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ANALYTICAL CHEMISTRY CHEM 3811 CHAPTERS 1 AND 3 (REVIEW)

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## ANALYTICAL CHEMISTRY CHEM 3811 CHAPTERS 1 AND 3 (REVIEW)

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**ANALYTICAL CHEMISTRY CHEM 3811CHAPTERS 1 AND 3 (REVIEW)**DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university**CHAPTERS 1 AND 3**MEASUREMENTS, SIGNIFICANT FIGURES ERRORS, STOICHIOMETRY, CONCENTRATIONS**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)**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)**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)**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)**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)**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**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 d’Unites) liter, meter, gram More convenient to use**FUNDAMENTAL SI UNITS**Physical Quantity Mass Length Time Temperature Amount of substance Electric current Luminous intensity Name of Unit Kilogram Meter Second Kelvin Mole Ampere Candela Abbreviation kg m s(sec) K mol A cd**DERIVED SI UNITS**Physical Quantity Force Pressure Energy Power Frequency 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)**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**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**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**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**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**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)**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**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)**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)**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**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**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 x1013 coefficient Exponent (power) 0.000000458 = 4.58 x10-7 Exponential term**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**SCIENTIFIC NOTATION**- Add exponents when multiplying exponential terms (5.4 x 104) x (1.23 x 102) = (5.4 x 1.23) x 10 4+2 = 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**DENSITY**- The amount of mass in a unit volume of a substance Density= Ratio of mass to volume = 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**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)**TEMPERATURE**- The degree of hotness or coldness of a body or environment - 3 common temperature scales Metric system Celsius and Kelvin English system Fahrenheit**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**TEMPERATURE**or or 10o, 40o, 60o are considered as 2 significant figures 100o is considered as 3 significant figures**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**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**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**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**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**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**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**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**ERRORS**Relative Uncertainty - Compares absolute uncertainty with its associated measurement - Dimensionless Percent Relative Uncertainty = Relative Uncertainty x 100**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 %**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**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 (5+1)=6 H atoms Place 3 in front of H2O to balance H atoms**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)+3=7 O atoms Place 3 in front of O2 to balance O atoms**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 (5+1)=6 H atoms 1+(3x2)=7 O atoms 2 C atoms (3x2)=6 H atoms (2x2)+3=7 O atoms Check to make sure equation is balanced When the coefficient is 1, it is not written**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)**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)**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 (Avogadro’s 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**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**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**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