Presentation on theme: "SAT II CHEM PREP PPT Mrs. Gupta"— Presentation transcript:
1 SAT II CHEM PREP PPT Mrs. Gupta Modified from Mark Rosengarten’s Powerpoint
2 Setup of the SAT II Chem Exam 85 total questions, 1 hour (about 42 s/question)- All multiple choice, 1/4th point taken off for every incorrect answer- if you can narrow down to two choices, then guess otherwise leave blank- scoring scale from
3 What to Bring to the Exam #2 pencil, eraserNo calculators allowed (brush up on your basic math skills)Your brain. Please don’t leave it at home.:)
4 How To PrepareDO NOT CRAM. Get your studying done with by the night before. Get a good night’s sleep and have breakfast the morning of the exam.Actively participate in any and all review classes and activities offered by your teacher.
5 Matter1) Properties of Phases2) Types of Matter3) Phase Changes
6 Properties of PhasesSolids: Crystal lattice (regular geometric pattern), vibration motion onlyLiquids: particles flow past each other but are still attracted to each other.Gases: particles are small and far apart, they travel in a straight line until they hit something, they bounce off without losing any energy, they are so far apart from each other that they have effectively no attractive forces and their speed is directly proportional to the Kelvin temperature (Kinetic-Molecular Theory, Ideal Gas Theory)
7 Solids The positive and negative ions alternate in the ionic crystal latticeof NaCl.
8 Liquids When heated, the ions move faster and eventually separate from each other toform a liquid. The ions areloosely held together by theoppositely charged ions, butthe ions are moving too fastfor the crystal lattice to staytogether.
9 Gases Since all gas molecules spread out the same way, equal volumes ofgas under equal conditions oftemperature and pressure willcontain equal numbers ofmolecules of gas L of anygas at STP (1.00 atm and 273K)will contain one mole(6.02 X 1023) gas molecules.Since there is space between gasmolecules, gases are affected bychanges in pressure.
10 Types of Matter Substances (Homogeneous) Mixtures Elements (cannot be decomposed by chemical change): Al, Ne, O, Br, HCompounds (can be decomposed by chemical change): NaCl, Cu(ClO3)2, KBr, H2O, C2H6MixturesHomogeneous: Solutions (solvent + solute)Heterogeneous: soil, Italian dressing, etc.
11 ElementsA sample of lead atoms (Pb). All atoms in the sample consist of lead, so the substance is homogeneous.A sample of chlorine atoms (Cl). All atoms in the sample consist of chlorine, so the substance is homogeneous.
12 CompoundsLead has two charges listed, +2 and +4. This is a sample of lead (II) chloride (PbCl2). Two or more elements bonded in a whole-number ratio is a COMPOUND.This compound is formed from the +4 version of lead. This is lead (IV) chloride (PbCl4). Notice how both samples of lead compounds have consistent composition throughout? Compounds are homogeneous!
13 MixturesA mixture of lead atoms and chlorine atoms. They exist in no particular ratio and are not chemically combined with each other. They can be separated by physical means.A mixture of PbCl2 and PbCl4 formula units. Again, they are in no particular ratio to each other and can be separated without chemical change.
14 The Atom1) Nucleons 2) Isotopes 3) Natural Radioactivity 4) Half-Life 5) Nuclear Power 6) Electron Configuation 7) Development of the Atomic Model
15 NucleonsProtons: +1 each, determines identity of element, mass of 1 amu, determined using atomic number, nuclear chargeNeutrons: no charge, determines identity of isotope of an element, 1 amu, determined using mass number - atomic number (amu = atomic mass unit)3216S and 3316S are both isotopes of SS-32 has 16 protons and 16 neutronsS-33 has 16 protons and 17 neutronsAll atoms of S have a nuclear charge of +16 due to the 16 protons.
16 IsotopesAtoms of the same element MUST contain the same number of protons.Atoms of the same element can vary in their numbers of neutrons, therefore many different atomic masses can exist for any one element. These are called isotopes.The atomic mass on the Periodic Table is the weight-average atomic mass, taking into account the different isotope masses and their relative abundance.Rounding off the atomic mass on the Periodic Table will tell you what the most common isotope of that element is.
17 Weight-Average Atomic Mass WAM = ((% A of A/100) X Mass of A) + ((% A of B/100) X Mass of B) + …What is the WAM of an element if its isotope masses and abundances are:X-200: Mass = amu, % abundance = 20.0 %X-204: Mass = amu, % abundance = 80.0%amu = atomic mass unit (1.66 × kilograms/amu)
18 Most Common IsotopeThe weight-average atomic mass of Zinc is amu. What is the most common isotope of Zinc? Zn-65!What are the most common isotopes of: C, H and O?FACT: one atomic mass unit (1.66 × kilograms) is defined as 1/12 of the mass of an atom of C-12.
19 Natural Radioactivity Alpha DecayBeta DecayPositron DecayGamma DecayCharges of Decay ParticlesNatural decay starts with a parent nuclide that ejects a decay particle to form a daughter nuclide which is more stable than the parent nuclide was.
20 Alpha DecayThe nucleus ejects two protons and two neutrons. The atomic mass decreases by 4, the atomic number decreases by 2.23892U
21 Beta DecayA neutron decays into a proton and an electron. The electron is ejected from the nucleus as a beta particle. The atomic mass remains the same, but the atomic number increases by 1.146C
22 Positron DecayA proton is converted into a neutron and a positron. The positron is ejected by the nucleus. The mass remains the same, but the atomic number decreases by 1.5326Fe
23 Gamma DecayThe nucleus has energy levels just like electrons, but the involve a lot more energy. When the nucleus becomes more stable, a gamma ray may be released. This is a photon of high-energy light, and has no mass or charge. The atomic mass and number do not change with gamma. Gamma may occur by itself, or in conjunction with any other decay type.
25 Half-LifeHalf life is the time it takes for half of the nuclei in a radioactive sample to undergo decay.Problem Types:Going forwards in timeGoing backwards in timeRadioactive Dating
26 Going Forwards in TimeHow many grams of a 10.0 gram sample of I-131 (half-life of 8 days) will remain in 24 days?#HL = t/T = 24/8 = 3Cut 10.0g in half 3 times: 5.00, 2.50, 1.25g
27 Going Backwards in Time How many grams of a 10.0 gram sample of I-131 (half-life of 8 days) would there have been 24 days ago?#HL = t/T = 24/8 = 3Double 10.0g 3 times: 20.0, 40.0, 80.0 g
28 Radioactive DatingA sample of an ancient scroll contains 50% of the original steady-state concentration of C-14. How old is the scroll?50% = 1 HL1 HL X 5730 y/HL = 5730y
29 Nuclear Power Artificial Transmutation Particle Accelerators Nuclear FissionNuclear Fusion
30 Artificial Transmutation 4020Ca + _____ > 4019K + 11H9642Mo + 21H > 10n + _____Nuclide + Bullet --> New Element + Fragment(s)The masses and atomic numbers must add up to be the same on both sides of the arrow.
31 Nuclear Fission23592U n 9236Kr Ba n + energyThe three neutrons given off can be reabsorbed by other U-235 nuclei to continue fission as a chain reactionA tiny bit of mass is lost (mass defect) and converted into a huge amount of energy.
33 Nuclear Fusion 21H + 21H 42He + energy Two small, positively-charged nuclei smash together at high temperatures and pressures to form one larger nucleus.A small bit of mass is destroyed and converted into a huge amount of energy, more than even fission.
34 Electron Configuration Basic ConfigurationValence ElectronsElectron-Dot (Lewis Dot) DiagramsExcited vs. Ground StateRules for Electron FillingPara and DiamagneticLewis Structures and Hybridization
35 Basic ConfigurationThe number of electrons is determined from the atomic number.Look up the basic configuration below the atomic number on the periodic table. (PEL: principal energy level = shell)He: 2 (2 e- in the 1st PEL)Na: (2 e- in the 1st PEL, 8 in the 2nd and 1 in the 3rd)Br: (2 e- in the 1st PEL, 8 in the 2nd, 18 in the 3rd and 7 in the 4th)
36 Valence ElectronsThe valence electrons are responsible for all chemical bonding.The valence electrons are the electrons in the outermost PEL (shell).He: 2 (2 valence electrons)Na: (1 valence electron)Br: (7 valence electrons)The maximum number of valence electrons an atom can have is EIGHT, called a STABLE OCTET.
37 Electron-Dot Diagrams The number of dots equals the number of valence electrons.The number of unpaired valence electrons in a nonmetal tells you how many covalent bonds that atom can form with other nonmetals or how many electrons it wants to gain from metals to form an ion.The number of valence electrons in a metal tells you how many electrons the metal will lose to nonmetals to form an ion. Caution: May not work with transition metals.EXAMPLE DOT DIAGRAMS(c) 2006, Mark Rosengarten
38 Example Dot Diagrams Carbon can also have this dot diagram, which it has when it forms organic compounds.
39 Excited vs. Ground State Configurations on the Periodic Table are ground state configurations.If electrons are given energy, they rise to higher energy levels (excited state).If the total number of electrons matches in the configuration, but the configuration doesn’t match, the atom is in the excited state.Na (ground, on table): 2-8-1Example of excited states: 2-7-2, , 2-6-3
40 Ways to Represent Electron Configuration Expanded Electron ConfigurationCondensed Electron ConfigurationsOrbital NotationElectron Dot StructureWrite the above four electron configurations for Zinc, Zinc ion and Cu ion.
41 Electron Configuration of Ions -Group configurations: s block ns1-2, p block ns2 np 1-6, d block ns0-2 (n-1) d 1-10, f block ns 0-2 (n-1) d 1 (n-2) f 1-14- Remember that outmost electrons are lost first (which means that it will always be s or p electrons lost, never d or f). Ex. Sc+ or Sc3+ electron configuration would be:
42 Rules for Electron Filling - Afbau’s Principle: Electrons tend to occupy the lowest energy orbitals first.- Hund’s Rule: Pairing of e in the degenerate orbitals does not take place till every orbital has one e.- Pauli’s Exclusion Principle: No two electrons can have all four same quantum numbers.
43 Diamagnetism, Paramagnetism Diamagnetism: does not show magnetic properties in external magnetic field. No unpaired electrons.Paramagnetism: shows magnetic properties in external magnetic field. Has unpaired electrons.Best way to predict dia or paramagnetism is by drawing orbital diagrams.
56 Hybridization Refers to mixing of orbitals. Atomic orbitals of central atom undergo change to accommodate incoming atoms.Hybridization could be sp, sp2, sp3, sp3d and sp3d2.How do you tell the hybridization on the central atom?
57 9.1 – 9.2: V.S.E.P.R. Valence-shell electron-pair repulsion theory Because e- pairs repel, molecular shape adjusts so the valence e- pairs are as far apart as possible around the central atom.Electron domains: areas of valence e- density around the central atom; result in different molecular shapesIncludes bonding e- pairs and nonbonding e- pairsA single, double, or triple bond counts as one domainSummary of LmABn (Tables ):L = lone or non-bonding pairsA = central atomB = bonded atomsBond angles notation used here:< xº means ~2-3º less than predicted<< xº means ~4-6º less than predicted
58 Tables 9.1 - 9.3 # of e- domains & # and type of hybrid orbitals e- domain geometryFormulaMolecular geometryPredicted bond angle(s)Example (Lewis structure with molecular shape)2Two sp hybrid orbitalsLinearAB2180ºBeF2CO2|XX|BBAA
67 What Is Light?Light is formed when electrons drop from the excited state to the ground state.The lines on a bright-line spectrum come from specific energy level drops and are unique to each element.Ex. Emission and Absorption Spectra ( line spectra)
68 EXAMPLE SPECTRUM This is the bright-line spectrum of hydrogen. The top numbers represent the energy level transition change that produces thelight with that color and the bottom number is thewavelength of the light (in nanometers, or 10-9 m).No other element has the same bright-line spectrum ashydrogen, so these spectra can be used to identifyelements or mixtures of elements.
69 Development of the Atomic Model Thompson ModelRutherford Gold Foil Experiment and ModelBohr ModelQuantum-Mechanical Model
70 Thompson ModelThe atom is a positively charged diffuse mass with negatively charged electrons stuck in it.
71 Rutherford ModelThe atom is made of a small, dense, positively charged nucleus with electrons at a distance, the vast majority of the volume of the atom is empty space.Alpha particles shotat a thin sheet of goldfoil: most go through(empty space). Somedeflect or bounce off(small + chargednucleus).
72 Bohr ModelElectrons orbit around the nucleus in energy levels (shells). Atomic bright-line spectra was the clue.
73 Quantum-Mechanical Model Electron energy levels are wave functions.Electrons are found in orbitals, regions of space where an electron is most likely to be found.You can’t know both where the electron is and where it is going at the same time.Electrons buzz around the nucleus like gnats buzzing around your head.
74 Orbital Quantum Numbers SymbolNameDescriptionMeaningEquationsnPrinciple Q.N.Energy level(i.e. Bohr’s theory)Shell numbern = 1, 2, 3, 4, 5, 6, 7n = 1, 2, 3, …lAngular Momentum Q.N.General probability plot (“shape” of the orbitals)Subshell numberl = 0, 1, 2, 3l = 0 means “s”l = 1 means “p”l = 2 means “d”l = 3 means “f”l = 0, 1, 2, …, n – 1Ex: If n = 1, l can only be 0; if n = 2, l can be 0 or 1.
75 Symbol Name Description Meaning Equations ml ms Magnetic Q.N.3-D orientation of the orbitals has 1p has 3d has 5f has 7ml = -l, -l +1, …, 0, l, …, +lThere are(2l + 1) values. msSpin Q.N.Spin of the electronParallel or antiparallel to fieldms = +½ or-½* s, p, d, and f come from the words sharp, principal, diffuse, and fundamental.
76 Permissible Quantum Numbers (4, 1, 2, +½)(5, 2, 0, 0)(2, 2, 1, +½)Not permissible; if l = 1, ml = 1, 0, or –1 (p orbitals only have 3 subshells)Not permissible; ms = +½ or –½Not permissible; if n = 2, l = 0 or 1 (there is no 2d orbital)
79 Phase Change Diagrams AB: Solid Phase BC: Melting (S + L) CD: Liquid PhaseDE: Boiling (L + G)EF: Gas PhaseNotice how temperature remains constant during a phase change? That’s because the PE is changing, not the KE.
80 Heat of Phase ChangeHow many joules would it take to melt 100. g of H2O (s) at 0oC?q=mHf = (100. g)(334 J/g) = JHow many joules would it take to boil 100. g of H2O (l) at 100oC?q=mHv = (100.g)(2260 J/g) = J
81 EvaporationWhen the surface molecules of a gas travel upwards at a great enough speed to escape.The pressure a vapor exerts when sealed in a container at equilibrium is called vapor pressure, and can be found on Table H.When the liquid is heated, its vapor pressure increases.When the liquid’s vapor pressure equals the pressure exerted on it by the outside atmosphere, the liquid can boil.If the pressure exerted on a liquid increases, the boiling point of the liquid increases (pressure cooker). If the pressure decreases, the boiling point of the liquid decreases (special cooking directions for high elevations).
82 Reference Table H: Vapor Pressure of Four Liquids (c) 2006, Mark Rosengarten
83 Phase diagrams: CO2 Lines: 2 phases exist in equilibrium Triple point: all 3 phases exist together in equilibrium (X on graph)Critical point, or critical temperature & pressure: highest T and P at which a liquid can exist (Z on graph)Temp (ºC)For most substances, inc P will cause a gas to condense (or deposit), a liquid to freeze, and a solid to become more dense (to a limit.)
84 Phase diagrams: H2OFor H2O, inc P will cause ice to melt.
85 The Periodic Table Metals Nonmetals Metalloids Chemistry of Groups ElectronegativityIonization Energy
86 MetalsHave luster, are malleable and ductile, good conductors of heat and electricityLose electrons to nonmetal atoms to form positively charged ions in ionic bondsLarge atomic radii compared to nonmetal atomsLow electronegativity and ionization energyLeft side of the periodic table (except H)
87 Nonmetals Are dull and brittle, poor conductors Gain electrons from metal atoms to form negatively charged ions in ionic bondsShare unpaired valence electrons with other nonmetal atoms to form covalent bonds and moleculesSmall atomic radii compared to metal atomsHigh electronegativity and ionization energyRight side of the periodic table (except Group 18)
88 MetalloidsFound lying on the jagged line between metals and nonmetals flatly touching the line (except Al and Po).Share properties of metals and nonmetals (Si is shiny like a metal, brittle like a nonmetal and is a semiconductor).
89 Chemistry of Groups Group 1: Alkali Metals Group 2: Alkaline Earth MetalsGroups 3-11: Transition ElementsGroup 17: HalogensGroup 18: Noble GasesDiatomic Molecules
90 Group 1: Alkali MetalsMost active metals, only found in compounds in natureReact violently with water to form hydrogen gas and a strong base: 2 Na (s) + H2O (l) 2 NaOH (aq) + H2 (g)1 valence electronForm +1 ion by losing that valence electronForm oxides like Na2O, Li2O, K2O
91 Group 2: Alkaline Earth Metals Very active metals, only found in compounds in natureReact strongly with water to form hydrogen gas and a base:Ca (s) + 2 H2O (l) Ca(OH)2 (aq) + H2 (g)2 valence electronsForm +2 ion by losing those valence electronsForm oxides like CaO, MgO, BaO
92 Groups 3-11: Transition Metals Many can form different possible charges of ionsIf there is more than one ion listed, give the charge as a Roman numeral after the nameCu+1 = copper (I) Cu+2 = copper (II)Compounds containing these metals can be colored.
93 Group 17: Halogens Most reactive nonmetals React violently with metal atoms to form halide compounds: 2 Na + Cl2 2 NaClOnly found in compounds in natureHave 7 valence electronsGain 1 valence electron from a metal to form -1 ionsShare 1 valence electron with another nonmetal atom to form one covalent bond.
94 Group 18: Noble GasesAre completely nonreactive since they have eight valence electrons, making a stable octet.Kr and Xe can be forced, in the laboratory, to give up some valence electrons to react with fluorine.Since noble gases do not naturally bond to any other elements, one atom of noble gas is considered to be a molecule of noble gas. This is called a monatomic molecule. Ne represents an atom of Ne and a molecule of Ne.
95 Diatomic MoleculesBr, I, N, Cl, H, O and F are so reactive that they exist in a more chemically stable state when they covalently bond with another atom of their own element to make two-atom, or diatomic molecules.Br2, I2, N2, Cl2, H2, O2 and F2The decomposition of water: 2 H2O 2 H2 + O2
96 ElectronegativityAn atom’s attraction to electrons in a chemical bond.F has the highest, at 4.0Fr has the lowest, at 0.7If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion.If the two atoms have an END of less than 1.7, they will share their unpaired valence electrons…covalent bond!
97 Ionization EnergyThe energy required to remove the most loosely held valence electron from an atom in the gas phase.High electronegativity means high ionization energy because if an atom is more attracted to electrons, it will take more energy to remove those electrons.Metals have low ionization energy. They lose electrons easily to form (+) charged ions.Nonmetals have high ionization energy but high electronegativity. They gain electrons easily to form (-) charged ions when reacted with metals, or share unpaired valence electrons with other nonmetal atoms.
98 IonsIons are charged particles formed by the gain or loss of electrons.Metals lose electrons (oxidation) to form (+) charged cations.Nonmetals gain electrons (reduction) to form (-) charged anions.Atoms will gain or lose electrons in such a way that they end up with 8 valence electrons (stable octet).The exceptions to this are H, Li, Be and B, which are not large enough to support 8 valence electrons. They must be satisfied with 2 (Li, Be, B) or 0 (H).
99 Metal Ions (Cations) Na: 2-8-1 Na+1: 2-8 Ca: 2-8-8-2 Ca+2: 2-8-8 Note that when the atom loses its valence electron, the next lower PEL becomes the valence PEL.Notice how the dot diagrams for metal ions lack dots! Place brackets around the element symbol and put the charge on the upper right outside!
100 Nonmetal Ions (Anions) Note how the ions all have 8 valence electrons. Also note the gained electrons as red dots. Nonmetal ion dot diagrams show 8 dots, with brackets around the dot diagram and the charge of the ion written to the upper right side outside the brackets.F: 2-7F-1: 2-8O: 2-6O-2: 2-8N: 2-5N-3: 2-8
101 Chemical BondingIntermolecular Bonding: Ionic, Covalent, Metallic and Covalent Network BondsIntermolecular Bonding: H bond, dipole-dipole interaction, LDFs
102 Ionic BondingIf two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion.The oppositely charged ions attract to form the bond. It is a surface bond that can be broken by melting or dissolving in water.Ionic bonding forms ionic crystal lattices, not molecules.
104 Covalent BondingIf two nonmetal atoms have an END of 1.7 or less, they will share their unpaired valence electrons to form a covalent bond.A particle made of covalently bonded nonmetal atoms is called a molecule.If the END is between 0 and 0.4, the sharing of electrons is equal, so there are no charged ends. This is NONPOLAR covalent bonding.If the END is between 0.5 and 1.7, the sharing of electrons is unequal. The atom with the higher EN will be d- and the one with the lower EN will be d+ charged. This is a POLAR covalent bonding. (d means “partial”)
106 Sigma and Pi bonds Sigma (s) bond: Pi (p) bond: Covalent bond that results from axial overlap of orbitals between atoms in a moleculeLie directly on internuclear axis“Single” bonds, could form between s-s orbital or s-p orbital or p-p orbital by axial overlappingEx: F2Pi (p) bond:Covalent bond that results from side-by-side overlap of orbitals between atoms in a molecule.Are “above & below” and “left & right” of the internuclear axis and therefore have less total orbital overlap, so they are weaker than s bonds. Forms between two p orbitals (py or pz)Make up the 2nd and 3rd bonds in double & triple bonds.Ex: O2 N2
107 Metallic BondingMetal atoms of the same element bond with each other by sharing valence electrons that they lose to each other.This is a lot like an atomic game of “hot potato”, where metal kernals (the atom inside the valence electrons) sit in a crystal lattice, passing valence electrons back and forth between each other).Since electrons can be forced to travel in a certain direction within the metal, metals are very good at conducting electricity in all phases.
108 Types of CompoundsIonic: made of metal and nonmetal ions. Form an ionic crystal lattice when in the solid phase. Ions separate when melted or dissolved in water, allowing electrical conduction. Examples: NaCl, K2O, CaBr2Molecular: made of nonmetal atoms bonded to form a distinct particle called a molecule. Bonds do not break upon melting or dissolving, so molecular substances do not conduct electricity. EXCEPTION: Acids [H+A- (aq)] ionize in water to form H3O+ and A-, so they do conduct.Network: made up of nonmetal atoms bonded in a seemingly endless matrix of covalent bonds with no distinguishable molecules. Very high m.p., don’t conduct.
111 Network SolidsNetwork solids are made of nonmetal atoms covalently bonded together to form large crystal lattices. No individual molecules can be distinguished. Examples include C (diamond) and SiO2 (quartz). Corundum (Al2O3) also forms these, even though Al is considered a metal. Network solids are among the hardest materials known. They have extremely high melting points and do not conduct electricity.
130 Attractive ForcesMolecules have partially charged ends. The d+ end of one molecule attracts to the d- end of another molecule.Ions are charged (+) or (-). Positively charged ions attract other to form ionic bonds, a type of attractive force.Since partially charged ends result in weaker attractions than fully charged ends, ionic compounds generally have much higher melting points than molecular compounds.Determining Polarity of MoleculesHydrogen Bond Attractions
131 Determining Polarity of Molecules (c) 2006, Mark Rosengarten
132 Hydrogen Bond Attractions A hydrogen bond attraction is a very strong attractive force between the H end of one polar molecule and the N, O or F end of another polar molecule. This attraction is so strong that water is a liquid at a temperature where most compounds that are much heavier than water (like propane, C3H8) are gases. This also gives water its surface tension and its ability to form a meniscus in a narrow glass tube.
134 Compounds1) Types of Compounds 2) Formula Writing 3) Formula Naming 4) Empirical Formulas 5) Molecular Formulas 6) Types of Chemical Reactions 7) Balancing Chemical Reactions 8) Attractive Forces
135 Formula WritingThe charge of the (+) ion and the charge of the (-) ion must cancel out to make the formula. Use subscripts to indicate how many atoms of each element there are in the compound, no subscript if there is only one atom of that element.Na+1 and Cl-1 = NaClCa+2 and Br-1 = CaBr2Al+3 and O-2 = Al2O3Zn+2 and PO4-3 = Zn3(PO4)2Try these problems!
136 Formulas to Write Ba+2 and N-3 NH4+1 and SO4-2 Li+1 and S-2 Cu+2 and NO3-1Al+3 and CO3-2Fe+3 and Cl-1Pb+4 and O-2Pb+2 and O-2
137 Formula NamingCompounds are named from the elements or polyatomic ions that form them.KCl = potassium chlorideNa2SO4 = sodium sulfate(NH4)2S = ammonium sulfideAgNO3 = silver nitrateNotice all the metals listed here only have one charge listed? So what do you do if a metal has more than one charge listed? Take a peek!
138 The Stock System CrCl2 = chromium (II) chloride Try CrCl3 = chromium (III) chloride Co(NO3)2 andCrCl6 = chromium (VI) chloride Co(NO3)3FeO = iron (II) oxide MnS = manganese (II) sulfideFe2O3 = iron (III) oxide MnS2 = manganese (IV) sulfideThe Roman numeral is the charge of the metal ion!
139 Math of Chemistry1) Formula Mass 2) Percent Composition 3) Mole Problems 4) Gas Laws 5) Neutralization 6) Concentration 7) Significant Figures and Rounding 8) Metric Conversions 9) Calorimetry
140 Formula MassGram Formula Mass = sum of atomic masses of all elements in the compoundRound given atomic masses to the nearest tenthH2O: (2 X 1.0) + (1 X 16.0) = 18.0 grams/moleNa2SO4: (2 X 23.0)+(1 X 32.1)+(4 X 16.0) = g/moleNow you try:BaBr2CaSO4Al2(CO3)3
141 Percent CompositionThe mass of part is the number of atoms of that element in the compound. The mass of whole is the formula mass of the compound. Don’t forget to take atomic mass to the nearest tenth! This is a problem for you to try.
142 Practice Percent Composition Problem What is the percent by mass of each element in Li2SO4?
144 Grams <=> MolesHow many grams will 3.00 moles of NaOH (40.0 g/mol) weigh?3.00 moles X 40.0 g/mol = 120. gHow many moles of NaOH (40.0 g/mol) are represented by 10.0 grams?(10.0 g) / (40.0 g/mol) = mol
145 Molecular FormulaMolecular Formula = (Molecular Mass/Empirical Mass) X Empirical FormulaWhat is the molecular formula of a compound with an empirical formula of CH2 and a molecular mass of 70.0 grams/mole?1) Find the Empirical Formula Mass: CH2 = 14.02) Divide the MM/EM: 70.0/14.0 = 53) Multiply the molecular formula by the result:5 (CH2) = C5H10
146 StoichiometryMoles of Target = Moles of Given X (Coefficent of Target/Coefficient of given)Given the balanced equation N2 + 3 H2 2 NH3, How many moles of H2 need to be completely reacted with N2 to yield 20.0 moles of NH3?20.0 moles NH3 X (3 H2 / 2 NH3) = 30.0 moles H2
147 Limiting Reactant controls the amount of product formed. CO(g) + 2H2 (g) Ch3OHIf 500 mol of CO react with 750 mol of H2, which is the limiting reactant?Use either given amount to calculate required amount of other.Compare calculated amount to amount givenb. How many moles of excess reactant remain unchanged?H2125 mol CO
148 Percent yield= (actual yield/ theoretical yield)*100 Theoretical yield is the maximum amount of product that can be produced from a given amount of reactantActual yield is the measured amount of a product obtained from a reactionTheoretical yield= g SnF2Actual yield = g SnF2Percent yield = g SnF2117.5 g SnF2*100
149 Determining empirical formula from combustion data When a compound containing C,H and O undergoes combustion, it forms CO2 and H2O. Then from the mass of CO2 and H2O, we can calculate the mass of C and Hand then find the mass of O by subtracting the sum of masses of C and H from total g present of that substance. From the mass of C,H and O, we can calculate the moles of C,H and O.Then the smallest whole number ratios of these moles will give the empirical formula.Ex. A g sample of the unknown produced g of CO2 and g of H2O. Determine the empirical formula of the compound. Ans. C7H6O2
150 Empirical FormulasIonic formulas: represent the simplest whole number mole ratio of elements in a compound.Ca3N2 means a 3:2 ratio of Ca ions to N ions in the compound.Many molecular formulas can be simplified to empirical formulasEthane (C2H6) can be simplified to CH3. This is the empirical formula…the ratio of C to H in the molecule.All ionic compounds have empirical formulas.
151 Molecular FormulasThe count of the actual number of atoms of each element in a molecule.H2O: a molecule made of two H atoms and one O atom covalently bonded together.C2H6O: A molecule made of two C atoms, six H atoms and one O atom covalently bonded together.Molecular formulas are whole-number multiples of empirical formulas:H2O = 1 X (H2O)C8H16 = 8 X (CH2)Calculating Molecular Formulas
161 Molecular FormulaActual ratio of atoms in a compound.Ex. H2O, C6H12O6To determine the molecular formula, divide the molar mass by empirical formula mass. This will give the number of empirical formula units (n) in actual molecule.n= Molar Mass/ Empirical Formula MassEx. Determine the empirical and molecular formula of each of the following:Ethylene glycol, the substance used as antifreeze has % C, 9.70 % H and % O , mm= gCaffeine, a stimulant in coffee has the following percent composition:49.50 % C, 5.15% H, % N and % O , molar mass= g
162 Types of Chemical Reactions Redox Reactions: driven by the loss (oxidation) and gain (reduction) of electrons. Any species that does not change charge is called the spectator ion.SynthesisDecompositionSingle ReplacementIon Exchange Reaction: driven by the formation of an insoluble precipitate. The ions that remain dissolved throughout are the spectator ions.Double Replacement
163 Synthesis Two elements combine to form a compound 2 Na + O2 Na2O Same reaction, with charges added in:2 Na0 + O20 Na2+1O-2Na0 is oxidized (loses electrons), is the reducing agentO20 is reduced (gains electrons), is the oxidizing agentElectrons are transferred from the Na0 to the O20.No spectator ions, there are only two elements here.
164 Decomposition A compound breaks down into its original elements. Na2O 2 Na + O2Same reaction, with charges added in:Na2+1O-2 2 Na0 + O20O-2 is oxidized (loses electrons), is the reducing agentNa+1 is reduced (gains electrons), is the oxidizing agentElectrons are transferred from the O-2 to the Na+1.No spectator ions, there are only two elements here.
165 Single ReplacementAn element replaces the same type of element in a compound.Ca + 2 KCl CaCl2 + 2 KSame reaction, with charges added in:Ca0 + 2 K+1Cl-1 Ca+2Cl K0Ca0 is oxidized (loses electrons), is the reducing agentK+1 is reduced (gains electrons), is the oxidizing agentElectrons are transferred from the Ca0 to the K+1.Cl-1 is the spectator ion, since it’s charge doesn’t change.
166 Double ReplacementThe (+) ion of one compound bonds to the (-) ion of another compound to make an insoluble precipitate. The compounds must both be dissolved in water to break the ionic bonds first.NaCl (aq) + AgNO3 (aq) NaNO3 (aq) + AgCl (s)The Cl-1 and Ag+1 come together to make the insoluble precipitate, which looks like snow in the test tube.No species change charge, so this is not a redox reaction.Since the Na+1 and NO3-1 ions remain dissolved throughout the reaction, they are the spectator ions.How do identify the precipitate?
167 Identifying the Precipitate The precipitate is the compound that is insoluble. AgCl is a precipitate because Cl- is a halide. Halides are soluble, except when combined with Ag+ and others.
168 Balancing Chemical Reactions Balance one element or ion at a timeUse a pencilUse coefficients only, never change formulasRevise if necessaryThe coefficient multiplies everything in the formula by that amount2 Ca(NO3)2 means that you have 2 Ca, 4 N and 12 O.Examples for you to try!
172 Manometers: measure P of a gas Closed-end: difference in Hg levels (Dh) shows P of gas in container compared to a vacuumclosed
173 2. Open-end:Difference in Hg levels (Dh) shows P of gas in container compared to Patm
174 Gas LawsMake a data table to put the numbers so you can eliminate the words.Make sure that any Celsius temperatures are converted to Kelvin (add 273).Rearrange the equation before substituting in numbers. If you are trying to solve for T2, get it out of the denominator first by cross-multiplying.If one of the variables is constant, then eliminate it.Try these problems!
175 Gas Law Problem 1A 2.00 L sample of N2 gas at STP is compressed to 4.00 atm at constant temp-erature. What is the new volume of the gas?V2 = P1V1 / P2= (1.00 atm)(2.00 L) / (4.00 atm)= L
176 Gas Law Problem 2To what temperature must a L sample of O2 gas at K be heated to raise the volume to L?T2 = V2T1/V1= (10.00 L)(300.0 K) / (3.000 L) = K
177 Gas Law Problem 3A 3.00 L sample of NH3 gas at kPa is cooled from K to K and its pressure is reduced to 80.0 kPa. What is the new volume of the gas?V2 = P1V1T2 / P2T1= (100.0 kPa)(3.00 L)(300. K) / (80.0 kPa)(500. K)= 2.25 L
178 Gay Lussac’s Law of Combining Volumes When measured at the same temperature and pressure, the ratio of the volumes of reacting gases are small whole numbers.N21 volume+3 H23 volumes→2 NH32 volumes1. Students write balanced equation for reaction. Note balancing coefficients are same as volume in Liters.
179 Gay Lussac’s Law of Combining Volumes When measured at the same temperature and pressure, the ratio of the volumes of reacting gases are small whole numbers.
180 Avogadro’s LawEqual volumes of different gases at the same temperature and pressure contain the same number of molecules.
181 22.4 L at STP is known as the molar volume of any gas. Mole-Mass-Volume RelationshipsVolume of one mole of any gas at STP = 22.4 L.22.4 L at STP is known as the molar volume of any gas.
182 atmospheresnTV aPNumber of molecules or moles affects other three quatities. Increase #, increase collision rate, increase pressure.
183 Determination of Density Using the Ideal Gas Equation Density = mass/volumeDensity varies directly with molar mass and pressure and inversely with Kelvin tempD = MP/ RT
184 Mole fraction (X):Ratio of moles of one component to the total moles in the mixture (dimensionless, similar to a %)Ex: What are the mole fractions of H2 and He in the previous example?
185 Collecting Gases “over Water” When a gas is bubbled through water, the vapor pressure of the water (partial pressure of the water) must be subtracted from the pressure of the collected gas:PT = Pgas + PH2O∴ Pgas = PT – PH2OSee Appendix B for vapor pressures of water at different temperatures.
186 Graham’s Law of Effusion The rates of effusion of gases at the same temperature and pressure are inversely proportional to the square roots of their molar masses.Rate of effusion of A = MBRate of effusion of B MAM = molar masses, density of a gas varies directly with molar mass, can replace molar mass with density
187 Neutralization10.0 mL of 0.20 M HCl is neutralized by 40.0 mL of NaOH. What is the concentration of the NaOH?#H MaVa = #OH MbVb, so Mb = #H MaVa / #OH Vb= (1)(0.20 M)(10.0 mL) / (1) (40.0 mL) = MHow many mL of 2.00 M H2SO4 are needed to completely neutralize 30.0 mL of M KOH?
188 Concentration Molarity Parts per Million Percent by Mass Percent by Volume
189 MolarityWhat is the molarity of a mL solution of NaOH (FM = 40.0) with 60.0 g of NaOH (aq)?Convert g to moles and mL to L first!M = moles / L = 1.50 moles / L = 3.00 MHow many grams of NaOH does it take to make 2.0 L of a M solution of NaOH (aq)?Moles = M X L = M X 2.0 L = molesConvert moles to grams: moles X 40.0 g/mol = 8.00 g
190 Parts Per Million100.0 grams of water is evaporated and analyzed for lead grams of lead ions are found. What is the concentration of the lead, in parts per million?ppm = ( g) / (100.0 g) X = 1.0 ppmIf the legal limit for lead in the water is 3.0 ppm, then the water sample is within the legal limits (it’s OK!)
191 Percent by MassA 50.0 gram sample of a solution is evaporated and found to contain grams of sodium chloride. What is the percent by mass of sodium chloride in the solution?% Comp = (0.100 g) / (50.0 g) X 100 = 0.200%
192 Percent By VolumeSubstitute “volume” for “mass” in the above equation.What is the percent by volume of hexane if 20.0 mL of hexane are dissolved in benzene to a total volume of 80.0 mL?% Comp = (20.0 mL) / (80.0 mL) X100 = 25.0%
193 Colligative Properties Vapor Pressure LoweringB.P. Elevation DTf= m. kf (or D Tb= m. kb)F.P. DepressionOsmotic PressureColligative properties depend upon # of particles (ions, atoms, molecule= particle)Which will have lowest B.P. 1M NaCl, 1 M C6H12O6 or 1M Na3PO4?
194 How many Sig Figs? Start counting sig figs at the first non-zero. All digits except place-holding zeroes are sig figs.Measurement# of Sig Figs0.115 cm3cm2cmcmcm5Measurement# of Sig Figs234 cm367000 cm2_45000 cm4560. cmcm5
195 What Precision?A number’s precision is determined by the furthest (smallest) place the number is recorded to.6000 mL : thousands place6000. mL : ones placemL : tenths place5.30 mL : hundredths place8.7 mL : tenths placemL : thousandths place
196 Rounding with addition and subtraction Answers are rounded to the least precise place.
197 Rounding with multiplication and division Answers are rounded to the fewest number of significant figures.
198 Metric ConversionsDetermine how many powers of ten difference there are between the two units (no prefix = 100) and create a conversion factor. Multiply or divide the given by the conversion factor.How many kg are in 38.2 cg?(38.2 cg) /( cg/kg) = kmHow many mL in dL?(0.988 dg) X (100 mL/dL) = 98.8 mL
199 CalorimetryThis equation can be used to determine any of the variables here. You will not have to solve for C, since we will always assume that the energy transfer is being absorbed by or released by a measured quantity of water, whose specific heat is given above.Solving for qSolving for mSolving for DT
200 Solving for qHow many joules are absorbed by grams of water in a calorimeter if the temperature of the water increases from 20.0oC to 50.0oC?q = mCDT = (100.0 g)(4.18 J/goC)(30.0oC) = J
201 Solving for mA sample of water in a calorimeter cup increases from 25oC to 50.oC by the addition of joules of energy. What is the mass of water in the calorimeter cup?q = mCDT, so m = q / CDT = (500.0 J) / (4.18 J/goC)(25oC) = 4.8 g
202 Solving for DTIf a 50.0 gram sample of water in a calorimeter cup absorbs joules of energy, how much will the temperature rise by?q = mCDT, so DT = q / mC = ( J)/(50.0 g)(4.18 J/goC) = 4.8oCIf the water started at 20.0oC, what will the final temperature be?Since the water ABSORBS the energy, its temperature will INCREASE by the DT: 20.0oC + 4.8oC = 24.8oC
203 Reaction RateReactions happen when reacting particles collide with sufficient energy (activation energy) and at the proper angle.Anything that makes more collisions in a given time will make the reaction rate increase.Increasing temperatureIncreasing concentration (pressure for gases)Increasing surface area (solids)Adding a catalyst makes a reaction go faster by removing steps from the mechanism and lowering the activation energy without getting used up in the process.
204 Heat of ReactionReactions either absorb PE (endothermic, +DH) or release PE (exothermic, -DH)Exothermic, PEKE, TempEndothermic, KEPE, TempRewriting the equation with heat included:4 Al(s) + 3 O2(g) 2 Al2O3(s) kJN2(g) + O2(g) kJ 2 NO(g)
205 5.3: Enthalpy, HSince most reactions occur in containers open to the air, w is often negligible. If a reaction produces a gas, the gas must do work to expand against the atmosphere. This mechanical work of expansion is called PV (pressure-volume) work.Enthalpy (H): change in the heat content (qp) of a reaction at constant pressureH = E + PVH = E + PV (at constant P)H = (qp + w) + (-w)H = qp
206 Sign conventionsH > 0 Heat is gained from surroundings + H in endothermic reactionH < 0 Heat is released to surroundings - H in exothermic reaction
207 5.4: Enthalpy of Reaction (Hrxn) Also called heat of reaction:Enthalpy is an extensive property (depends on amounts of reactants involved).Ex: CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (l)Hrxn = kJCombustion of 1 mol CH4 produces 890. kJ… of 2 mol CH4 → (2)(-890. kJ) = kJWhat is the H of the combustion of 100. g CH4?
209 5.6: Hess’ Law Ex. What is DHrxn of the combustion of propane? If a rxn is carried out in a series of steps,Hrxn = (Hsteps) = H1 + H2 + H3 + …Germain Hess ( )Ex. What is DHrxn of the combustion of propane?C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l)3 C (s) + 4 H2 (g) C3 H8 (g) H1 = kJC (s) + O2 (g) CO2 (g) H2 = kJH2 (g) + ½ O2 (g) H2O (l) H3 = kJC3H8 (g) 3 C (s) + 4 H2 (g) H1 = kJBorn in Switzerland, moved to St Petersburg at age 3 with his family. His father was an artist. Hess was also interested in geology.3[ ] ( )4[ ] ( )Hrxn = ( ) + 4( ) = kJ
210 5.7: Enthalpy of Formation (Hf) Formation: a reaction that describes a substance formed from its elementsNH4NO3 (s)Standard enthalpy of formation (Hf): forms 1 mole of compound from its elements in their standard state (at 298 K)C2H5OH (l)Hf = kJHf of the most stable form of any element equals zero. H2, N2 , O2 , F2 , Cl2 (g)Br2 (l), Hg (l)C (graphite), P4 (s, white), S8 (s), I2 (s)Ex: 2 N2 (g) + 4 H2 (g) + 3 O2 (g) 22 C (graphite) + 3 H2 (g) + ½ O2 (g)
212 5.5: Calorimetry q = C DT q = m c DT Measurement of heat flow Heat capacity, C: amount of heat required to raise T of an object by 1 Kq = C DTSpecific heat (or specific heat capacity, c): heat capacity of 1 g of a substanceq = m c DTEx: How much energy is required to heat 40.0 g of iron (c = 0.45 J/(g K) from 0.0ºC to 100.0ºC?q = m c DT = (40.0 g)(0.45 J/(g K))(100.0 – 0.0 ºC)= 1800 J
213 Potential Energy Diagrams Steps of a reactions:Reactants have a certain amount of PE stored in their bonds (Heat of Reactants)The reactants are given enough energy to collide and react (Activation Energy)The resulting intermediate has the highest energy that the reaction can make (Heat of Activated Complex)The activated complex breaks down and forms the products, which have a certain amount of PE stored in their bonds (Heat of Products)Hproducts - Hreactants = DH EXAMPLES
214 Making a PE Diagram X axis: Reaction Coordinate (time, no units) Y axis: PE (kJ)Three lines representing energy (Hreactants, Hactivated complex, Hproducts)Two arrows representing energy changes:From Hreactants to Hactivated complex: Activation EnergyFrom Hreactants to Hproducts : DHENDOTHERMIC PE DIAGRAMEXOTHERMIC PE DIAGRAM
215 Endothermic PE Diagram If a catalyst is added?
216 Endothermic with Catalyst The red line represents the catalyzed reaction.
217 Exothermic PE DiagramWhat does it look like with a catalyst?
218 Exothermic with a Catalyst The red line represents the catalyzed reaction. Lower A.E. and faster reaction time!
219 19.1: Spontaneous Processes Reversible reaction: can proceed forward and backward along same path (equilibrium is possible)Ex: H2O freezing & melting at 0ºCIrreversible reaction: cannot proceed forward and backward along same pathEx: ice melting at room temperatureSpontaneous reaction: an irreversible reaction that occurs without outside interventionEx: Gases expand to fill a container, ice melts at room temperature (even though endothermic), salts dissolve in water
220 Entropy Entropy (S): a measure of molecular randomness or disorder S is a state function: DS = Sfinal - Sinitial+ DS = more randomness- DS = less randomnessFor a reversible process that occurs at constant T:Units: J/K
221 Examples of spontaneous reactions: Particles are more evenly distributedParticles are no longer in an ordered crystal latticeIons are not locked in crystal latticeGases expand to fill a container:Ice melts at room temperature:Salts dissolve in water:
222 19.3: 3rd Law of Thermodynamics The entropy of a crystalline solid at 0 K is 0.How to predict DS:Sgas > Sliquid > SsolidSmore gas molecules > Sfewer gas moleculesShigh T > Slow TEx: Predict the sign of DS for the following:CaCO3 (s) → CaO (s) + CO2 (g)N2 (g) + 3 H2 (g) → 2 NH3 (g)N2 (g) + O2 (g) → 2 NO (g)+, solid to gas-, fewer moles produced?
223 DG = DH - TDS DG° = DH° - TDS° 19.5: Gibbs free energy, GRepresents combination of two forces that drive a reaction: DH (enthalpy) and DS (disorder)Units: kJ/molDG = DH - TDS DG° = DH° - TDS°(absolute T)Josiah Willard Gibbs ( )Called “free energy” because DG represents maximum useful work that can be done by the system on its surroundings in a spontaneous reaction. (See p. 708 for more details.)
224 Determining Spontaneity of a Reaction If DG is: reaction is spontaneous (proceeds in the forward directionPositiveForward reaction is non-spontaneous; the reverse reaction is spontaneousZero The system is at equilibrium
225 19.6: Free Energy & Temperature DG depends on enthalpy, entropy, and temperature:DG = DH - TDSDH DS DG and reaction outcome- + Always (-); spontaneous at all T2 O3 (g) → 3 O2 (g)+ - Always +; non-spontaneous at all T3 O2 (g) → 2 O3 (g)- - Spontaneous at low T; non-spontaneous at high TH2O (l) → H2O (s)+ + Spontaneous only at high T ; non-spontaneous at low TH2O (s) → H2O (l)
226 Solubility CurvesSolubility: the maximum quantity of solute that can be dissolved in a given quantity of solvent at a given temperature to make a saturated solution.Saturated: a solution containing the maximum quantity of solute that the solvent can hold. The limit of solubility.Supersaturated: the solution is holding more than it can theoretically hold OR there is excess solute which precipitates out. True supersaturation is rare.Unsaturated: There are still solvent molecules available to dissolve more solute, so more can dissolve.How ionic solutes dissolve in water: polar water molecules attach to the ions and tear them off the crystal.
227 SolubilitySolubility: go to the temperature and up to the desired line, then across to the Y-axis. This is how many g of solute are needed to make a saturated solution of that solute in 100g of H2O at that particular temperature.At 40oC, the solubility of KNO3 in 100g of water is 64 g. In 200g of water, double that amount. In 50g of water, cut it in half.
228 Supersaturated If 120 g of NaNO3 are added to 100g of water at 30oC: 1) The solution would be SUPERSATURATED, because there is more solute dissolved than the solubility allows2) The extra 25g would precipitate out3) If you heated the solution up by 24oC (to 54oC), the excess solute would dissolve.
229 Unsaturated If 80 g of KNO3 are added to 100g of water at 60oC: 1) The solution would be UNSATURATED, because there is less solute dissolved than the solubility allows2) 26g more can be added to make a saturated solution3) If you cooled the solution down by 12oC (to 48oC), the solution would become saturated
230 How Ionic Solutes Dissolve in Water Water solvent molecules attach to the ions (H end to the Cl-, O end to the Na+)Water solvent holds the ions apart and keeps the ions from coming back together
231 Formulas, Naming and Properties of Acids Arrhenius Definition of Acids: molecules that dissolve in water to produce H3O+ (hydronium) as the only positively charged ion in solution.HCl (g) + H2O (l) H3O+ (aq) + Cl-Properties of AcidsNaming of AcidsFormula Writing of Acids
232 Properties of AcidsAcids react with metals above H2 on Table J to form H2(g) and a salt.Acids have a pH of less than 7.Dilute solutions of acids taste sour.Acids turn phenolphthalein CLEAR, litmus RED and bromthymol blue YELLOW.Acids neutralize bases.Acids are formed when acid anhydrides (NO2, SO2, CO2) react with water for form acids. This is how acid rain forms from auto and industrial emissions.
233 Naming of Acids (polyatomic ion) -ate +ic acid Binary Acids (H+ and a nonmetal)hydro (nonmetal) -ide + ic acidHCl (aq) = hydrochloric acidTernary Acids (H+ and a polyatomic ion)(polyatomic ion) -ate +ic acidHNO3 (aq) = nitric acid(polyatomic ion) -ide +ic acidHCN (aq) = cyanic acid(polyatomic ion) -ite +ous acidHNO2 (aq) = nitrous acid
234 Formula Writing of Acids Acids formulas get written like any other. Write the H+1 first, then figure out what the negative ion is based on the name. Cancel out the charges to write the formula. Don’t forget the (aq) after it…it’s only an acid if it’s in water!Hydrosulfuric acid: H+1 and S-2 = H2S (aq)Carbonic acid: H+1 and CO3-2 = H2CO3 (aq)Chlorous acid: H+1 and ClO2-1 = HClO2 (aq)Hydrobromic acid: H+1 and Br-1 = HBr (aq)Hydronitric acid:Hypochlorous acid:Perchloric acid:
235 Formulas, Naming and Properties of Bases Arrhenius Definition of Bases: ionic compounds that dissolve in water to produce OH- (hydroxide) as the only negatively charged ion in solution.NaOH (s) Na+1 (aq) + OH-1 (aq)Properties of BasesNaming of BasesFormula Writing of Bases
236 Properties of BasesBases react with fats to form soap and glycerol. This process is called saponification.Bases have a pH of more than 7.Dilute solutions of bases taste bitter.Bases turn phenolphthalein PINK, litmus BLUE and bromthymol blue BLUE.Bases neutralize acids.Bases are formed when alkali metals or alkaline earth metals react with water. The words “alkali” and “alkaline” mean “basic”, as opposed to “acidic”.
237 Naming of BasesBases are named like any ionic compound, the name of the metal ion first (with a Roman numeral if necessary) followed by “hydroxide”.Fe(OH)2 (aq) = iron (II) hydroxideFe(OH)3 (aq) = iron (III) hydroxideAl(OH)3 (aq) = aluminum hydroxideNH3 (aq) is the same thing as NH4OH:NH3 + H2O NH4OHAlso called ammonium hydroxide.
238 Formula Writing of Bases Formula writing of bases is the same as for any ionic formula writing. The charges of the ions have to cancel out.Calcium hydroxide = Ca+2 and OH-1 = Ca(OH)2 (aq)Potassium hydroxide = K+1 and OH-1 = KOH (aq)Lead (II) hydroxide = Pb+2 and OH-1 = Pb(OH)2 (aq)Lead (IV) hydroxide = Pb+4 and OH-1 = Pb(OH)4 (aq)Lithium hydroxide =Copper (II) hydroxide =Magnesium hydroxide =
240 pH A change of 1 in pH is a tenfold increase in acid or base strength. A pH of 4 is 10 times more acidic than a pH of 5.A pH of 12 is 100 times more basic than a pH of 10.
241 16.2: Dissociation of Water Autoionization of water:H2O (l) ↔ H+ (aq) + OH- (aq)KW = ion-product constant for waterH3O+ (aq) or H+ (aq) = hydronium
242 Indicators At a pH of 2: Methyl Orange = red Bromthymol Blue = yellow Phenolphthalein = colorlessLitmus = redBromcresol Green = yellowThymol Blue = yellowMethyl orange is red at a pH of 3.2 and below and yellow at a pH of 4.4 and higher. In between the two numbers, it is an intermediate color that is not listed on this table.
243 Alternate TheoriesArrhenius Theory: acids and bases must be in aqueous solution.Alternate Theory: Not necessarily so!Acid: proton (H+1) donor…gives up H+1 in a reaction.Base: proton (H+1) acceptor…gains H+1 in a reaction.HNO3 + H2O H3O+1 + NO3-1Since HNO3 lost an H+1 during the reaction, it is an acid.Since H2O gained the H+1 that HNO3 lost, it is a base.
244 16.11: Lewis Acids & Bases Lewis acid: “e- pair acceptor” Brønsted-Lowry acid = H+ donorArrhenius acid = produces H+Lewis base: “e- pair donor”B-L base = H+ acceptorArrhenius base = produces OH-Ex:NH3 + BF3 → NH3BF3Lewis base Lewis acid Lewis salt6 CN- + Fe3+ → Fe(CN)63-Lewis base Lewis acid Coordination compoundGilbert N. Lewis (1875 – 1946)Picture of Gilbert N Lewis (USA) who was the first to isolate D2O.
245 15.1: Chemical Equilibrium Occurs when opposing reactions are proceeding at the same rateForward rate = reverse rate of reactionEx:Vapor pressure: rate of vaporization = rate of condensationSaturated solution: rate of dissociation = rate of crystallizationExpressing concentrations:Gases: partial pressures, PXSolutes in liquids: molarity, [X]
246 Reversible Reactions and Rate Forward rateReaction RateTimeEquilibrium is established:Forward rate = Backward rateBackward rateWhen equilibrium is achieved:[A] ≠ [B] and kf/kr = Keq
247 15.2: Law of Mass ActionDerived from rate laws by Guldberg and Waage (1864)For a balanced chemical reaction in equilibrium:a A + b B ↔ c C + d DEquilibrium constant expression (Keq):Cato Guldberg Peter Waage ( ) ( )orBut Waage and Guldberg were also related through two marriages; Guldberg married his cousin Bodil Mathea Riddervold, daughter of cabinet minister Hans Riddervold, and the couple had three daughters. Waage married Bodil's sister, Johanne Christiane Tandberg Riddervold by whom he had five children, and after her death in 1869, he became Guldberg's brother-in-law a second time, in 1870, by marrying one of Guldberg's sisters, Mathilde Sofie Guldberg, by whom he had six children.Keq is strictly based on stoichiometry of the reaction (is independent of the mechanism).Units: Keq is considered dimensionless (no units)
248 Relating Kc and Kp Convert [A] into PA: where Dn = = change in coefficents of products – reactants (gases only!) = (c+d) - (a+b)
249 Magnitude of KeqSince Keq a [products]/[reactants], the magnitude of Keq predicts which reaction direction is favored:If Keq > 1 then [products] > [reactants] and equilibrium “lies to the right”If Keq < 1 then [products] < [reactants] and equilibrium “lies to the left”
250 Relationship Between Q and K Reaction Quotient (Q): The particular ratio of concentration terms that we write for a particular reaction is called reaction quotient.For a reaction, A B, Q= [B]/[A]At equilibrium, Q= KReaction Direction: Comparing Q and KQ<K, reaction proceeds to right, until equilibrium is achieved (or Q=K)Q>K, reaction proceeds to left, until Q=K
251 Value of K For the reference rxn, A>B, For the reverse rxn, B >A,For the reaction,2A > 2BFor the rxn,A > CC > BK(ref)= [B]/[A]K= 1/K(ref)K= K(ref)2K (overall)= K1 X K2
252 15.3: Types of EquilibriaHomogeneous: all components in same phase (usually g or aq)N2 (g) + H2 (g) ↔ NH3 (g)132Fritz Haber (1868 – 1934)German chemist, who received the Nobel Prize in Chemistry in 1918 for his development of synthetic ammonia, important for fertilizers and explosives. He is also credited as the "father of chemical warfare" for his work developing and deploying chlorine and other poison gases during World War I; this role is thought to have provoked his wife to commit suicide. Despite his contributions to the German war effort, Haber was forced to emigrate from Germany in 1933 by the Nazis because of his Jewish background; many of his relatives were killed by the Nazis in concentration camps, gassed by Zyklon B. Though he had converted from Judaism in an effort to become fully accepted, he was forced to emigrate from Germany by the Nazis in 1933 on account of his being Jewish in their eyes. He died in the process of emigration.The Haber process now produces 500 million tons of nitrogen fertilizer per year, mostly in the form of anhydrous ammonia, ammonium nitrate, and urea. 1% of the world's annual energy supply is consumed in the Haber process (Science 297(1654), Sep 2002). That fertilizer is responsible for sustaining 40% of the Earth's population, as well as various deleterious environmental consequences.
253 CaCO3 (s) ↔ CaO (s) + CO2 (g) Heterogeneous: different phasesCaCO3 (s) ↔ CaO (s) + CO2 (g)Definition: What we use:Concentrations of pure solids and pure liquids are not included in Keq expression because their concentrations do not vary, and are “already included” in Keq (see p. 548).Even though the concentrations of the solids or liquids do not appear in the equilibrium expression, the substances must be present to achieve equilibrium.
254 15.4: Calculating Equilibrium Constants Steps to use “ICE” table:“I” = Tabulate known initial and equilibrium concentrations of all species in equilibrium expression“C” = Determine the concentration change for the species where initial and equilibrium are knownUse stoichiometry to calculate concentration changes for all other species involved in equilibrium“E” = Calculate the equilibrium concentrations
255 NH3 (aq) + H2O (l) ↔ NH41+ (aq) + OH1- (aq) Ex: Enough ammonia is dissolved in 5.00 L of water at 25ºC to produce a solution that is M ammonia. The solution is then allowed to come to equilibrium. Analysis of the equilibrium mixture shows that [OH1-] is 4.64 x 10-4 M. Calculate Keq at 25ºC for the reaction:NH3 (aq) + H2O (l) ↔ NH41+ (aq) + OH1- (aq)
256 NH3 (aq) + H2O (l) ↔ NH41+ (aq) + OH1- (aq) InitialChangeEquilibriumNH3 (aq)H2O (l)NH41+ (aq)OH1- (aq)XM0 M0 MX- x+ x+ xXM4.64 x 10-4 M4.64 x 10-4 Mx = 4.64 x 10-4 M
257 EquilibriumWhen the rate of the forward reaction equals the rate of the reverse reaction.(c) 2006, Mark Rosengarten
258 Examples of Equilibrium Solution Equilibrium: when a solution is saturated, the rate of dissolving equals the rate of precipitating.NaCl (s) Na+1 (aq) + Cl-1 (aq)Vapor-Liquid Equilibrium: when a liquid is trapped with air in a container, the liquid evaporates until the rate of evaporation equals the rate of condensation.H2O (l) H2O (g)Phase equilibrium: At the melting point, the rate of solid turning to liquid equals the rate of liquid turning back to solid.H2O (s) H2O (l)
259 Le Châtelier’s Principle If a system at equilibrium is stressed, the equilibrium will shift in a direction that relieves that stress.A stress is a factor that affects reaction rate. Since catalysts affect both reaction rates equally, catalysts have no effect on a system already at equilibrium.Equilibrium will shift AWAY from what is addedEquilibrium will shift TOWARDS what is removed.This is because the shift will even out the change in reaction rate and bring the system back to equilibriumNEXT
260 Steps to Relieving Stress 1) Equilibrium is subjected to a STRESS.2) System SHIFTS towards what is removed from the system or away from what is added.The shift results in a CHANGE OF CONCENTRATION for both the products and the reactants.If the shift is towards the products, the concentration of the products will increase and the concentration of the reactants will decrease.If the shift is towards the reactants, the concentration of the reactants will increase and the concentration of the products will decrease.NEXT
261 Examples For the reaction N2(g) + 3H2(g) 2 NH3(g) + heat Adding N2 will cause the equilibrium to shift RIGHT, resulting in an increase in the concentration of NH3 and a decrease in the concentration of N2 and H2.Removing H2 will cause a shift to the LEFT, resulting in a decrease in the concentration of NH3 and an increase in the concentration of N2 and H2.Increasing the temperature will cause a shift to the LEFT, same results as the one above.Decreasing the pressure will cause a shift to the LEFT, because there is more gas on the left side, and making more gas will bring the pressure back up to its equilibrium amount.Adding a catalyst will have no effect, so no shift will happen.
262 Oxidation Numbers Rules for Assigning Oxidation States The oxidation state of an atom in an uncombined element is 0.The oxidation state of a monatomic ion is the same as its charge.Oxygen is assigned an oxidation state of –2 in most of its covalent compounds. Important exception: peroxides (compounds containing the O2 2- group), in which each oxygen is assigned an oxidation state of –1)In its covalent compounds with nonmetals, hydrogen is assigned an oxidation state of +1For a compound, sum total of ON s is zero.For an ionic species (like a polyatomic ion), the sum of the oxidation states must equal the overall charge on that ion.
263 16.6: Weak AcidsWeak acids partially ionize in water (equilibrium is somewhere between ions and molecules). HA (aq) ↔ A- (aq) + H+ (aq)Ka = acid-dissociation constant in waterWeak acids generally have Ka < 10-3See Appendix D for full listing of Ka values
290 16.9: Salt Solutions as Acids & Bases Hydrolysis: acid/base reaction of ion with water to produce H+ or OH-Anion (A-) = a conjugate baseA- (aq) + H2O (l) ↔ HA (aq) + OH- (aq)Cation (B+) = a conjugate acidB+ (aq) + H2O (l) ↔ BOH (aq) + H+ (aq)
291 17.1: Common Ion EffectAddition of a “common ion”: solubility of solids decrease because of Le Châtelier’s principle.Ex: AgCl (s) ↔ Ag+ (aq) + Cl- (aq)Addition of Cl- shifts equilibrium toward solid
292 17.4: Solubility Equilibria Dissolving & precipitating of saltsSolubility rules discussed earlier are generalized qualitative observations of quantitative experiments.Ex: PbCl2 (s) ↔ Pb2+ (aq) + 2 Cl- (aq)Ksp = [Pb2+][Cl-]2 = 1.6 x 10-5Ksp = solubility-product constant (found in App. D)Recall that both aqueous ions and solid must be present in solution to achieve equilibriumChanges in pH will affect the solubility of salts composed of a weak acid or weak base ion.
301 17.2: Buffers:Solutions that resist drastic changes in pH upon additions of small amounts of acid or base.Consist of a weak acid and its conjugate base (usually in salt form)Ex: acetic acid and sodium acetate:HC2H3O2 + NaC2H3O2Or consist of a weak base and its conjugate acid (usually in salt form)Ex: ammonia and ammonium chloride:NH3 + NH4Cl
312 Practice Problem on Titration: If 7.3 mL of 1.25 M HNO3 is required to neutralize mL of a potassiumhydroxide solution, what is the molarity of the potassium hydroxide?0.044 M KOH
313 Titration of a Weak Base and Strong Acid 14pH 7Volume of HCl added (mL)Half Equivalence Point , pH= pKapka or pkb of weak acid or base in a buffer should be clsoe to the desired pH of the buffer solution.
314 Redox: Reduction occurs when an atom gains one or more electrons. Ex: Oxidation occurs when an atom or ion loses one or more electrons.LEO goes GERCopper metal reacts with silver nitrate to form silver metal and copper nitrate:Cu + 2 Ag(NO3) 2 Ag + Cu(NO3)2.
315 Identifying OX, RD, SI Species Ca0 + 2 H+1Cl-1 Ca+2Cl-12 + H20Oxidation = loss of electrons. The species becomes more positive in charge. For example, Ca0 Ca+2, so Ca0 is the species that is oxidized.Reduction = gain of electrons. The species becomes more negative in charge. For example, H+1 H0, so the H+1 is the species that is reduced.Spectator Ion = no change in charge. The species does not gain or lose any electrons. For example, Cl-1 Cl-1, so the Cl-1 is the spectator ion.
316 Oxidizing Agent and Reducing Agent: Oxidizing agent gets reduced itself and reducing agent gets oxidized itself, so a strong oxidizing agent should have a great tendency to accept e and a strong reducing agent should be willing to lose e easily. What are strong oxidizing agents- metals or non metals? Why?Which is the strongest oxidizing agent and which is the strongest reducing agent?
317 Agents Ca0 + 2 H+1Cl-1 Ca+2Cl-12 + H20 Since Ca0 is being oxidized and H+1 is being reduced, the electrons must be going from the Ca0 to the H+1.Since Ca0 would not lose electrons (be oxidized) if H+1 weren’t there to gain them, H+1 is the cause, or agent, of Ca0’s oxidation. H+1 is the oxidizing agent.Since H+1 would not gain electrons (be reduced) if Ca0 weren’t there to lose them, Ca0 is the cause, or agent, of H+1’s reduction. Ca0 is the reducing agent.
318 Steps for Balancing a Redox Reaction: Half Reaction Method In half reaction method, oxidation and reduction half- reactions are written and balanced separately before combining them into a balanced redox reaction. It is a good method for balancing redox reactions because this method can be used both for reactions carried out in acidic and basic medium .
319 Steps for Balancing Redox Reaction Using Half Reaction Method IN ACIDIC MEDIUM: Step 1: Write unbalanced equation in ionic form.Step 2: Write separate half reactions for the oxidation and reduction processes. (Use Oxidation Numbers for identifying oxidation and reduction reactions)Step 3: Balance atoms in the half reactionsFirst, balance all atoms except H and OBalance O by adding H2OBalance H by adding H+Step 4: Balance Charges on each half reaction, by adding electrons.Step 5: Multiply each half reaction by an appropriate number to make the number of electrons equal in both half reactions.Step 6: Add two half reactions and simplify where possible by canceling species appearing in both sides.Step 7: Check equation for same number of atoms and charges on both sides.
320 Writing Half-Reactions Ca0 + 2 H+1Cl-1 Ca+2Cl-12 + H20Oxidation: Ca0 Ca+2 + 2e-Reduction: 2H+1 + 2e- H20The two electrons lost by Ca0 are gained by the two H+1 (each H+1 picks up an electron).PRACTICE SOME!
321 Practice Half-Reactions Don’t forget to determine the charge of each species first!4 Li + O2 2 Li2OOxidation Half-Reaction:Reduction Half-Reaction:Zn + Na2SO4 ZnSO4 + 2 Na
322 Steps for Balancing Redox Reaction Using Half Reaction Method IN BASIC MEDIUM: For balancing redox reactions in basic solutions, all the steps are the same as acidic medium balancing, except you add one more step to it. The H+ ions can then be “neutralized” by adding an equal number of OH- ions to both sides of the equation. Ex.
323 Standard Cell Potential Just as the water tends to flow from a higher level to a lower level, electrons also move from a higher “potential” to a lower potential. This potential difference is called the electromotive force (EMF) of cell and is written as Ecell.The standard for measuring the cell potentials is called a SHE (Standard Hydrogen Electrode).Description of SHE (Standard Hydrogen Electrode)Reaction 2H+(aq, 1M)+ 2e - H2(g, 101kPa) E0= 0.00 V
324 Standard Reduction Potentials Many different half cells can be paired with the SHE and the standard reduction potentials for each half cell is obtained. Check the table for values of reduction potential for various substances:Would substances with high reduction potential be strong oxidizing agents or strong reducing agents? Why?
326 Activity SeriesFor metals, the higher up the chart the element is, the more likely it is to be oxidized. This is because metals like to lose electrons, and the more active a metallic element is, the more easily it can lose them.For nonmetals, the higher up the chart the element is, the more likely it is to be reduced. This is because nonmetals like to gain electrons, and the more active a nonmetallic element is, the more easily it can gain them.
327 Metal Activity3 K0 + Fe+3Cl-13REACTIONMetallic elements start out with a charge of ZERO, so they can only be oxidized to form (+) ions.The higher of two metals MUST undergo oxidation in the reaction, or no reaction will happen.The reaction 3 K + FeCl3 3 KCl + Fe WILL happen, because K is being oxidized, and that is what Table J says should happen.The reaction Fe + 3 KCl FeCl3 + 3 K will NOT happen.Fe0 + 3 K+1Cl-1NO REACTION
328 Voltaic Cells (Galvanic Cells) A voltaic cell converts chemical energy from a spontaneous redox reaction into electrical energy. Ex: Cu and Zn voltaic cell (More positive reduction potential is the cathode)Key Words:CathodeAnodeSalt BridgeHow a Voltaic Cell Works: An Ox, Red CatRepresenting Electrochemical Cells
329 Voltaic CellsProduce electrical current using a spontaneous redox reactionUsed to make batteries!Materials needed: two beakers, piece of the metals (anode, - electrode and cathode + electrode), solution of each metal, porous material (salt bridge), solution of a salt that does not contain either metal in the reaction, wire and a load to make use of the generated current!Use Reference Table J to determine the metals to useHigher = (-) anode (lower reduction potential)Lower = (+) cathode (higher reduction potential)
331 Electrolytic CellsUse electricity to force a nonspontaneous redox reaction to take place.Uses for Electrolytic Cells:Decomposition of Alkali Metal CompoundsDecomposition of Water into Hydrogen and OxygenElectroplatingDifferences between Voltaic and Electrolytic Cells:ANODE: Voltaic (-) Electrolytic (+)CATHODE: Voltaic (+) Electrolytic (-)Voltaic: 2 half-cells, a salt bridge and a loadElectrolytic: 1 cell, no salt bridge, IS the load
332 Decomposing Alkali Metal Compounds 2 NaCl 2 Na + Cl2The Na+1 is reduced at the (-) cathode, picking up an e- from the batteryThe Cl-1 is oxidized at the (+) anode, the e- being pulled off by the battery (DC)
333 Decomposing Water 2 H2O 2 H2 + O2 The H+ is reduced at the (-) cathode, yielding H2 (g), which is trapped in the tube.The O-2 is oxidized at the (+) anode, yielding O2 (g), which is trapped in the tube.
334 ElectroplatingThe Ag0 is oxidized to Ag+1 when the (+) end of the battery strips its electrons off.The Ag+1 migrates through the solution towards the (-) charged cathode (ring), where it picks up an electron from the battery and forms Ag0, which coats on to the ring.
335 Spontaneity of Redox Reactions: E0 = E0red( reduction process-cathode) – E0red (oxidation process-anode)A positive value of E0 indicates a spontaneous process and a negative value of E0 indicates a nonspontaneous value.Steps for Predicting Spontaneity of Redox ReactionsFirst write the reaction as oxidation and reduction half reactions.Then plug standard reduction potential values in the equation given above.Check for the spontaneity by a positive or a negative value of E0Ex:
336 Hydrocarbons Molecules made of Hydrogen and Carbon Carbon forms four bonds, hydrogen forms one bondHydrocarbons come in three different homologous series:Alkanes (single bond between C’s, saturated)Alkenes (1 double bond between 2 C’s, unsaturated)Alkynes (1 triple bond between 2 C’s, unsaturated)These are called aliphatic, or open-chain, hydrocarbons.Count the number of carbons and add the appropriate suffix!
337 Alkanes CH4 = methane C2H6 = ethane C3H8 = propane C4H10 = butane C5H12 = pentaneTo find the number of hydrogens, double the number of carbons and add 2.
338 Methane Meth-: one carbon -ane: alkane The simplest organic molecule, also known as natural gas!
340 Propane Prop-: three carbons -ane: alkane Also known as “cylinder gas”, usually stored under pressure and used for gas grills and stoves. It’s also very handy as a fuel for Bunsen burners!
341 Butane But-: four carbons -ane: alkane Liquefies with moderate pressure, useful for gas lighters. You have probably lit your gas grill with a grill lighter fueled with butane!
342 Pentane Pent-: five carbons -ane: alkane Draw Hexane: Draw Heptane: Your Turn!!!Draw Hexane:Draw Heptane:
343 Alkenes C2H4 = Ethene C3H6 = Propene C4H8 = Butene C5H10 = Pentene To find the number of hydrogens, double the number of carbons.
344 EtheneTwo carbons, double bonded. Notice how each carbon has four bonds? Two to the other carbon and two to hydrogen atoms.Also called “ethylene”, is used for the production of polyethylene, which is an extensively used plastic. Look for the “PE”, “HDPE” (#2 recycling) or “LDPE” (#4 recycling) on your plastic bags and containers!
345 PropeneThree carbons, two of them double bonded. Notice how each carbon has four bonds?If you flipped this molecule so that the double bond was on the right side of the molecule instead of the left, it would still be the same molecule. This is true of all alkenes.Used to make polypropylene (PP, recycling #5), used for dishwasher safe containers and indoor/outdoor carpeting!
346 ButeneThis is 1-butene, because the double bond is between the 1st and 2nd carbon from the end. The number 1 represents the lowest numbered carbon the double bond is touching.This is 2-butene. The double bond is between the 2nd and 3rd carbon from the end. Always count from the end the double bond is closest to.ISOMERS: Molecules that share the same molecular formula, but have different structural formulas.
347 PenteneThis is 1-pentene. The double bond is on the first carbon from the end.This is 2-pentene. The double bond is on the second carbon from the end.This is not another isomer of pentene. This is also 2-pentene, just that the double bond is closer to the right end.
348 Alkynes C2H2 = Ethyne C3H4 = Propyne C4H6 = Butyne C5H8 = Pentyne To find the number of hydrogens, double the number of carbons and subtract 2.
349 Ethyne Now, try to draw propyne! Any isomers? Let’s see! Also known as “acetylene”, used by miners by dripping water on CaC2 to light up mining helmets. The “carbide lamps” were attached to miner’s helmets by a clip and had a large reflective mirror that magnified the acetylene flame.Used for welding and cutting applications, as ethyne burns at temperatures over 3000oC!
350 Propyne This is propyne! Nope! No isomers. OK, now draw butyne. If there are any isomers, draw them too.
351 Butyne Well, here’s 1-butyne! And here’s 2-butyne! Is there a 3-butyne? Nope! That would be 1-butyne. With four carbons, the double bond can only be between the 1st and 2nd carbon, or between the 2nd and 3rd carbons.Now, try pentyne!
352 Pentyne Naming: Check this link out 1-pentyne2-pentyneNow, draw all of the possible isomers for hexyne!
353 IsomersIsomers are compounds that have same molecular formula (same number of atoms) but a different structure.There are three types of isomers:Structural Isomers: Same number of atoms, arranged differently.Geometric Isomers (Cis- trans-): Happens in = or triple bonded compounds since these are inflexible bonds. Ex.Optical Isomers ( D- and L-): Need a central atom that is “Chiral” (all four groups attached to it are different). These are non super imposable mirror images. Usually this central atom is C.
355 Substituted Hydrocarbons Hydrocarbon chains can have three kinds of “dingly-danglies” attached to the chain. If the dingly-dangly is made of anything other than hydrogen and carbon, the molecule ceases to be a hydrocarbon and becomes another type of organic molecule.Alkyl groupsHalide groupsOther functional groupsTo name a hydrocarbon with an attached group, determine which carbon (use lowest possible number value) the group is attached to. Use di- for 2 groups, tri- for three.
358 Organic Families Each family has a functional group to identify it. Alcohol (R-OH, hydroxyl group)Organic Acid (R-COOH, primary carboxyl group)Aldehyde (R-CHO, primary carbonyl group)Ketone (R1-CO-R2, secondary carbonyl group)Ether (R1-O-R2)Ester (R1-COO-R2, carboxyl group in the middle)Amine (R-NH2, amine group)Amide (R-CONH2, amide group)These molecules are alkanes with functional groups attached. The name is based on the alkane name.
361 Positioning of Functional Group PRIMARY (1o): the functional group is bonded to a carbon that is on the end of the chain.SECONDARY (2o): The functional group is bonded to a carbon in the middle of the chain.TERTIARY (3o): The functional group is bonded to a carbon that is itself directly bonded to three other carbons.
362 Organic AcidThese are weak acids. The H on the right side is the one that ionized in water to form H3O+. The -COOH (carboxyl) functional group is always on a PRIMARY carbon.Can be formed from the oxidation of primary alcohols using a KMnO4 catalyst.
363 AldehydeAldehydes have the CO (carbonyl) groups ALWAYS on a PRIMARY carbon. This is the only structural difference between aldehydes and ketones.Formed by the oxidation of primary alcohols with a catalyst. Propanal is formed from the oxidation of 1-propanol using pyridinium chlorochromate (PCC) catalyst.*
364 KetoneKetones have the CO (carbonyl) groups ALWAYS on a SECONDARY carbon. This is the only structural difference between ketones and aldehydes.Can be formed from the dehydration of secondary alcohols with a catalyst. Propanone is formed from the oxidation of 2-propanol using KMnO4 or PCC catalyst.*
365 EtherEthers are made of two alkyl groups surrounding one oxygen atom. The ether is named for the alkyl groups on “ether” side of the oxygen. If a three-carbon alkyl group and a four-carbon alkyl group are on either side, the name would be propyl butyl ether. Made with an etherfication reaction.
366 EsterEsters are named for the alcohol and organic acid that reacted by esterification to form the ester. If the alcohol was 1-propanol and the acid was hexanoic acid, the name of the ester would be propyl hexanoate. Esters contain a COO (carboxyl) group in the middle of the molecule, which differentiates them from organic acids.
367 AmineComponent of amino acids, and therefore proteins, RNA and DNA…life itself!- Essentially ammonia (NH3) with the hydrogens replaced by one or more hydrocarbon chains, hence the name “amine”!
368 Amide Synthetic Polyamides: nylon, kevlar Natural Polyamide: silk! For more information on polymers, go here.
370 CombustionHappens when an organic molecule reacts with oxygen gas to form carbon dioxide and water vapor. Also known as “burning”.
371 Substitution Alkane + Halogen Alkyl Halide + Hydrogen Halide The halogen atoms substitute for any of the hydrogen atoms in the alkane. This happens one atom at a time. The halide generally replaces an H on the end of the molecule.C2H6 + Cl2 C2H5Cl + HClThe second Cl can then substitute for another H:C2H5Cl + HCl C2H4Cl2 + H2
372 Addition Alkene + Halogen Alkyl Halide The double bond is broken, and the halogen adds at either side of where the double bond was. One isomer possible.(c) 2006, Mark Rosengarten
373 Etherification* Alcohol + Alcohol Ether + Water A dehydrating agent (H2SO4) removes H from one alcohol’s OH and removes the OH from the other. The two molecules join where there H and OH were removed.Note: dimethyl ether and diethyl ether are also produced from this reaction, but can be separated out.
374 Esterification Organic Acid + Alcohol Ester + Water A dehydrating agent (H2SO4) removes H from the organic acid and removes the OH from the alcohol. The two molecules join where there H and OH were removed.
375 Saponification The process of making soap from glycerol esters (fats). Glycerol ester + 3 NaOH soap + glycerolGlyceryl stearate + 3 NaOH sodium stearate + glycerolThe sodium stearate is the soap! It emulsifies grease…surrounds globules with its nonpolar ends, creating micelles with - charge that water can then wash away. Hard water replaces Na+ with Ca+2 and/or other low solubility ions, which forms a precipitate called “soap scum”.Water softeners remove these hardening ions from your tap water, allowing the soap to dissolve normally.
376 PolymerizationA polymer is a very long-chain molecule made up of many monomers (unit molecules) joined together.The polymer is named for the monomer that made it.Polystyrene is made of styrene monomerPolybutadiene is made of butadiene monomerAddition PolymersCondensation PolymersRubber
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