Keys to the Study of Chemistry

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Keys to the Study of Chemistry

• Chapter 1

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Chapter 1 Keys to the Study of Chemistry
1.1 Some Fundamental Definitions 1.2 Chemical
Arts and the Origins of Modern Chemistry 1.3 The
Scientific Approach Developing a Model 1.4
Chemical Problem Solving 1.5 Measurement in
Scientific Study 1.6 Uncertainty in Measurement
Significant Figures
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CHEMISTRY
Is the study of matter, its properties, the
changes that matter undergoes, and the
energy associated with these changes.
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Definitions

• Matter - anything that has mass and volume
• Properties of matter
• Physical properties - describes the substance as
  is. ex. color, odor, density, mp, bp,
  hardness, conductivity,...
• those which the substance shows by itself without
  interacting with another substance such as color,
  melting point, boiling point, density
• Chemical properties - describes how the substance
  changes or reacts to form other substances ex.
  Iron will react with oxygen to form rust
• those which the substance shows as it interacts
  with, or transforms into, other substances such
  as flammability, corrosiveness

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Definitions

• Physical change - change of physical appearance
  not composition
• results in change of physical properties
• ex Ice ? Water ? Water vapor
• it's still H2O
• includes changes of state and mechanical changes
  (breaking, bending)
• Chemical change - a reaction - alters the
  identity of the substance
• ex. electric current through water produces
  hydrogen and oxygen
• Ex. iron rusting - CC
• dry ice sublimes - PC

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Sample Problem 1.1
Distinguishing Between Physical and Chemical
Change
(a) Frost forms as the temperature drops on a
humid winter night.
(b) A cornstalk grows from a seed that is watered
and fertilized.
(c) Dynamite explodes to form a mixture of gases.
(d) Perspiration evaporates when you relax after
jogging.
(e) A silver fork tarnishes in air.
PLAN
Does the substance change composition or just
change form?
SOLUTION
(a) physical change
(b) chemical change
(c) chemical change
(d) physical change
(e) chemical change
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Definitions

• intensive properties - do not depend on the
  amount of the substance
• can be used to identify a substance
• ex. mp, density, color
• extensive properties - depend on amount
• ex mass, volume

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States of Matter

• How many are there?
• Matter can exist in four possible states
• Solid
• Liquid
• Gas
• Plasma the sun, lightning, lab experiments,
  star trek, etc.

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States of Matter(in the macroscopic scale)

• Solid - definite shape and volume (not
  compressible)
• Liquid - definite volume, not shape (not
  compressible)
• Gas - no definite shape or volume (compressible,
  expandable)

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States of Matter(in the submicroscopic scale)

• Solid - particles(atoms or mol.) are packed
  closely in a definite arrangement
• Liquid - particles close together, changing
  location
• Gas - particles far apart, moving very quickly

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Figure 1.2
The Physical States of Matter
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States of Matter(in the submicroscopic scale)

• What happens to the size of an iron plate when
  heated?
• Why?
• What happens to the size of a whole in an iron
  plate when heated?
• Are you sure?

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States of Matter(in the submicroscopic scale)

• What happens to the size of an iron plate when
  heated?
• Why?
• What happens to the size of a whole in an iron
  plate when heated?
• Are you sure?

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States of Matter

• State of matter depends on
• T
• P
• nature of the substance
• (intermolecular forces determine how strongly
  particles hold together, not changeable)
• T and P can be changed to alter phase

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Energy in the Study of Matter

• Energy - the ability to do work
• Kinetic Energy The energy of motion.
• 1/2mv2
• a rolling car
• a vibrating atom (thermal energy)
• Potential Energy Stored energy due to position.
• car on a hill
• an unstable chemical structure (chemical energy)
  due to relative positions, attractions, and
  repulsions of its particles.

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Energy in the Study of Matter

• In nature situations of lower potential energy
  are typically more stable.
• Cars tend to roll down hills converting potential
  energy into kinetic energy.
• Chemical reactions tend to roll down hills too,
  converting chemical energy into thermal energy.
  Although not always.

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Energy in the Study of Matter

• Cars can be moved back to the top by adding
  energy.
• Many reactions can be reversed by adding energy.
• NOTE reactions have exactly opposite energy
  requirements when reversed.
•  
• metabolism C6H12O6 (s) 6 O2 (g) ? 6 CO2
  (g) 6 H2O (g) ?H -2540 kJ
•  
• photosynthesis 6 CO2 (g) 6 H2O (g) ? C6H12O6
  (s) 6 O2 (g) ?H 2540 kJ
• combustion of hydrogen and oxygen, and hydrolysis
  of water

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The Law of Conservation of Energy

• a.k.a. 1st Law of Thermodynamics
• - Energy is neither created nor destroyed in
  chemical reactions (only changed in form)
• Energy released or absorbed in processes always
  existed and always will exist in some form
• ex car on hill to car rolling to friction
  produced heat
• Consider 4 examples of energy conversion on page 7

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Exceptions to 1st Law?

• The Conversion of Matter to Energy -
• The Einstein equation
• E mc2
• relationship between matter and energy
• useful for nuclear reactions
• proven nearly 40 years later by the atomic bomb
• 1 g of matter to energy heat a house for 1000
  years
• less than 1 conversion in nuclear explosions
• Result The sum total of matter and energy in
  the universe is constant.

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Metric/SI System

• Well organized system
• cm-ml-g-cal
• conversion within unit types based on multiples
  of ten.
• prefixes can be used to express multiples of ten.
  
• the specific metric units used for scientific
  measurement are SI units(Systeme International)

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Table 1. 2
SI - Base Units
Physical Quantity
Unit Name
Abbreviation
mass
kg
kilogram
meter
length
m
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SI Units

• all other SI units are derived from these 7
• ex volume in SI is derived, m3, but we usually
  use L(dm3) or ml(cm3)
• mass is kg 2.2 pounds chemistry use gram

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Common Decimal Prefixes Used with SI Units
Table 1.3
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Some Interesting Quantities
Figure 1. 10
Length Volume Mass
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Exact and Inexact Numbers

• Numbers that have to be measured are always
  inexact. (how tall? 6.????....)most English to
  metric conversions are inexact
• Exact numbers come from defined values (3ft/yd)
  or integer counts of values (14 students)  
• Inexact numbers have limits, exact number dont

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Significant Figures

• Indicates the exactness of a measurement.
• In calculations inexact data yields inexact
  answers. How exact is answer?

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The Number of Significant Figures in a
Measurement Depends Upon the Measuring Device
Figure 1.14
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Rules for Determining Which Digits are
Significant
except zeros that are used only to position the
decimal point.
All digits are significant

• Make sure that the measured quantity has a
  decimal point.
• Start at the left of the number and move right
  until you reach the first nonzero digit.
• Count that digit and every digit to its right as
  significant.

Zeros that end a number and lie either after or
before the decimal point are significant thus
1.030 ml has four significant figures, and 5300.
L has four significant figures also. Numbers
such as 5300 L are assumed to only have 2
significant figures. A terminal decimal point is
often used to clarify the situation, but
scientific notation is the best!
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Rules for Significant Figures

• 1. nonzero digits always sig
• ex 45 - 2 sig figs , 1.37 - 3 sig figs
• 2. captive zeros must be significant
• ex 1001 - 4 sig figs , 1.0005 - 5 sig figs
• 3. leading zeros are not significant
• ex .004 - 1 sig digit, 0.0045 2 sig digits
• 4. trailing zeros are significant if there is a
  decimal point
• ex .00400 - 3 sig figs, 1000. - 4 sig figs
• 5. zeros at end, no decimal point ???
• ex 1000 1,2,3,or 4 sig figs?? must assume not
  significant
• use exponential notation to remedy
• ex 1000 with one sig figs is 1 x 103
• ex 1000 with four sig figs is 1.000 x 103

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Sample Problem 1.7
Determining the Number of Significant Figures
(b) 0.1044 g
(a) 0.0030 L
(c) 53.069 mL
(e) 57,600. s
(d) 0.00004715 m
(f) 0.0000007160 cm3
PLAN
Determine the number of sf by counting digits and
paying attention to the placement of zeros.
SOLUTION
2sf
4sf
5sf
(f) 7.160x10-7 cm3
(d) 4.715x10-5 m
4sf
(e) 5.7600x104 s
5sf
4sf
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Rules for Significant Figures in Answers
1. For addition and subtraction. The answer has
the same number of decimal places as there are
in the measurement with the fewest decimal
places.
Example adding two volumes
106.78 mL 106.8 mL
Example subtracting two volumes
863.0879 mL 863.1 mL
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Rules for Significant Figures in Answers
2. For multiplication and division. The number
with the least certainty limits the certainty of
the result. Therefore, the answer contains the
same number of significant figures as there are
in the measurement with the fewest significant
figures.
Multiply the following numbers
9.2 cm x 6.8 cm x 0.3744 cm
23.4225 cm3 23 cm3
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Issues Concerning Significant Figures
be sure to correlate with the problem
FIX function on some calculators
graduated cylinder lt buret pipet
60 min 1 hr
numbers with no uncertainty
1000 mg 1 g
These have as many significant digits as the
calculation requires.
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Rules for Rounding Off Numbers
1. If the digit removed is more than 5, the
preceding number increases by 1. 5.379 rounds to
5.38 if three significant figures are retained
and to 5.4 if two significant figures are
retained.
2. If the digit removed is less than 5, the
preceding number is unchanged. 0.2413 rounds to
0.241 if three significant figures are retained
and to 0.24 if two significant figures are
retained.
3.If the digit removed is 5, the preceding number
increases by 1 if it is odd and remains unchanged
if it is even. 17.75 rounds to 17.8, but 17.65
rounds to 17.6. If the 5 is followed only by
zeros, rule 3 is followed if the 5 is followed
by nonzeros, rule 1 is followed 17.6500 rounds
to 17.6, but 17.6513 rounds to 17.7
4. Be sure to carry two or more additional
significant figures through a multistep
calculation and round off only the final answer.
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Sample Problem 1.8
Significant Figures and Rounding
PLAN
In (a) we subtract before we divide for (b) we
are using an exact number.
SOLUTION
2.104 cm2
4.16 g/ cm3
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Precision and Accuracy Errors in Scientific
Measurements

• Systematic error -
• Values that are either all higher or all lower
  than the actual value.
• Random Error -
• In the absence of systematic error, some values
  that are higher and some that are lower than the
  actual value.
• Accuracy -
• Refers to how close a measurement is to the real
  value.
• - Accurate measurements have low systematic and
  random error

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Precision and Accuracy Errors in Scientific
Measurements

• Precision
• measurements that are in close agreement
• i.e. reproducible
• - measurements have more significant figures
• the sloppiness of measurement is the degree of
  precision
• precise measurements have low random error, but
  may have systematic error

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Precision and Accuracy Errors in Scientific
Measurements

• When several measurements of the same item agree
  with each other, they are precise but not
  necessarily accurate (perhaps the scale is
  wrong).
• A Thermometer marked to tenths of degrees is more
  precise than one marked with whole degrees, but
  not necessarily more accurate.
• If equipment is calibrated better precision leads
  to better accuracy.

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Figure 1.16
Precision and Accuracy in the Laboratory
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Precision and Accuracy in the Laboratory
Figure 1.16 continued
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Dimensional Analysis

• Method by which units (dimensions) are used to
  help solve problems, and check answers.

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Calculations with Dimensional Analysis

• 1. Use equivalence statement to get conversion
  factor.
• 2. Pick conversion factor that cancels
  appropriate unit.
• 3. Multiply quantity by conversion factor.
• 4. Check Sig Figs.
• 5. Ask whether your answer makes sense.

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Relationships to Know

• English unit relationships
• Metric prefixes
• 1cm3 1mL
• 2.54 cm 1 in.
• 454 g 1 lb.
• 1L 1.057 qt.

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Example Conversions

• Convert
• .3704 m to cm
• 5.6 cm to inches
• .0478 mg to ?g
• 351 in3 to cm3, and L
• 200 ml to fluid ounces
• 95 kmph to ft/s

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Example Problems

• Solve
• A faucet dripping 53 drops every minute adds how
  many gallons to your monthly water bill? It
  takes approximately 19 drops to equal 1 ml.
• How much does it cost a month to run a 200W yard
  light that only runs about 10 hours every night?
  The cost of electricity is .06048/KWHr.
• A water solution contains 12 NaOH by mass, and
  has a density of 1.131g/ml. What volume of this
  solution contains 3500 g NaOH?

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Sample Problem 1.2
Converting Units of Length
PLAN
Known - length (in cm) of wire and cost per
length (in ft) We have to convert cm to inches
and inches to ft followed by finding the cost for
the length in ft.
SOLUTION
length (cm) of wire
Length (in) length (cm) x conversion factor
2.54 cm 1 in
325 cm x
128 in
length (in) of wire
12 in 1 ft
Length (ft) length (in) x conversion factor
length (ft) of wire
128 in x
10.7 ft
1 ft 0.15
Price () length (ft) x conversion factor
Price () of wire
10.7 ft x
1.60
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Sample Problem 1.4
Converting Units of Mass
PLAN
The sequence of steps may vary but essentially
you have to find the length of the entire cable
and convert it to mass.
SOLUTION
length (km) of fiber
8.84 x 106m
length (m) of fiber
1.05 x 104lb
mass (lb) of fiber
mass (kg) of cable
mass (lb) of cable
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Sample Problem 1.3
Determining the Volume of a Solid by Displacement
of Water
PLAN
The volume of galena is equal to the change in
the water volume before and after submerging the
solid.
SOLUTION
volume (mL) before and after addition
(24.5 - 19.9)mL volume of galena
volume (mL) of galena
4.6 mL x
4.6 cm3
volume (cm3) of galena
volume (L) of galena
4.6 mL x
4.6x10-3 L
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Density

• Which is heavier a pound of feathers or a pound
  of lead?
• What is density?
• Mass per unit volume
• Density mass/volume
• Intensive/characteristic property
• Solids g/cm3, liquids g/ml, gases g/L
• An object will float in/on another if it is less
  dense.
• Q A student took an unknown liquid and
  discovered that a volume of 9.02 mL had a mass of
  8.31g. Calculate the density.
• A d m/v 8.31g/9.02mL 0.92 g/mL

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Densities of Some Common Substances
Table 1.5
Substance Physical State
Density (g/cm3)
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Specific Gravity

• S. G. (Density of Substance)/(Density of Water)
  
• Density of water 1.00 g/mL at 4o C
• Note Specific gravity has no units.

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Density

• What is the mass of 2.5 L of ethanol?
• A d m/v m d x v
• (0.789g/mL)x(2.5L)x(1000mL/L)
• (0.789g/mL)x(2.5L)x(1000mL/L)
• 1972.5 g
• 2.0 x 103 g

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Density

• Q Calculate the volume of a 100.g sample of a
  substance which has a density of 0.92 g/mL.
• A d m/v
• v m/d 100.g/0.92g/mL

110 mL
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Sample Problem 1.5
Calculating Density from Mass and Length
PLAN
Density is expressed in g/cm3 so we have to the
mass in grams and the volume in cm3.
SOLUTION
lengths (mm) of sides
1.49g
1.49x103mg x
mass (mg) of Li
lengths (cm) of sides
20.9mm x
2.09cm
Similarly the other sides will be 1.11cm and
1.20cm, respectively.
mass (g) of Li
volume (cm3)
2.09 x 1.11 x 1.20 2.76cm3
density (g/cm3) of Li
density of Li
0.540g/cm3
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Temperature Measurements

• Common Units of Temperature
• Fahrenheit (oF)
• Celsius (oC)
• Kelvin (K)
• Boiling Point of Water
• 212oF , 100oC, 373.15 K
• Freezing Point of Water
• 32oF, 0oC, 273.15 K

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Kelvin Temperature

• Based on Celsius, but with a zero that is zero.
• No negatives.
• No degree used.
• O K is the lowest possible temperature, or
  absolute zero.

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A Comparison of Temperature Scales
Insert Figure 3.21
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Conversion Among Temperature Units

• Because the temperature scales have different
  zero points, formulas must be used to carry out
  the conversions.
• K oC 273.15
• oC K - 273.15
• oC

(oF - 32)
5
(oF - 32)
or
9
1.8
9
(oC)
32
oF
5
or
1.8(oC) 32
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Example Temperature Conversions

• Convert 350oF to oC and K.
• oC (350-32 )(5/9) (318)(5/9) 177oC
• K 177 273 450 K
• Convert -40oC to oF
• oF (9/5)(-40) 32 9x(-8) 32
• -7232 -40o F
• Convert 298 K to oC
• oC 298 - 273 25o C
  

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The Freezing and Boiling Points of Water
Figure 1.12