Chapter 1 Chemistry

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Chapter 1Chemistry Measurement

• Access all online homework (HW) assignments at
  www.saplinglearning.com
• 1. Graded HW Practice Assignment(10pts) Math
  Review(100pts) due before 9/24, 11pm
• 2. Graded HW 1 (chap 1, 100 pts) due before 9/30,
  11 pm
• 3. End-of-Chapter problems Not Graded. See
  the next slide.
• 4. 2 Internet Assignments Not Graded. See
  Blackboard course documents
• 5. Exam 1 over chapters 12 in a little over two
  weeks 22 of Grade
• 6. Will finish chapter 1 next week.

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Suggested Chapter 1 End-of-Chapter Homework

• - Text Chapter 1 Pages 31 to 40 these are not
  to be turned in. If possible, finish them before
  I end the lectures on chapters 12. I will
  devote class time to answering questions on
  these.
• 6 7 9 12 13 14 18 22
  24 28 34 44 51 53 55 57
  61 68 69 73 75 79 81
  83 85 89 91
• 97 103 109 119 129 137 147

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I. Introduction

• A. Definitions
• Matter - Anything which occupies space and has
  mass.
• Chemistry - The study of the composition, the
  structure, the properties, and the reactions of
  matter.
• Chemistry The Study of Stuff.
• B. Things Chemists Do
• 1. Theoretical Models, Mathematics,
  Instruments
• 2. Design Chemical plants, Equipment,
  Instruments
• 3. Publish, Mentor, Communicate - Pass on
  chemical information

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I. Introduction B. Things Do Continued

• 4. Synthesis Chemical Companies, Research,
  Pharmaceutical Industries, Petroleum Industries,
  etc
• 5. Purify / Separate Needed for most chemical
  tasks Classical Instrumental
• 6. Qualitative Analysis (What is Present)
  Research, Product Control, Regulatory Agencies,
  Forensics Classical Instrumental Methods
• 7. Quantitative Analysis (How Much is Present) -
  Research, Regulatory Agencies, Forensics, Product
  Control Classical Instrumental Methods

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I. Introduction C. Chemical Interfaces
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I. Introduction Continued

• D. Divisions of Chemistry
• 1. Inorganic 2. Organic
• 3. Biochemical 4. Physical
• 5. Analytical 6. Nuclear

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II. Scientific Method (Different than our Text
use mine)

• 1. Define Area of Study.
• 2. Perform Background Research - Literature
  Search (CA, BA, PM).
• 3. Make a Hypothesis (a proposal about the
  course of action).
• 4. Conduct Experiments Collect Data Make
  Observations - to Test Hypothesis.
• 5. Propose a Theory (Explanation that best fits
  data to date).
• 6. Theory may become law after decades of
  support.
• Note 1) Scientists may be biased. 2)
  Scientists frequently use models verbal or
  graphic descriptions of a concept. 3) Laws are a
  rarity. 4) Do the two Blackboard assignments a)
  search the scientific literature for references
  to Tylenol and Liver, and b) tour the
  webelements site.

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III. Matter A. Categories
2 or More Substances (Not Pure)
1 Substance (Pure)
Cu or O2
H2O
Salt in H2O (Solution)
Sand in H2O (Insoluble)
Examples
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III. Matter B. States Three States of Matter

• Gas Has no definite volume and no definite
  shape is compressible completely fills its
  container. Example Water at 110 oC under normal
  pressure.
• Liquid Has definite volume and no definite
  shape is incompressible conforms to the shape
  of its container. Example Water at 20 oC.
• Solid Has definite volume and definite shape is
  incompressible does not change shape or volume
  from one container to another. Example Water at
  - 50 oC.
• - Can change a substance from a solid to a gas by
  changing temperature or pressure. Upon heating,
  the particles which make up that substance pick
  up energy move faster. As the particles move
  faster they tend to break attractive forces
  between the particles become somewhat
  independent.

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III. Matter B. Changes C. Separation

• B. Changes
• 1) Physical Change A change in the form of
  matter but not in its chemical identity.
  Examples melting, boiling
• 2) Chemical Change A change in which matter is
  converted to a different kind of matter.
  Examples food digestion paper burning
• C. Separation
• - Chemists frequently need to separate mixtures
  into pure substances.
• - Heterogeneous mixtures can sometimes be
  separated by simple methods involving filtration
  or density Sand Water Gold Dirt
• - Homogeneous mixtures are more difficult to
  separate. Some can be separated by distillation
  or by solubility. Usually we use chromatographic
  methods to separate homogeneous mixtures.

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III. Matter C. Separation of
Homogeneous Mixtures

• Distillation can be used for separation of a
  homogeneous mixture of low boiling liquids such
  as acetone (bp 56 oC) and water (bp 100 oC) - as
  in fingernail polish remover.

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III. Matter C. Separation of homogeneous mixtures

• - Chromatographic methods are powerful separation
  methods.
• - Chromatography the selective distribution of
  components in a mixture between two phases
  Stationary (SP) and Mobile (MP).
• - Compounds X and O have different affinities for
  the stationary phase inside the column and the
  mobile phase which flows through the column. If
  O has a higher affinity for the stationary phase,
  then it will stay on the column longer and
  eventually be separated from X.

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C. Separation - Column Chromatography (CC)

• Column Chromatography (CC) for Rough Separation
  of Components in a Homogeneous Mixture

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III. Matter C. Separation Thin Layer
Chromatography (TLC)

• TLC can be used for Rough Separation of
  Components of Homogeneous Mixture such as Ink.
  Stationary phase solid coated on the glass
  plate Mobile phase solvent creeping up the
  plate.

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III. Matter C. Separation Gas Chromatography
(GC)
1 µL
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III. Matter C. Separation Liquid
Chromatography (LC)
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C. Separation Summary

• CC TLC Used for qualitative analysis and
  rough separation of the components of a
  homogeneous mixture.
• GC Used for qualitative and quantitative
  analysis of complex mixtures of somewhat volatile
  chemicals in a solution. Can automate.
• LC Used for qualitative and quantitative
  analysis of complex mixtures of nonvolatile
  chemicals in a solution. Can automate.
• Output Chromatogram plot of Intensity/Signal
  from detector versus time. Qualitative data from
  time detector signal and Quantitative data from
  detector signal intensity.
• Detector The detector produces a signal
  proportional to the amount of chemical coming off
  of column. A good detector will also aid us in
  determining the identity of the separated
  chemicals coming off of the column. Example
  Mass Spectrometer (MS).

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C. Separation - Chromatogram

• LC chromatogram of ephedrine alkaloids. A plot of
  UV absorbance versus time. Retention time is
  characteristic of compound area under a peak is
  proportional to the amount of compound. EPH
  ephedrine PSE pseudoephedrine.

Abs _at_ 255 nm
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IV. Measurements

• A. Introduction
• - Much of the development of modern chemistry
  was based on well designed experiments.
• - In these experiments measurements were made of
  mass, volume, length, time, temperature.
• - We have to know how to make, interconvert,
  manipulate, estimate, correctly report
  (significant figures - SF) judge measurements
  (accuracy and precision).
• - We will use both fixed decimal and scientific
  notation.

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IV. Measurements B. Accuracy Precision

• 1. Accuracy
• - In an experiment many measurements are made on
  the same system, and an average value, X, is
  calculated.
• - The closeness of the average value to the true
  value is a measure of accuracy.
• 2. Precision
• - The grouping of similar measurements is the
  precision.
• - We represent precision with Standard Deviation
  (s).
• - X 1 s range in which 68 of all future
  measurements should fall if used same measuring
  system and techniques.
• s ? (X - xi)2
• n-1
• - Given measurements of 5.1, 5.2 5.0, calculate
  X s.
• Answer X 5.1 0.10 (convention
  use 2 sf for std dev)

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IV. Measurements Accuracy Precision
Reasonable
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IV. Measurements C. Significant Figures (SF)

• 1. Definition All digits known with certainty
  plus the first uncertain one. Note you assume
  that the last digit is in doubt by 1
• 2. Examples 162 3 SF 92 2 SF 9.634 x 107
  4 SF
• 3. Zeros (pay careful attention)
• a) Are significant when between other digits
  1023 4 SF
• b) Are significant when to right of decimal at
  end
• 12.0 3 SF 10.0100 6 SF
• c) Not significant when left of first digit
  000.011 2 SF
• d) Dont know in following case 100
• - Can use a decimal, underline, or exponential
  notation
• 100. or 1.00x102 or 100 3 SF

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IV. Measurements C. Significant Figures (SF)
Continued

• - Why Use To not mislead others using the
  numbers.
• - Examples
• 1. Silver Ore from Mexico
• 2. Bookcase Finger Method gave width as 153
  inches
• Room width was actually 143.5 inches!
• - Note SF in experiment limited by
• 1) Technique Used
• 2) Measuring Devices
• - Use
• 1. Addition Subtraction Answer limited to
  first doubtful digit can lose SF on subtraction
  gain SF on addition.
• 2. Multiplication Division Answer limited
  to Smallest of SF.
• - When have multiple steps, then do addn/subtn
  first.

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IV. Measurements C. Significant Figures -
Examples

• - Addition and Subtraction (subtract 9.31 from
  9.3331)
• 9 . 3 3 3 1
• - 9 . 3 1
• 0 . 0 2 3 1 0.02 (1 SF)
  Cut Off at First Doubtful Digit
• - Multiplication Division
• 1.011 x 7.01 / 6.9 1.0 (2SF)
  Fewest SF limits answer
• - Note
• 2.001 x 3.099 / (8.999 - 8.998) 6.201099x103
  6x103 (1 SF) Why?

2SF
4SF
3SF
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V. Problem Solving - Two Basic Methods

• 1) Formula - Solve a given formula for the
  variable of interest, plug in along with the
  units and solve. Check units estimate answer.
• Example How many grams would 24.5 mL of mercury
  weigh if the density of mercury is 13.6 g/mL?
  You are given d mass/mL
• d m/mL m d x mL 13.6 g x 24.5
  mL 333 g
• 
  mL
• 2) Conversion Factor (CF) - Write down initial
  data multiply by CFs such that units cancel
  until get to completion. Check units estimate
  answer. CFs have to be given or memorized to
  enough sf.
• Example How many yards are in 0.0031 miles?
• CF 5280. ft 1.000 mi 3.00 ft
  1.00 Yd
• 0.0031 mi x 5280. ft x 1.00 yd
  5.4 yd
• 1.000 mi
  3.00 ft

3 SF, g
2 SF, yd
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V. Problem Solving Examples

• Note Do not let your conversion factor limit
  the of SF
• 1) Convert 2.0x109 ton/yr to ton/hr 1.0yr365d
  1.0d 24hr
• 2.0x109 ton x 1.0 yr x 1.0 d
  2.3x105 ton
• yr 365 d 24
  hr hr
• 2.0x109 ton 1.0 yr 1.0 d
  2.3x105 ton
• yr 365 d 24 hr
  hr
• - The above is a clear, neat, easy to follow
  format please use.

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V. Problem Solving Examples Continued

• 2) Convert 6.0 Ft2 to in2
  1.0 Ft 12 in
• 6.0 ft2 x 12 in x 12 in
  8.6 x 102 in2
• 1 ft
  1 ft
• 6.0 ft2 12 in 12 in
  8.6 x 102 in2
• 1 ft 1 ft
• Common mistakes
• 6.0 ft2 x 12 in / 1 ft 72 (ft x
  in)
• 6.0 ft2 x 12 in2 / 1 ft2 72 (in2)
• 36 ft2 x 12 in / 1 ft x 12 in / 1 ft
  5.2x103 (in2)
• 6.0 ft2 x 1 ft / 12 in x 1 ft / 12 in
  0.042 (ft4 / in2)

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VI. Units - A. SI English

• Note 1) Following CF are exact numbers
• 2) Need to memorize these
• English - Random, difficult system.
• 16 oz 1 lb 3 ft 1 yd
  12 in 1 ft
• 5280 ft 1 mi 1 qt 4 cups
  2 cups 1 pt
• 4 qts 1 gal 2 pts 1 qt
  1 ton 2000 lbs
• SI - System used by most of the world. Each is
  10x times whatever follows it 1 centameter (cm)
  10-2 m. Know these.
• Prefixes kilo k 103 1km 103 m
• centa c 10-2 1cm 10-2 m
• milli m 10-3 1mm 10-3 m
• micro µ 10-6 1µm 10-6 m
• nano n 10-9 1nm 10-9 m
• pico p 10-12 1pm 10-12 m

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VI. Units - A. SI English

• Conversion Factors for SI Prefixes (using grams ,
  g)
• 1 kg 1000 g
• 1x102 cg 1 g or 1 cg 10-2 g
• 1x103 mg 1g or 1 mg 10-3 g
• 1x106 µg 1g or 1 µg 10-6 g
• 1x109 ng 1g or 1 ng 10-9 g
• 1x1012pg 1g or 1 pg 10-12 g
• 1) Know k, c, m, µ, n, p prefixes
• 2) Many toxic chemicals are measured in µg to pg
  amounts.
• 3) We frequently need to measure small distances
  like the wavelength of ultraviolet light ( 200
  nm).

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VI. Units - A. SI English
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VI. Units - B. Length, Volume, Mass

• - Mass gram g
• 454 g 1.00 lb (English-SI Bridge
  Memorize)
• -Volume liter L
• 1.00 L 1.06 qt (English-SI Bridge
  Memorize)
• - Length meter m
• 1.00 inch 2.54 cm (English-SI Bridge
  Memorize)
• Note 1) 1.00 g water 1.00 mL _at_ 4oC Memorize
• 2) 1.00 mL 1.00 cm3 for all substances -
  Memorize

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VI. Units - B. Length, Volume, Mass1000 mL
1000 cm3 1 L
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VI. Units - Volume Measuring Devices
3 - 4 SF
Auto Pipet, Syringe 3 SF
Volumetric Pipet 4 SF
Volumetric Flask 4 SF
Mohr Pipet 3 SF
2 - 3 SF
1 SF
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VI. Units - C. Conversions

• - Example How many inches high is a 1.75 m
  fence?
• 1.75 m x 100. cm x 1.00 in 68.9
  in
• 1.00 m 2.54 cm
• - Example How many g are in 1.00 oz?
• 1.00 oz x 1.00 lb x 454 g
  28.4 g
• 16.0 oz 1.00 lb
• - Example How many L are in 2.4x102 cups?
• 2.4x102 cups x 1.0 qt x 1.00 L
  57 L
• 4.0 cups 1.06 qt
• D. Density
• - Conversion Factor for converting between mass
  volume
• - Characteristic physical property that can be
  used to identify a substance.
• Examples dwater 1.00 g/mL dbenzene 0.80
  g/mL dgold 19.3 g/cm3 dAl
  2.70 g/cm3
• Note cm3 generally used for solids mL for
  liquids (both same).

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VI. Units - D. Density Continued

• Calculations
• Can do two ways Use the formula with algebraic
  manipulation or use the Conversion Factor
  method Ill use the formula method.
• - Examples
• 1) Your fiancé gives you a yellow ring that has
  a volume of 2.9 cm3 and a mass of 7.8 g. Is it
  Gold or Aluminum painted yellow?
• d m/v 7.8 g / 2.9 cm3 2.7 g/cm3
  (It is Al)
• 2) How many g are in 200. mL of oil? Given d of
  oil 0.92g/mL
• d m/V
• m d x V 0.92 g x 200. mL 1.8
  x 102 g
• mL

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VI. Units - E. Temperature
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VI. Units - E. Temperature Continued

• Scale oF oC K
• BP H2O 212 100. 373
• MP H2O 32.0 0.00 273
• Lowest -460 -273 0.00
• - Need to know following formulas
• o F (1.80 x oC) 32.0
• K oC 273
• Examples
• - Convert 31 oC to oF oF (1.80 x 31)
  32.0 88 oF
• - Convert 25 oC to K K 25 273 298 K