ANALYTICAL CHEMISTRY CHEM 3811 CHAPTERS 1 AND 3 (REVIEW)

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ANALYTICAL CHEMISTRY CHEM 3811 CHAPTERS 1 AND 3
(REVIEW)
DR. AUGUSTINE OFORI AGYEMAN Assistant professor
of chemistry Department of natural
sciences Clayton state university
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CHAPTERS 1 AND 3 MEASUREMENTS, SIGNIFICANT
FIGURES ERRORS, STOICHIOMETRY, CONCENTRATIONS
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ANALYTICAL CHEMISTRY
- Deals with the separation, identification,
quantification, and statistical treatment of
the components of matter Two Areas of
Analytical Chemistry Qualitative Analysis -
Deals with the identification of materials in a
given sample (establishes the presence of a
given substance)
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ANALYTICAL CHEMISTRY
- Deals with the separation, identification,
quantification, and statistical treatment of
the components of matter Quantitative
Analysis - Deals with the quantity (amount) of
material (establishes the amount of a substance
in a sample) - Some analytical methods offer
both types of information (GC/MS)
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ANALYTICAL CHEMISTRY
Analytical Methods - Gravimetry (based on
weight) - Titrimetry (based on volume) -
Electrochemical (measurement of potential,
current, charge, etc) - Spectral (the use of
electromagnetic radiation) - Chromatography
(separation of materials) - Chemometrics
(statistical treatment of data)
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ANALYTICAL CHEMISTRY
General Steps in Chemical Analysis - Formulating
the question (to be answered through chemical
measurements) - Selecting techniques (find
appropriate analytical procedures) -
Sampling (select representative material to be
analyzed) - Sample preparation (convert
representative material into a suitable form for
analysis)
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ANALYTICAL CHEMISTRY
General Steps in Chemical Analysis -
Analysis (measure the concentration of analyte in
several identical portions) (multiple samples
identically prepared from another
source) (replicate samples splits of sample from
the same source) - Reporting and
interpretation (provide a complete report of
results) - Conclusion (draw conclusions that are
consistent with data from results)
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MEASUREMENT
Measurement - Is the determination of the
dimensions, capacity, quantity, or extent of
something - Is a quantitative observation and
consists of two parts a number and a scale
(called a unit) Examples mass, volume,
temperature, pressure, length, height, time
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MEASUREMENT SYSTEMS
Two measurement systems English System of
Units (commercial measurements) pound, quart,
inch, foot, gallon Metric System of Units
(scientific measurements) SI units (Systeme
International dUnites) liter, meter, gram More
convenient to use
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FUNDAMENTAL SI UNITS
Physical Quantity Mass Length Time Temperature Am
ount of substance Electric current Luminous
intensity
Name of Unit Kilogram Meter Second Kelvin Mole Am
pere Candela
Abbreviation kg m s (sec) K mol A cd
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DERIVED SI UNITS
Physical Quantity Force Pressure Energy Power Fr
equency
Name of Unit Newton Pascal Joule Watt Hertz
Abbreviation N (m-kg/s2) Pa (N/m2 kg/(m-s2) J
(N-m m2-kg/s2) W (J/s m2-kg/s3) Hz (1/s)
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METRIC UNITS
Prefix Giga Mega Kilo Deci Centi Milli Micro Nano
Pico Femto
Abbreviation G M k d c m µ n p f
Notation 109 106 103 10-1 10-2 10-3 10-6 10-9 10-
12 10-15
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UNIT CONVERSIONS
Length/Distance 2.54 cm 1.00 in. 12 in. 1
ft 1 yd 3 ft 1 m 39.4 in. 1 km 0.621 mile 1
km 1000 m
Time 1 min 60 sec 1 hour 60 min 24 hours 1
day 7 days 1 week
Volume 1 gal 4 qt 1 qt 0.946 L 1 L .0265
gal 1 mL 0.034 fl. oz.
Mass 1 Ib 454 g 1 Ib 16 oz 1 kg 2.20 Ib 1
oz 28.3 g
24 hours 1 day
or

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UNIT CONVERSIONS
Convert 34.5 mg to g
How many gallons of juice are there in 20 liters
of the juice?
Convert 4.0 gallons to quarts
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SIGNIFICANT FIGURES
Exact Numbers - Values with no uncertainties -
There are no uncertainties when counting objects
or people (24 students, 4 chairs, 10 pencils) -
There are no uncertainties in simple
fractions (1/4, 1/7, 4/7, 4/5) Inexact Numbers
- Associated with uncertainties - Measurement
has uncertainties (errors) associated with it -
It is impossible to make exact measurements
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SIGNIFICANT FIGURES
Measurements contain 2 types of information -
Magnitude of the measurement - Uncertainty of the
measurement - Only one uncertain or estimated
digit should be reported Significant Figures
digits known with certainty one uncertain
digit
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RULES FOR SIGNIFICANT FIGURES
1. Nonzero integers are always significant 2.
Leading zeros are not significant 0.0045 (2
sig. figs.) 0.00007895 (4 sig. figs.) The
zeros simply indicate the position of the decimal
point 3. Captive zeros (between nonzero digits)
are always significant 1.0025 (5 sig figs.)
12000587 (8 sig figs)
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RULES FOR SIGNIFICANT FIGURES
4. Trailing zeros (at the right end of a number)
are significant only if the number contains a
decimal point 2.3400 (5 sig figs) 23400
(3 sig figs) 5. Exact numbers (not obtained
from measurements) are assumed to have
infinite number of significant figures
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RULES FOR SIGNIFICANT FIGURES
Rounding off Numbers 1. In a series of
calculations, carry the extra digits through
to the final result before rounding off to the
required significant figures 2. If the
first digit to be removed is less than 5, the
preceding digit remains the same (2.53 rounds
to 2.5 and 1.24 rounds to 1.2)
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RULES FOR SIGNIFICANT FIGURES
Rounding off Numbers 3. If the first digit to be
removed is greater than 5, the preceding
digit increases by 1 (2.56 rounds to 2.6
and 1.27 rounds to 1.3) 4. If the digit to be
removed is exactly 5 - The preceding number is
increased by 1 if that results in an even
number (2.55 rounds to 2.6 and 1.35000
rounds to 1.4) - The preceding number remains the
same if that results in an odd number (2.45
rounds to 2.4 and 1.25000 rounds to 1.2)
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RULES FOR SIGNIFICANT FIGURES
- The certainty of the calculated quantity is
limited by the least certain measurement, which
determines the final number of significant
figures
Multiplication and Division - The result
contains the same number of significant figures
as the measurement with the least number of
significant figures 2.0456 x 4.02 8.223312
8.22 3.20014 1.2 2.6667833 2.7
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RULES FOR SIGNIFICANT FIGURES
- The certainty of the calculated quantity is
limited by the least certain measurement, which
determines the final number of significant
figures
Addition and Subtraction - The result contains
the same number of decimal places as the
measurement with the least number of decimal
places
4.03
5.5
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SCIENTIFIC NOTATION
- Used to express too large or too small numbers
(with many zeros) in compact form - The product
of a decimal number between 1 and 10 (the
coefficient) and 10 raised to a power
(exponential term) 24,000,000,000,000 2.4 x
1013
coefficient
Exponent (power)
0.000000458 4.58 x 10-7
Exponential term
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SCIENTIFIC NOTATION
- Provides a convenient way of writing the
required number of significant figures 6300000
in 4 significant figures 6.300 x 106 2400 in
3 significant figures 2.40 x 103 0.0003 in 2
significant figures 3.0 x 10-4
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SCIENTIFIC NOTATION
- Add exponents when multiplying exponential
terms (5.4 x 104) x (1.23 x 102) (5.4 x 1.23)
x 10 42 6.6 x 106 - Subtract exponents when
dividing exponential terms (5.4 x 104)/(1.23 x
102) (5.4/1.23) x 10 4-2 4.4 x 102
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DENSITY
- The amount of mass in a unit volume of a
substance
Ratio of mass to volume
Density
Units Solids grams per cubic centimeter
(g/cm3) Liquids grams per milliliter
(g/mL) Gases grams per liter (g/L)
- Density of 2.3 g/mL implies 2.3 grams per 1 mL
- Density usually changes with change in
temperature
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DENSITY
The amount of mass in a unit volume of a substance
For a given liquid - Objects with density less
than that of the liquid will float - Objects
with density greater than that of the liquid will
sink - Objects with density equal to that of the
liquid will remain stationary (neither float nor
sink)
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TEMPERATURE
- The degree of hotness or coldness of a body
or environment - 3 common temperature
scales Metric system Celsius and
Kelvin English system Fahrenheit
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TEMPERATURE
Celsius Scale (oC) - Reference points are the
boiling and freezing points of water (0oC and
100oC) - 100 degree interval Kelvin Scale (K)
- Is the SI unit of temperature (no degree
sign) - The lowest attainable temperature on the
Kelvin scale is 0 (-273 oC) referred to as the
absolute zero Fahrenheit Scale (oF) - Water
freezes at 32oF and boils at 212oF - 180 degree
interval
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TEMPERATURE
or
or
10o, 40o, 60o are considered as 2 significant
figures 100o is considered as 3 significant
figures
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LOGARITHMS
n 10a implies log n a - The logarithm
(base 10) of n is equal to a (written as log on
calculators) log 1000 3 since 1000 103 log
0.01 -2 since 0.01 10-2
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LOGARITHMS
log 436 2.639 2 is the characteristic 0.639 is
the mantissa - The number of digits in the
mantissa should be equal to the number of
significant figures in the original number
(436) log 4368 3.6403 log 0.4368 -0.3597
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ANTILOGARITHMS
n 10a implies antilog a n - n is the
antilogarithm of a (written as antilog or 10x or
INV log on calculators) antilog 3 1000 since
103 1000 antilog -2 0.01 since 10-2 0.01
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ANTILOGARITHMS
antilog 2.639 436 - The number of significant
figures in the answer should be equal to the
number of digits in the mantissa antilog 6.65
4466835.922 4.5 x 106 antilog -3.230
0.0005888436 5.89 x 10-4
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ERRORS
- Two classes of experimental errors systematic
and random Systematic Error - Also called
determinate error - Repeatable in a series of
measurements - Can be detected and
corrected Examples uncalibrated buret, pipet,
analytical balance, pH meter power fluctuations,
temperature variations
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ERRORS
- Two classes of experimental errors systematic
and random Randon Error - Also called
indeterminate error - Always present and cannot
be corrected Examples Taking readings from an
instrument, reading between markings
(interpolation), electrical noise in instruments
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ERRORS
Precision - Provides information on how closely
individual measurements agree with one
another (measure of reproducibility of a
result) Accuracy - Refers to how closely
individual measurements agree with the true
value (correct value) (systematic errors reduce
the accuracy of a measurement) - Precise
measurements may NOT be accurate - Our goal is
to be accurate and precise
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ERRORS
Absolute Uncertainty - The margin of uncertainty
associated with a measurement - If estimated
uncertainty in a buret reading is 0.05 mL
then absolute uncertainty 0.05 mL - If
estimated uncertainty in an analytical balance is
0.0001 g then absolute uncertainty 0.0001g
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ERRORS
Relative Uncertainty - Compares absolute
uncertainty with its associated measurement -
Dimensionless
Percent Relative Uncertainty Relative
Uncertainty x 100
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ERRORS
For a buret reading of 41.45 0.05 mL Absolute
uncertainty 0.05 mL
Percent Relative Uncertainty 0.001 x 100 0.1
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BALANCING CHEMICAL EQUATIONS
- Whole numbers are placed on the left side of
the formula (called coefficients) to balance the
equation (subscripts remain unchanged) - The
coefficients in a chemical equation are the
smallest set of whole numbers that balance the
equation
C2H5OH(l) O2(g)
2CO2(g) H2O(g)
2 C atoms
2 C atoms
Place the coefficient 2 in front of CO2 to
balance C atoms
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BALANCING CHEMICAL EQUATIONS
- Whole numbers are placed on the left side of
the formula (called coefficients) to balance the
equation (subscripts remain unchanged) - The
coefficients in a chemical equation are the
smallest set of whole numbers that balance the
equation
C2H5OH(l) O2(g)
2CO2(g) 3H2O(g)
3(1x2)6 H atoms
(51)6 H atoms
Place 3 in front of H2O to balance H atoms
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BALANCING CHEMICAL EQUATIONS
- Whole numbers are placed on the left side of
the formula (called coefficients) to balance the
equation (subscripts remain unchanged) - The
coefficients in a chemical equation are the
smallest set of whole numbers that balance the
equation
C2H5OH(l) 3O2(g)
2CO2(g) 3H2O(g)
1(3x2)7 O atoms
(2x2)37 O atoms
Place 3 in front of O2 to balance O atoms
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BALANCING CHEMICAL EQUATIONS
- Whole numbers are placed on the left side of
the formula (called coefficients) to balance the
equation (subscripts remain unchanged) - The
coefficients in a chemical equation are the
smallest set of whole numbers that balance the
equation
C2H5OH(l) 3O2(g)
2CO2(g) 3H2O(g)
2 C atoms (51)6 H atoms 1(3x2)7 O atoms
2 C atoms (3x2)6 H atoms (2x2)37 O atoms
Check to make sure equation is balanced When the
coefficient is 1, it is not written
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BALANCING CHEMICAL EQUATIONS
- States of reactants and products - Physical
states of reactants and products are represented
by (g) gas (l) liquid (s) solid (aq) aqueous
or water solution
C2H5OH(l) 3O2(g) ? 2CO2(g) 3H2O(g)
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BALANCING CHEMICAL EQUATIONS
Balance the following chemical equations Fe(s)
O2(g) ? Fe2O3(s) C12H22O11(s) O2(g) ?
CO2(g) H2O(g) (NH4)2Cr2O7(s) ? Cr2O3(s)
N2(g) H2O(g)
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MOLAR MASS
- Add atomic masses to get the formula mass (in
amu) molar mass (in g/mol) - That is the
mass, in g, of 1 mole of the substance 1 mole
6.02214179 x 1023 entities (atoms or molecules)
Usually rounded to 6.022 x 1023 (Avogadros
number) This implies that 6.022 x 1023 amu 1.00
g Atomic mass (amu) mass of 1 atom molar mass
(g) mass of 6.022 x 1023 atoms
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MOLAR MASS
Calculate the mass of 2.4 moles of NaNO3 Molar
mass NaNO3 22.99 14.01 3(16.00) 85.00 g
/mol NaNO3
204 g NaNO3 2.0 x 102 g NaNO3
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CHEMICAL FORMULA
Consider Na2S2O3 - Two atoms of sodium, two
atoms of sulfur, and three atoms of oxygen are
present in one molecule of Na2S2O3 - Two moles
of sodium, two moles of sulfur, and three moles
of oxygen are are present in one mole of Na2S2O3
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CHEMICAL FORMULA
How many moles of sodium atoms, sulfur atoms, and
oxygen atoms are present in 1.8 moles of a sample
of Na2S2O3? I mole of Na2S2O3 contains 2 moles
of Na, 2 moles of S, and 3 moles of O
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CHEMICAL CALCULATIONS
Calculate the number of molecules present in
0.075 g of urea, (NH2)2CO Given mass of urea
- convert to moles of urea using molar mass -
convert to molecules of urea using Avogadros
number
7.5 x 1020 molecules (NH2)2CO
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CHEMICAL CALCULATIONS
How many grams of carbon are present in a 0.125 g
of vitamin C, C6H8O6 Given mass of vitamin C
- convert to moles of vitamin C using molar
mass - convert to moles of C (1 mole C6H8O6
contains 6 moles C) - convert moles carbon to g
carbon using atomic mass
0.0511 g carbon
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CHEMICAL EQUATIONS (STOICHIOMETRIC CALCULATIONS)
Given C3H8(g) 5O2(g) ? 3CO2(g)
4H2O(g) - 1 molecule of C3H8 reacts with 5
molecules of O2 to produce 3 molecules of CO2 and
4 molecules of H2O - 1 mole of C3H8 reacts with
5 moles of O2 to produce 3 moles of CO2 and 4
moles of H2O
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CHEMICAL EQUATIONS (STOICHIOMETRIC CALCULATIONS)
Given C3H8(g) 5O2(g) ? 3CO2(g)
4H2O(g) What mass of oxygen will react with 96.1
g of propane?
- make sure the equation is balanced - calculate
moles of propane from given mass and molar mass -
determine moles of oxygen from mole ratio
(stoichiometry) - calculate mass of oxygen
349 g O2
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CHEMICAL EQUATIONS (STOICHIOMETRIC CALCULATIONS)
Given C3H8(g) 5O2(g) ? 3CO2(g)
4H2O(g) What mass of CO2 will be produced from
96.1 g of propane?
- make sure the equation is balanced - calculate
moles of propane from given mass and molar mass -
determine moles of CO2 from mole ratio
(stoichiometry) - calculate mass of CO2
288 g CO2
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CONCENTRATION OF SOLUTIONS
- The amount of solute dissolved in a given
quantity of solvent or solution Molarity
(M) - The number of moles of solute per liter of
solution
- A solution of 1.00 M (read as 1.00 molar)
contains 1.00 mole of solute per liter of
solution
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CONCENTRATION OF SOLUTIONS

• Calculate the molarity of a solution made by
  dissolving 2.56 g of
• NaCl in enough water to make 2.00 L of solution
• - Calculate moles of NaCl using grams and molar
  mass
• Convert volume of solution to liters
• - Calculate molarity using moles and liters

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CONCENTRATION OF SOLUTIONS
After dissolving 1.56 g of NaOH in a certain
volume of water, the resulting solution had a
concentration of 1.60 M. Calculate the volume of
the resulting NaOH solution - Convert grams NaOH
to moles using molar mass - Calculate volume (L)
using moles and molarity
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CONCENTRATION OF IONS
Consider 1.00 M NaCl 1.00 M Na and 1.00 M
Cl- 1.00 M ZnCl2 1.00 M Zn2 and 2.00 M
Cl- 1.00 M Na2SO4 2.00 Na and 1.00 M SO42-
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CONCENTRATION OF IONS
Calculate the number of moles of Na and SO42-
ions in 1.50 L of 0.0150 M Na2SO4
solution 0.0150 M Na2SO4 solution contains 2 x
0.0150 M Na ions and 0.0150 M SO42- ions Moles
Na 2 x 0.0150 M x 1.50 L 0.0450 mol
Na Moles SO42- 0.0150 M x 1.50 L 0.0225 mol
SO42-
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PERCENT COMPOSITION
- May also be represented by (m/m) mass of
solution mass of solute mass of solvent
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PERCENT COMPOSITION
A sugar solution is made by dissolving 5.8 g of
sugar in 82.5 g of water. Calculate the percent
by mass concentration of sugar.
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PERCENT COMPOSITION
- May also be represented by (v/v) - Due to the
way molecules are packed and differences in
distances between molecules (bond lengths), the
volume of the resulting solution is almost
always less than the sum of the volume of
solute and the volume of solvent
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PERCENT COMPOSITION
Calculate the volume percent of solute if 345 mL
of ethyl alcohol is dissolved in enough water to
produce 1257 mL of solution
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PARTS PER MILLION (PPM)
Percent can be defined as parts per hundred
1 ppm 1 µg/mL or 1 mg/L
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PARTS PER MILLION (PPM)
If 0.250 L of aqueous solution with a density of
1.00 g/mL contains 13.7 µg of pesticide, express
the concentration of pesticide in ppm ppm
µg/mL 0.250 L 250 mL Density 1.00 g/mL
Implies mass solution 250 g
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PARTS PER BILLION (PPB)
1 ppb 1 ng/mL or 1 µg/L
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PARTS PER MILLION (PPB)
If 0.250 L of aqueous solution with a density of
1.00 g/mL contains 13.7 µg of pesticide, express
the concentration of pesticide in ppb ppm
µg/L Volume of solution 0.250 L Density 1.00
g/mL Implies mass solution 250 g
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DILUTION
- Consider a stock solution of concentration M1
and volume V1 - If water is added to dilute to a
new concentration M2 and volume V2 - moles
before dilution moles after dilution - Implies
that M1V1 M2V2
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DILUTION
Calculate the volume of 3.50 M HCl needed to
prepare 500.0 mL of 0.100 M HCl (3.50 M)(V1)
(0.100 M)(500.0 mL) V1 14.3 mL
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CONCENTRATION OF SOLUTIONS
Mole Fraction (?) - Fraction of moles of a
component of solution
The sum of mole fractions of all components 1
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CONCENTRATION OF SOLUTIONS
Given that the total moles of an aqueous solution
of NaCl and other solutes is 1.75 mol. Calculate
the mole fraction of NaCl if the solution
contains 4.56 g NaCl.
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MOLALITY (m)
Moles of solute per kg of solvent Unit m or
molal
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MOLALITY (m)
What is the molality of a solution that contains
2.50 g NaCl in 100.0 g water? - Calculate moles
NaCl - Convert g water to kg water - Divide to
get molality
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CONVERTING CONCENTRATION UNITS
Calculate the molality of a 6.75 (m/m) solution
of ethanol (C2H5OH) in water Mass water 100 g
solution 6.75 g ethanol 93.25 g water
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CONVERTING CONCENTRATION UNITS
Calculate the mole fraction of a 6.75 (m/m)
solution of ethanol (C2H5OH) in water Mass
water 100 g solution 6.75 g ethanol 93.25 g
water
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CONVERTING CONCENTRATION UNITS
Practice Question Given that the mole fraction
of ammonia (NH3) in water is 0.088 Calculate the
molality of the ammonia solution
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CONVERTING CONCENTRATION UNITS
- Molarity is temperature dependent (changes
with change in temperature) - Volume increases
with increase in temperature hence molarity
decreases On the other hand - Molality - Mass
percent - Mole fraction are temperature
independent
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CHEMICAL EQUILIBRIUM
- Occurs when there is product build-up during a
chemical reaction - The product molecules
interact with one another to re-produce reactants
forward reaction
A B
C D
reverse reaction
Chemical Equilibrium - When the rate of product
formation (forward reaction) is equal to the
rate of reactant formation (reverse reaction)
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CHEMICAL EQUILIBRIUM
- Reactant and product concentrations are usually
not equal - Such reactions are known as
reversible reactions - Forward reaction rate
decreases with time as reactants are used up -
Reverse reaction rate increases with time as
products are being formed - Concentrations are
reached when both forward and reverse rates
become equal
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EQUILIBRIUM CONSTANT
- Describes the extent of reaction in a given
system - For a chemical reaction of the form aA
bB ? cC dD - The equilibrium constant
(Keq) is given by
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EQUILIBRIUM CONSTANT
- denotes concentration in moles/liter (M) -
Product concentrations in the numerator -
Reactant concentrations in the denominator -
Concentrations are raised to the powers of the
respective coefficients - Gases (in bars) and
substances in solution (in mol/L) are written in
Keq expressions
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EQUILIBRIUM CONSTANT
- Pure solids, pure liquids (e.g. water), and
solvents are not written since they are
constant - Keq changes with change in
temperature - For exothermic forward reactions
(heat released) Keq decreases with increasing
temperature - For endothermic forward reactions
(heat absorbed) Keq increases with increasing
temperature
84
EQUILIBRIUM CONSTANT
Large Keq - Greater product concentrations than
reactant concentrations - Equilibrium position
lies to the right Small Keq - Smaller product
concentrations than reactant concentrations -
Equilibrium position lies to the
left Intermediate Keq (near unity) - Both
products and reactants are in significant
amounts - Equilibrium position lies neither to
the right nor to the left
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EQUILIBRIUM CONSTANT
- Longer arrows can be used to indicate the
predominant species - Longer forward reaction
arrow for large Keq - Longer reverse reaction
arrow for small Keq CO2 H2O
H2CO3
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LE CHATELIERS PRICIPLE
- If a stress (change of conditions) is applied
to a system in equilibrium the system will
readjust (change the equilibrium position) in
the direction that best reduces the stress
imposed on the system - If more products form
as a result of the applied stress the
equilibrium is said to have shifted to the
right - If more reactants form as a result of
the applied stress the equilibrium is said to
have shifted to the left
87
LE CHATELIERS PRICIPLE
Concentration Changes For a reaction mixture at
equilibrium - Addition of reactant(s) shifts the
equilibrium position to the right - Removal of
product(s) shifts the equilibrium position to the
right - Addition of product(s) shifts the
equilibrium position to the left - Removal of
reactant(s) shifts the equilibrium position to
the left
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LE CHATELIERS PRICIPLE
Temperature Changes Exothermic Reactions - Heat
is a product - Increase in temperature shifts
the equilibrium position to the left - Decrease
in temperature shifts the equilibrium position to
the right
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LE CHATELIERS PRICIPLE
Temperature Changes Endothermic Reactions -
Heat is a reactant - Increase in temperature
shifts the equilibrium position to the right -
Decrease in temperature shifts the equilibrium
position to the left
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LE CHATELIERS PRICIPLE
Pressure Changes - Gases must be involved in the
chemical reaction - The total number of moles
of the gaseous state must change - Equilibrium
is shifted in the direction of fewer moles
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LE CHATELIERS PRICIPLE
Pressure Changes Higher moles of gaseous
reactants than products - Increase in pressure
shifts the equilibrium position to the right -
Decrease in pressure shifts the equilibrium
position to the left Higher moles of gaseous
products than reactants - Increase in pressure
shifts the equilibrium position to the left -
Decrease in pressure shifts the equilibrium
position to the right
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LE CHATELIERS PRICIPLE
Pressure Changes No change in equilibrium
position occurs if - There is no reactant nor
product in the gaseous state - Number of moles
of gaseous reactants equals number of moles of
gaseous products - Pressure is increased by
adding a nonreactive (inert) gas
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LE CHATELIERS PRICIPLE
Addition of Catalysts - Catalysts do not change
equilibrium positions - Catalysts speed up both
forward and reverse reactions so have no net
effect - Catalysts allow equilibrium to be
established more quickly by lowering the
activation energy