is the diameter of the nucleus  than the atom as a whole? Comparison of subatomic particles (i.e. the things atoms are made from) Particle  Symbol Charge  Relative mass Electron e -1 1 Proton p +1 1836 Neutron* n 0 1839

* ask me to tell you a very poor neutron joke - it starts with ‘a neutron walks into a bar’….Electrons are much lighter than the neutrons and protons (that, in turn,  ‘inhabit’ the nucleus) ⇒ ELECTRONS MOVE MUCH MORE  QUICKELY THAN THE NEUCLIUS EVER CAN (this is called the  Born – Oppenheimer Approximation). This is why electrons are said to either ‘orbit’ the nucleus or exist  as ‘blurred out’ electron ‘clouds’. This is the main difference  between the ‘classical’ and ‘modern’ models of atomic structure. We  will study the modern ‘electron cloud’ model in depth later in the  course.

60 Question: Are ATOMS* electrically charged? Answer: ______ Question: What then must be true for EVERY atom in terms of the number  of electrons and protons it contains?

*Aside: We saw/will see in lab that ions are made by electrically charging atoms or molecules, we will  study this concept in more detail later. Question: Where can we find out the number of protons, Z, (and therefore  also the number of electrons) an atom has?  The Periodic TableNote how the P. Table is fundamentally arranged in terms of  increasing atomic number (Z)

61   Task: Use the P. Table to determine how many protons and electrons the  following types of atoms (elements) have: Atom  #p  #e  Atom  #p  #e  Atom  #p  #e Carbon (C):

Silicon (Si):

Lead (Pb):

Discussion: The atomic number (Z) indicates the number of protons an  individual atom has. What other type of subatomic particle is also found  within an atom’s nucleus? How is the number of these particles within any  nucleus represented and/or determined?

62 COMPLETE ATOMIC SYMBOL: Example: Write the complete atomic symbol for an atom of Carbon that  contains 6 protons and 6 neutrons.  ‘Shorthand’ version of the complete atomic symbol: Task: Carbon–14 has a mass number of 14. Use this  information to write its complete atomic symbol. Do the same for U-235 and Cl-35. * remind me to tell a story about U-235 and U-23463 Understanding Isotopes An element has a FIXED number of protons in its nucleus. (This information is contained within the element’s Atomic  Number. E.g. All hydrogen (H) atoms have 1 proton in their nuclei,  while all carbon (C) atoms have 6 protons in their nuclei).   HOWEVER, an element can have a VARIABLE number of  neutrons in its nuclei.    (This does NOT alter the identity of the element (#p same), but  DOES make the element heavier or lighter (# n changed))

The AVERAGE atomic mass value for ALL an element’s isotopes  is displayed in the periodic table. E.g. Chlorine has a mass number of 35.45 amu* – there are NO single chlorine atoms in existence with a mass of 35.45 amu (i.e. no such  thing as 0.45 of a neutron!), but there are Cl isotopes with mass  numbers of 35 and 37 – their weighted average is 35.45 amu

The complete atomic symbol’s mass number’ (A) and the  respective Element’s ‘box weight’ in the periodic table do NOT convey the same information.   The complete atomic symbol denotes the mass of ONE isotope of the  element in amu, while the p. table gives is the average mass of ALL isotopes of the element in amu.

*Note: an amu is an atomic mass unit – the mass of a single proton or  neutron. This is ≈1.66053873 x 10-24g. It is much simpler to count atomic masses in amu – “an atom of carbon -12  (which contains 6 p and 6 n, so has a mass number of 12) weighs 12 amu” is  better than saying “an atom of carbon -12 weigh 1.992648 x 10-23 grams”!64 Task: Complete the following table for the isotopes of Carbon. (Tip: what  are the values of #p and #e ALWAYS for carbon? Where would you find  this information?) Complete atomic Symbol #p  #e  #n





I have a very poor 14C joke; ask at your own peril….  Determining Relative isotopic abundance Typical Question: Naturally occurring magnesium has the following isotopic  abundances: Isotope Abundance Mass (amu) 24Mg 78.99 23.98504 25Mg 10.00 24.98584 26Mg 11.01 25.98259 What is the average atomic mass of Mg?65 The isotopic abundances of any series of isotopes always add up  to 100%  The sum of the weighted abundances is equal to the average  atomic mass (as found in the P. Table).   Determine each isotope’s weighted abundance by multiplying its  FRACTIONAL ABUNDANCE by its isotopic mass

Edit the table supplied to make two new columns - fractional  abundance and weighted abundance.    Determine the weighted abundances and then combine them to  find the element’s average atomic mass.

Answer: Isotope Abundance Fractional  Abundance Mass  (amu) Weighted  Abundance (amu) 24Mg 78.99 0.7899 x 23.98504 = 25Mg 10.00 0.1000 x 24.98584 = 26Mg 11.01 0.1101 x 25.98259 = Sum of weighted abundances = _________________      Check: Is the sum of weighted abundances equal to the average atomic mass  for Mg from the P. Table?66 Task (complete outside of class): As the above table illustrates, the amu  masses of individual atoms (isotopes), or even molecules, can be measured  with a high degree of precision. This is made possible through the technique  of mass spectroscopy. Make notes on the design and function of a mass  spectrometer, as well as how the mass spectrum of Chlorine may be used to  determine its average atomic mass (p 68, 69).67   “Symbol” The following question were taken from your 1st practice  midterm:

Write the complete atomic symbol for the isotope that contains 29 protons and 34  neutrons.  Complete atomic symbol:  Answer:63 Cu 2968 Introduction to Molecules and Ions Reading: Ch 3 sections 1- 6 Homework: Chapter 3: 23, 25, 27, 29*, 31, 33*,  35*, 39, 43, 49*, 51*, 53

* = ‘important’ homework question   Molecules, Molecular Elements and Molecular Compounds Recap: What is a molecule? What is a molecular compound? What is a  molecular element? Molecule: Molecular Element: Molecular Compound:

Molecules and their Chemical FormulasThere are two ways of describing the components (i.e. the number  and type of atoms) found inside any molecule:    Molecular Formula: the actual number and type of atoms in a  compound, e.g. hydrogen peroxide = H2O2 Empirical Formula: the lowest whole number ratio of each type of  atom in a compound e.g. hydrogen peroxide = HO

69 Task: Complete the following table Name Molecular formula Empirical formula Hydrogen peroxide

Dinitrogen tetroxide N2O4

Benzene C6H6

Butane C4H10

Tetraphosphorus  decoxide   P4O10

Note: Empirical formulas most often pertain to molecular / covalent compounds, as ionic compounds’  formulas are typically in their lowest ratio to begin with (this will be discussed further below) Picturing Molecules – Structural Formulas   A structural formula is simply a more detailed version of the  molecule’s corresponding molecular formula.   The major difference is that structural formulas also indicate  the spatial relationship, and bonding, between atoms in a  molecule

Eg: 70 Name and  Molecular  Formula Electron density map of the  ‘real’ molecule Structural Formula Water (H2O)

Task: Using the electron density maps as a guide, complete the following  table Key   Name and  Molecular  Formula Electron density map of the  ‘real’ molecule Structural Formula Carbon  Dioxide (CO2)

Methane  (CH4)

Ethane  (C2H6)

Ammonia (NH3)

71 Naming Molecular Elements and Compounds Task: Write the formula of and name as many molecular  elements and compounds as you can Formula Name Formula Name

Discussion: What relationships do you see between the names and formulas  of molecular compounds?

72 Prefix Table  Number of atoms Prefix* Example 1






*Prefixes are dropped for the first single atom in a formula. E.g. CO2 is named ‘Carbon  dioxide’, not ‘Mono Carbon dioxide’. Tasks: 

Name the Following: NF3



Write formulas for the following: Chlorine dioxide

Chlorine pentafluoride

Dihydrogen monosulfide*

* If named using ionic nomenclature, also known as _________________73 Location, Location, Location! ONLY a non-metal bonded to another non-metal (top RHS p.  table) make molecular materials with covalent bonds. E.g. CO,  H2O, SO3 THESE MATERIALS ARE NAMED IN ACCORDANCE WITH THE ABOVE ‘MOLECULAR’ RULES  

Metallic vs Non metallic Elements in the Periodic TableONLY a non-metal (top RHS) bonded to metal (LHS) make giant  compounds with ionic bonds. E.g. NaCl, CaO   THESE MATERIALS ARE NAMED IN ACCORDANCE WITH THE ‘IONIC’ RULES DISCUSSED  IMMEDIATELY BELOW

74 Ions and Ionic Compounds Questions: What are ions? How are they made? Ion:

*Atomic Ions:

* Ask me to tell you a very poor ion joke….. Atomic (micro) scale diagram of Ionization and macro scale crystal  growth (slide)In reality, electron(s) are EXCHANGED between atoms in order to  become ionic compounds. I.E. what is lost by the metal (to become an  Mn+ cation) is gained by the non-metal (to become An anion)

75 Making and Naming Ionic Formulas List of Common atomic ions (must learn): See appendixes Group I Group VII Group II Group VI Group III Group V

Naming atomic ions: An atomic (+ve) cation has the same  name as the metal it was formed from. An atomic (-ve) anion  has the same root as the non-metal it was formed from, but and  –ide ending. Examples: Metal atom Metal cation Non-metal atom Non-metal anion Na




Ionic formulas are made by combining ANY cation (+ve) with any  anion (-ve).  The order in ANY ionic formula is cation first, anion second, in both formula and name. i.e. (cation)(anion)   Examples: NaCl (sodium chloride)    LiF ( )

76 Ionic formulas ALWAYS have a ZERO net charge – i.e. the (+) and  (-) ionic charges in ANY formula cancel. If the above rule is followed, the ionic compound must exist and is  probably sitting on a shelf in the chemistry stock room!

Task: Construct and name as many ionic compounds as possible from the  following ions:  Li+ Ca2+ Al3+ Cl- O2- N3- List of Common molecular ions (must learn): See attached handout. Trick – many molecular ions appear on the data sheet (see  handout). Just keep using (homework) and/or looking  (fridge) at the rest

Naming molecular ions:  There is ONLY one molecular cation – (NH4)+, ammonium. Molecular anions with NO (or very few*) oxygen atoms in their  structure have the –ide ending. Examples: -OH (hydroxide)*, CN- (cyanide) Molecular anions with ‘lots’ of oxygen atoms in their  structure have the –ate ending. Examples: (SO4)2-(sulfate),  (NO3)-(nitrate), (CO3)2-(carbonate), (PO4)3-(phosphate)77 Recall: Ionic formulas ALWAYS have a ZERO net charge – i.e.  the ionic charges in ANY formula cancel. This is true for molecular ions too – just treat the whole molecular ion  as if it were an atomic ion when making the formula. Name the  resulting compound in a similar way also.

Task: Construct and name as many ionic compounds as possible from the  following ions:  Li+ Mg2+ (NH4)+(NO3)-(SO4)2-(PO4)3- Naming Ionic compounds containing a cation of variable charge Metallic elements from the center of the periodic table (the transition  series, between groups II and III) can form atomic ions with a range  of +ve charges. Examples: Fe2+ and Fe3+, Cu+and Cu2+.

Question: Can you see a potential problem with regard to writing the names  and formulas of ionic compounds containing such cations? Answer:78 Ionic formulas featuring a variable charge (oxidation state) cation  include the charge of the cation (written in Roman numerals) in the  formula name. E.g.: Cu2O = Copper(I) oxide

Task: Complete the following table: Name Formula Name Formula Iron (II) Sulfate

Copper (I) Phosphate



Acids and bases Discussion: Are acids and bases typically ionic or molecular compounds  (trick question!)? What is ‘special’ about them and their formulas?

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