Unit 0 Basics of Chemistry - Ms. DiOrio's AP Chemistry...
Transcript of Unit 0 Basics of Chemistry - Ms. DiOrio's AP Chemistry...
Unit 0Basics of Chemistry
Ms. DiOrio
Rm 109
Ch. 1: Chemical Foundations
Lab Safety and Techniques
Lab Safety
• Waft
• Always wear:• Safety Goggles
• Hair Tied Back
• Closed toed shoes
• Apron or gloves when necessary
• Never point heated substance towards another student
• Rinse acid or base spill on skin immediately with water!
• Notify instructor of spills, breakage, fire, etc.
• Check glassware to see if cool• Always wash hands when finished with
lab• NO HORSEPLAY!
Lab Techniques
• Massing on Balance
• Lighting Bunsen Burner
• Reading Glassware
• Titrations
• Preparation of solution from solid
• Preparation of solution from concentrated solutions
Lab Glassware
• Buchner funnel
• Filtration flask
• Volumetric flask
• Volumetric pipet
• Mohr pipet
Lab Equipment
• Drying oven
• Spectrophotometer
• Hot/stir plate
• Desiccator
Measurement
SI Units
• Systéme International d’Unités
• A different base unit is used for each quantity
Metric System
• Prefixes convert the base units into units that are appropriate for the item being measured
Mass vs. Weight
Mass• Measures the amount of material• Balance UsedWeight• The force of the object due to gravity• Scale Used
• **You can be weightless but you cannot be massless
Temperature
• A measure of the average kinetic energy of the particles in a sample
• Measured in Kelvin, Celsius, and Fahrenheit
Temperature
Celsius• Based on the properties of water (0℃ melting point and 100℃ boiling point)Kelvin• Based on the properties of gases• There are no negative Kelvin temperaturesFahrenheit• No used in scientific measurements
Temperature Conversions
MEMORIZE!!!!
℉ =95 ℃ + 32
℃ =59 (℉ − 32)
𝐾 = ℃ + 273.15
Volume
• The most commonly used metric units for volume are the liter (L) and the milliliter (mL)• A liter is a cube 1 dm long on each side
• A milliliter is a cube 1 cm long on each side
Derived Units
• Made by combining the seven SI units to make new units
• Ex: Volume, density, speed, force, pressure, and energy
Measurement Derived Unit Symbol
Density Grams/milliliter g/mL
Concentration Moles/liter mol/L
Area Meters cubed m2
Volume Decimeter squared dm3 ( = Liter )
Speed (velocity) Meters/second m/s
Acceleration Meters/second sq. m/s2
Density
• A physical property used to help characterize a substance
• Units of g/cm3 or g/mL
Uncertainty in Measurements
• Different measuring devices have different uses and different degrees of accuracy and precision
Significant Figures
• Digits that are measured
• Determines how precise and instrument is
• When measuring à go one place beyond markers (except with electronic equipment)
• When calculating à round to sig figs at the end of calculation• You do not want to overstate the accuracy of answers
• Avoiding rounding error throughout calculations
Significant Figures
1. All nonzero digits are significant
2. Zeroes between two significant figures are themselves significant
3. Zeroes at the beginning of a number are never significant
4. Zeroes at the end of a number are significant if a decimal point is written in the number
Atlantic-Pacific Rule
1410
25085932
2 1
1 -
2 1 -
5 4 3 2 1
2 sig figs
1 sig fig
2 sig figs
5 sig figs
1.4.080.002508.593005400.
- 12 sig figs
1 sig fig
3 sig figs
6 sig figs
4 sig figs
1 2 3 4 5 6
- - - 1 2 3
1 2 3 4
1 2
Significant Figures
Multiplication / Division
• Answers are rounded to the least number of sig figs in any of the numbers used in the calculation
Addition / Subtraction
• Answers are rounded to the least significant decimal place
**Remember, carry all digits through until the very end of a multistep calculation
Sig Fig Review
v 0.0234v 0.0230v 0.12v 1.000
v 1.2v 100.5v 5050v 5000
3 Sig Figs
3 Sig Figs
2 Sig Figs
4 Sig Figs
2 Sig Figs
4 Sig Figs
3 Sig Figs
1 Sig Fig
Prefix Conversion Steps
1. Write down what is given in scientific notation
2. Put the units given in the denominator3. Put the units wanted in the numerator4. Put a 1 with the larger unit and the positive 10x
difference with the smaller unit5. Solve
1. Exponents in numerator are added, exponents in denominator are subtracted
Prefix Conversion Practice
v Convert 467g into Mg
vConvert 0.0056 pm into cm
4.67 x 102 g 1 Mg106 g
5.6 x 10-3 pm 1 cm109 pm
= 4.67 x 10-4 Mg
= 5.6 x 10-12 cm
Dimensional Analysis
• Converts one quantity to another
• Most commonly utilized conversion factors• Ex: 1 in = 2.54 cm
Dimensional Analysis Steps
1. Start with the number and units that are given
2. Set up a conversion with the units given in the denominator to cancel out
3. Put units you want in the numerator
4. Solve
Dimensional Analysis with Cubes
• Don’t forget, if a unit is cubed your conversion must also be cubed
Convert 4.5 mm3 to nm3
4.5 mm3 (106 nm)3
(1 mm)3 = 4.5 x 1018 nm3
Challenge Question
• What is the length in micrometers of a 2.4 g cube of gold whose density is 19.3 g/mL? 2.4 g 1 mL 1 cm3
19.3 g 1 mL= 0.124 cm3
Length = 0.124 cm31
= 0.499 cm 104 nm1 cm
= 0.499 x 104 nm = 5.0 x 103 nm
Error Analysis
Error
Accuracy
• Refers to the proximity of a measurement to the true value of a quantity
Precision
• Refers to the proximity of several measurements to each other
Error
• The difference between a measured value and a true value
Error & Accuracy vs. Precision
• If the significant digits of a measurement (expressed in however many sig figs) agree with the true value, then the measurement is accurate
• A poorly used instrument which delivers 6 significant digits may be precisebut will not be helpful in obtaining an accurate value
• In general, the greater the number of measurements, the closer the estimate and the smaller the error
Types of Error
Random• Non-reproducible and not the same magnitude for any two measurements (except
by chance)• Often tend to cancel out if enough samples are taken• Repetition of work is the best way to minimize the effects on accuracyDeterminate• Affects each measurement in the same way and to the same extent• Includes systematic, proportional, and constant error
Types of Determinate ErrorTypes of Error Explanation Can be reduced by…
SystematicOften occurs in one
directionDifferent magnitude
Calibration of equipment
Proportional Constant relative error such as incorrect concentration
Using controls with known values
ConstantReproducible error
Seen with fixed amount of reagents used
Using a “blank”
Error
All measurements have some amount of error.• This is due to both human error (estimation)
and error built into an instrument• Most instruments are labeled with the
uncertainty
**If not stated, assume uncertainty is ±1 in the last significant digit
Instrument Uncertainty (±)
Balance, top loading 0.01 g
Balance, analytical 0.001 g
Thermometer 0.2
10 mL graduated cylinder 0.1 mL
25 mL graduated cylinder 0.3 mL
50 mL graduated cylinder 0.4 mL
100 mL graduated cylinder 0.6 mL
100 mL gas measuring tube 0.2 mL
50 mL buret 0.05 mL
1 mL pipet 0.006 mL
10 mL pipet 0.02 mL
25 mL pipet 0.03 mL
100 mL volumetric flask 0.08 mL
250 mL volumetric flask 0.12 mL
Logic of Uncertainty in Calculations
Addition and Subtraction:
23.2 ± 0.3 mL + 5.6 ± 0.1 mL =
Smallest Possible Value: Largest Possible Value:
Range:
Report as:
28.4 mL 29.2 mL0.8 mL28.8 ± 0.4 mL
Uncertainty with Addition and Subtraction
• The previous example only works if the direction of both errors is the same. To correct for the sign of the uncertainty, it is more correct to use the following equation:
Example:
𝑈𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦𝑖𝑛𝑟𝑒𝑠𝑢𝑙𝑡 = (𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦> )?+(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦? )? + ⋯+ (𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦A )?�
23.2 ± 0.3 mL + 5.6 ± 0.1 mL = 𝑈𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦 = (0.3)?+(0.1)?�
= ± 0.3 mL
28.8 ± 0.3 mL
Uncertainty with Multiplication and Division
• The error propagation with these operations is not as straight forward
• Use the following equation:
𝑈𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦𝑖𝑛𝑟𝑒𝑠𝑢𝑙𝑡 = 𝐶𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑𝑟𝑒𝑠𝑢𝑙𝑡𝑥(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦> )?
(𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡> )?+
(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦? )?
(𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡? )?+ ⋯+
(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦A )?
(𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡A )?�
Q-Test: Rejection of Data
• Sometime one value in a set of data appears different in a non-random way The q-test is an objective statistical test for deciding if a result should be rejected as an “outlier”
• Only the survivors of the q-test are used in the mean, standard deviation, and t-test
Q-Test
1. Arrange the data in order2. Calculate the difference (d) between the suspect value
and the nearest neighbor3. Calculate the range (r) between the suspect value and
the farthest neighbor4. Calculate the experimental value for Qexp
QIJK = dr
5. Compare Qexp to Qtable
6. Reject the data point if Qexp > Qtable
Confidence Level90% 95% 99%
n Q Q Q3 0.89 0.94 0.994 0.68 0.76 0.895 0.56 0.64 0.786 0.48 0.56 0.78
7 0.43 0.43 0.64For n>7, use value for 7
Q-Test Practice
• Should any of the data be thrown out at a 95% confidence level?
Trial Volume (mL)1 14.82 12.63 13.94 13.25 14.1
Highest Value: 14.8
Lowest Value: 12.6
QIJK = dr
Confidence Level90% 95% 99%
n Q Q Q3 0.89 0.94 0.994 0.68 0.76 0.89
5 0.56 0.64 0.786 0.48 0.56 0.78
7 0.43 0.43 0.64For n>7, use value for 7
QIJK = 14.8−14.114.8 − 12.6
= 0.32
0.32 < 0.64
QIJK = 13.2−12.614.8 − 12.6
= 0.27
0.27 < 0.64
Standard Deviation (s or 𝜎)
• A method for calculating the precision of a series of measurements
• Where x is the measurement and x is the mean (measurement-mean) and n is the number of samples
**The smaller the standard deviation, the more precise the data
Calculate the Standard Deviation
Mean:
Standard Deviation:
Trial Volume (mL)1 45.62 46.23 45.74 45.2
45.675
=(45.6 − 45.675)?+(46.2 − 45.675)?+(45.7 − 45.675)?+(45.2 − 45.675)?
4 − 1�
s = 0.411
One Sample T-Test
• Used to estimate the reliability of the derived result from a small set of data• Sets up a “confidence interval” for a result (the probability that the true value is
close to the mean (𝜇))
• x = mean• t = value from table• s = standard deviation• n = number of trials
𝑥 ± 𝑡(𝑠𝑛�)
**Report results using this formula
T-Table
• Confidence level or % probability that the actual value lies within the interval from the mean
T-Test Practice
• Use a T-Test at a 95% confidence level to find the mean and range of uncertainty.
Trial Volume (mL)1 45.62 46.23 45.74 45.2
𝑥 ± 𝑡(𝑠𝑛�)
45.7 ± 3.18(0.4114�)
𝟒𝟓. 𝟕 ± 𝟎. 𝟕𝒎𝑳
The Beauty of Modern Technology
• Standard Deviation and T-Tests can also be calculated in excel or with a TI graphing calculator
Excel
=STDEV.S(range)
TI- Press [2nd], [LIST], - Scroll to MATH and select 7:stdDev(- Input in brackets within parenthesesseparated by commas
The Beauty of Modern Technology
• Standard Deviation and T-Tests can also be calculated in excel or with a TI graphing calculator
Excel
=T.TEST(range)
TI
- Input Data to L1- Press [STAT] - Scroll to TESTS and select 2:T-Test- Scroll down to Calculate
Summary
• % Error – simple calculation to determine accuracy (known actual value)
• Uncertainty in Measurements – table of values for instruments, equations for add/sub and mult/div
• Q-Test – determines if data can be thrown out of statistical analysis
• Standard Deviation – complex determination of error and analysis of mean
• T-Test – used in small data sets to determine value of mean in relation to true value
Ch. 2: Atoms, Molecules and Ions
Classification of Matter
Matter
• We define matter as anything that has mass and takes up space
Atoms of an element Molecules of an element Molecules of a compound Mixture of elements and a compound
Matter
• Atoms are the building blocks of matter
• Each element is made of the same kind of atom
• A compound is made of two or more different kinds of elements
States of Matter
Properties and Changes of Matter
Types of Properties
• Physical Properties• Can be obtained without changing a substance into another substance
• Boiling point, density, mass, volume, etc.
• Chemical Properties• Can only be observed when a substance is changed into another substance
• Flammability, corrosiveness, reactivity with acid, etc.
Types of Properties
• Intensive Properties• Are independent of the amount of the substance that is present
• Density, boiling point, color, etc.
• Extensive Properties• Depend upon the amount of the substance present
• Mas Volume, energy, etc.
Types of Changes
• Physical Changes• These are changes in matter that do not change the composition of a substance
• Changes of state, temperature, volume, etc.
• Chemical Changes• Chemical changes result in new substances
• Combustion, oxidation, decomposition, etc.
Separation of Mixtures
Evaporation
• Evaporation uses difference in the boiling points of substances to separate a mixture into its components• The substance with the
lower boiling point is lost to the atmosphere
Distillation
• Distillation uses difference in the boiling points of substances to separate a homogeneous mixture into its components
• All components are retained
Filtration
• In filtration, solid substances are separated from liquids and solutions
Types of Filtration
• Decanting• Involves pouring off the top layer of liquid
• The simplest way to separate a liquid and solid and should only be used for qualitative work
• Gravity Filtration• Uses the natural downward flow of the liquid through a porous medium (usually paper) to
separate components
• Vacuum Filtration• Creates a partial vacuum in the collection flask to speed up separation
• Uses Buchner funnel and filtration flask
Chromatography
• Separates substances on the basis of difference in polarity
• Types of chromatograph differ by their mobile and stationary phases
Absorbent vs. Adsorbent
• Absorbent• The ability to carry a liquid in small chambers
• Ex: sponge
• Related to volume and diffusion
• Adsorbent• When particles get caught in small pores by weak Van der Waals forces (i.e. London
Dispersion) to form a layer• Ex: Chromatography
• Related to surface area and adhesion
Chromatography
Substances with similar structures (polar or nonpolar) will be attracted to each other
Types of Chromatography
• Thin Layer Chromatography (TLC)• Simple method using a gel or silica as the stationary phase and solvent as mobile phase
• Paper Chromatography• Similar to TLC but using filter paper as the stationary phase
Uses for Chromatography
• Identify unknowns based on comparison
• Determine the purity of a substance
• Monitor the progress of a reaction
Calculating Rf Values
Retention Factor (Rf)
• Value to help determine the relative distance of compounds in TLC
𝑅Y = 𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒𝑡𝑟𝑎𝑣𝑒𝑙𝑒𝑑𝑏𝑦𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒𝑡𝑟𝑎𝑣𝑒𝑙𝑒𝑑𝑏𝑦𝑠𝑜𝑙𝑣𝑒𝑛𝑡
• Best for compare compounds on the same stationary phase• Can compare if TLC has constant solvent system, stationary
phase, amount of material spotted, and temperature.
Calculating Rf Values
**Always mark in pencil the origin line (before starting) and the solvent front (after sample is run before solvent dries)
History of Atomic Theory
Democritus (~400 BC)
• First suggested the existence of atoms
• All matter is made up of small indestructible particles
John Dalton (1803-1807)
• First Atomic Theory based on proof.
1. All matter is made up of atoms
2. Atoms of the same element are the same, atoms of different elements are different
3. Atoms combine in small whole number ratios to form compounds
4. Atoms are rearranged in a chemical reaction
Thomson’s Experiment
Thomson’s Experiment
Thomson’s Experiment
• Passing an electric current makes a beam appear to move from the negative to the positive end
Thomson’s Experiment
• By adding an electric field, he found that the moving pieces were negative
J.J. Thomson (1897)
• Used cathode ray tubes to discover the electron
• Measured the charge to mass ratio of the electron
• Since atoms are neutral species, he suggested that negative electrons are floating in a positively charged medium
Millikan’s Oil Drop Experiment
• From the mass of the drop and the charge of the plates, he calculated the mass of an electron
9.1 x 10-31 kg
Radioactivity
• Bequerel (~1890s)• Discovered radioactivity by accident
• Curie (~1900s)• Isolated radioactivity
• Rutherford (~1900)• Discovered three types:• Alpha – Helium nucleus (+2 charge, large mass)• Beta – High speed electron• Gamma – High energy light
Rutherford’s Gold Foil Experiment
Rutherford’s Gold Foil Experiment
What he expected
Rutherford’s Gold Foil Experiment
Because he thought the mass was evenly distributed in the atom
Rutherford’s Gold Foil Experiment
What he got
Earnest Rutherford (1908-1913)
1. Atom contains a dense nucleus
2. The atom is mostly empty space
Niels Bohr (~1910s)
• Electrons have certain energies which allow them to stay in certain orbits around the nucleus
• Still used to explain energy levels in atoms
Modern View of the Atom
• The atom is mostly empty space
• Two regions• Nucleus – protons and neutrons
• Electron Cloud – region where you might find an electron
Quantum Model (Electron Cloud Model)
The Periodic Table
Elements
Atomic Number
Atomic Mass
Atomic Symbol
Periodic Table
Metals
Nonmetals
Metalloids
Alkali Metals
Alkali Earth Metals
Transition Metals
Halogens
Noble Gases
Inner Transition Metals
Common Charges
Chemical Bondingand Nomenclature
Chemical Bonds
Covalent Bonding• Sharing electrons• 2 non-metals• Makes molecules
Ionic Bonding• Gaining/losing electrons• Metal and non-metal (or polyatomic)• Held together by opposite charges• Makes compounds
The forces that hold atoms together
Types of Formulas
• Chemical Formula• The number and type of atoms in a
molecule
• Structural Formula• Shows the connects but not
necessarily the shape
• Skeletal Structure• Abbreviated from structural
• C and H are not written (C at bend)
• Ball and Stick Model• Shows the three dimensional shape
with accurate angles
Butane
Empirical vs. Molecular Formula
Empirical Formula
• The reduced formula of a compound
Molecular Formula
• The unreduced or actual formula of a compound
**The empirical and molecular formulas can be the same!**Empirical formula can be found from the molecular formula, but data is
needed to go from empirical to molecular
Molecular àEmpirical
• Divide all subscripts in the MF by the greatest common factor
MF Si4O8 C25H45 CO2 C8H10N4O2 Ba3(PO4)2
EF SiO2 C5H9 CO2 C4H5N2O Ba3(PO4)2
Empirical à Molecular
Formula Molar Mass GCF
MF
EF
Determine the molecular formula of a compound with an empirical formula of NH2 and a formula mass of 32.06 amu.
**EF formula and MF molar mass given1. Calculate the molar mass of the EF2. Divide MF molar mass by EF molar
mass to find the greatest common factor (GCF)
3. Multiply all subscripts in the EF by the GCF
NH2
N2H4 32.06
16.032
1
23
Ions
Cations
• Positive ions
• Lose electrons
Anions
• Negative ions
• Gain electrons
Atoms of groups of atoms with a chargeIonic solids are called salts
Polyatomic Ions
• Groups of atoms covalently bonded that have an overall charge
• Bromine and iodine are just like chlorine• Bromate (BrO3
-1), perbromate (BrO4-1)
• Adding hydrogen to polyatomics• Sulfate (SO4
-2), hydrogen sulfate or bisulfate (HSO4-1)
Hypochlorite ClO-1 -2Chlorite ClO2
-1 -1Chlorate ClO3
-1 basePerchlorate ClO4
-1 +1
Polyatomic Ions to Memorize
Naming Ionic Compounds
From formula to name:
1. Name cationa. Transition metals must have roman numeral or -ic/-ous
2. Name aniona. Monatomic ion à end in –ide
b. Polyatomic ion à leave name as is
Transition Metals
• Have multiple possible ion charges• Exceptions: Ag+1 Zn+2 Cd+2
Element (ous) (ic) NameIron (Fe) +2 +3 Ferr-Lead (Pb) +2 +4 Plumb-Tin (Sn) +2 +4 Stann-Mercury +1 +2 Mercur-Copper +1 +2 Cupr-
Examples
Name the following:
• CaS
• AlPO4
• FeS
• CuI3
• KClO4
• YBrO2
• Cr(ClO)6
Calcium sulfide
Aluminum phosphate
Iron (II) sulfide OR Ferrous sulfideCopper (III) iodide
Potassium perchlorate
Yttrium (II) bromite
Chromium (VI) hypochlorite
Naming Ionic Compounds
From name to formula:
1. Determine charge of cation and anion
2. Criss-cross charges as positive subscripts1. Put () around polyatomics with charge on the outside
3. Reduce subscripts if needed
Examples
Write the formula for the following:• Tungsten (II) nitrite• Stannous chromate• Potassium permanganate• Zinc phosphate• Sodium oxide• Aluminum hydrogen sulfate
W(NO2)2
SnCrO4
KMnO4
Na2OZn3(PO4)2
Al(HSO4)3
Naming Covalent Compounds
From formula to name:
1. Write first element with prefixa. Exception à never write mono with
the first element
2. Write second element with prefix and end in -ide
Covalent Prefixesmono- 1 hexa- 6
di- 2 hepta- 7tri- 3 octa- 8
tetra- 4 nona- 9penta- 5 deca- 10
Examples
Name the following:
• CO2
• CO
• CCl4
• N2O4
• XeF6
• N4O4
• P2O10
Carbon dioxide
Carbon monoxide
Carbon tetrachloride
Dinitrogen tetraoxide
Xenon hexafluoride
Tetranitrogen tetraoxide
Diphosphorus decaoxide
Naming Covalent Compounds
From name to formula:
1. Prefixes indicate subscript on each element
2. If no prefix on first element, assume 1
3. DO NOT REDUCE!
Covalent Prefixesmono- 1 hexa- 6
di- 2 hepta- 7tri- 3 octa- 8
tetra- 4 nona- 9penta- 5 deca- 10
Examples
Write the formula for the following:
• Dinitrogen pentachloride
• Carbon disulfide
• Triboron heptafluoride
• Phosphorus trioxide
N2Cl5CS2
B3Fl5PO3
Acids
• Substances that produce H+ when dissolved in water
• All acids begin with H
• Two types of acids:• Oxyacids (contains oxygen, usually a polyatomic)
• Non-oxyacids (can be binary acids à H with a monatomic ion, or non-oxygen polyatomics)
Naming Oxyacids
From formula to name:
1. Name polyatomic / anion1. Change -ate to –ic acid
2. Change –ite to –ous acid
2. Do NOT include hydrogen in the name
From name to formula:
1. Look for acid at end of name
2. Determine polyatomic1. If –ic acid, then use –ate polyatomis
2. If –ous acid, then use –ite polyatomis
3. Use # of H to balance out charge of polyatomic
Naming Non-oxyacids
From formula to name:
1. Write hydro-
2. Name anion1. Change –ide to –ic acid
From name to formula:
1. Name contains hydro-
2. Determine –ide anion formula
3. Use # of H to balance out charge
Examples
Name the following:
• H2CrO4
• HCl
• H2S
• HMnO4
• HCN*
• HNO2
Chromic acid
Hydrochloric acid
Hydrosulfuric acid
Permanganic acid
Hydrocyanic acid ORHydrogen cyanide
Nitrous acid
Examples
Write the formula for the following:• Hydrofluoric acid• Dichromic acid• Carbonic acid• Hydrophosphoric acid• Perchloric acid• Phosphorous acid
HFH2Cr2O7
H2CO3
HClO4
H3P
H3PO3
Hydrates
• Salts containing coordinated water
• Name must include water
Naming Hydrates
From formula to name:
1. Write name of salt
2. Write hydrate with covalent prefix to indicate # of water molecules
From name to formula:
1. Determine formula for salt
2. Put a dot, then xH2O1. X is prefix from name for # of
water molecules
Examples
Write the formula for the following:
• Calcium chloride dihydrate
• Chromium (III) nitrate hexahydrate
Write the name for the following:
• CuSO4 • 5H2O
• Na2O • 3H2O
CaCl2 • 2H2O
Cr(NO3)3 • 6H2O
Copper (II) sulfate pentahydrate
Sodium oxide trihydrate
Organic Nomenclature
Organic Molecules
• Hydrocarbon• Organic molecules containing just
hydrogen and carbon
• Substituent• General term for a functional group or
R group
• Functional Group• Group containing oxygen, nitrogen,
and/or sulfur which gives the molecule different properties/reactivities
• R group• A hydrocarbon chain of undetermined
length attached to another hydrocarbon or functional group
• Aromatics• Molecules which contain alternating
double bonds, usually cyclical
• Ex: Benzene
Covalently bonded molecules containing hydrogen, carbon, oxygen, nitrogen, and/or sulfur
Simple Organic Naming
From formula to name:
1. Use organic prefixes for # of carbons
2. End in suffix based on # hydrogens / types of bonds
From name to formula:
1. Use prefix to determine # of carbons
2. Based on suffix, use formula to determine # of hydrogen
Organic Prefixes
Meth- 1 Hex- 6Eth- 2 Hept- 7Prop- 3 Oct- 8But- 4 Non- 9Pent- 5 Dec- 10 **Use cyclo before prefix if
hydrocarbon is in a ring
Organic Suffix
Ending C-C Bond Formula-ane Single bond CnH2n+2
-ene Double bond CnH2n
-yne Triple bond CnH2n-2
Functional GroupsFunctional Group Name Example
Aldehyde ______al Ethanal
Alcohol _____ol Ethanol
Carboxylic Acid ____oic acid Ethanoic
acid
Ketone _____one Propanone
Functional GroupsFunctional Group Name Example
Ether R-yl R’-yl ether Dimethyl ether
Amine _____ylamine Methylamine
Ester R’-yl _____oate Methyl propanoate
Amide _____amide Ethanamide
Examples
Write the names for the following:
• CH3CH2NH2
• -
• CH3CH2COOCH2CH2CH2CH3
• -
ethylamine
pentanone
butyl propanoate
ethyl methyl ether
Examples
Draw the structures for the following:
• Pentanol
• Butanoic acid-
• Ethyl propyl ether
• Cyclopentanamide
Challenge Problem
IUPAC Nomenclature
• Systematic method to indicate where something is occurring in an organic structure
• Using a numbering system of the “parent” chain
IUPAC Nomenclature
From formula to name:
1. Find the longest hydrocarbon chain (This is the parent chain)
2. Number the chain so that substituents coming off the parent chain get the lowest numbered carbons
3. Write the #, substituent, -, then parent chain1. If substituent is a functional group, use functional group naming with the #
Examples
1-chlorobutane
2-methylhexane
2-butene
3-hexanol
1-bromocyclopentane
4-ethyloctane
1
1
1
1
1
1
2
2
2
2
2
2
3
3
3
3
3
3
4
4
4
4
4
4
5
5
5
5
6
6
678
Parent chain
Substituent
• Instead of putting the number for functional groups in front of the name, they can be placed immediately before the prefix ending
Alternate Naming
2-butene but-2-ene
• Instead of putting the number for functional groups in front of the name, they can be placed immediately before the prefix ending
Alternate Naming
3-hexanol hexan-3-ol
Multiple Substituents
• When the parent chain has multiple substituents, they are named ALPHABETICALLY rather than in number order
4-ethyl-2-methylhexane
1 2 3 4 5 6
ethylmethyl
Multiple Substituents
• This is where the alternate numbering comes in handy
4-ethyl-2-methylhex-3-ene
1 2 3 4 5 6
ethylmethyl