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CHAPTER 7 AND 16 AP CHEMISTRY.

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Presentation on theme: "CHAPTER 7 AND 16 AP CHEMISTRY."— Presentation transcript:

1 CHAPTER 7 AND 16 AP CHEMISTRY

2 LIGHT Wavelength = λ nm/ wave, m/wave, or nm or m
λf= c; c = x 108 m/s Frequency = f Waves/s or 1/s = c/ λ High frequency = short wavelength Low frequency = long wavelength The wavelength of the radiation which produces the yellow color of a sodium vapor light is nm. What is the frequency of the radiation? / 109 nm / 2.998x 108 m 589.0 nm/ 1 m / s 5.090 x /s

3 PHOTON ENERGY E = hf Planck’s constant h = 6.626 x 10-34 J.s
From the last question, give the amount of energy, in Joules, that a photon would emit by an excited sodium atom? In kilojoules? 3.373 x J 3.373 x kJ Continuous spectrum All wavelengths Line spectrum Discrete energies light going through a prism ROY G BIV

4 HYDROGEN ATOM Bohr model used hydrogen atom
Electron orbits around nucleus Energy levels: Lyman 364 and Balmer 656 Zero energy where a photon and electron are in completely separate energy states within an atom Ground state lowest amount of energy an electron has Excited state electron getting extra energy E = hf = Ehi - Elo Electrons can only give off certain energies

5 EXAMPLE A green light of wavelength 456 nm is observed in emission spectrum of hydrogen. Calculate the energy difference, Ehi -Elo, between the two states that are responsible for this line in kJ/mole / 109 nm / 2.998x 108 m 456 nm / 1 m / s 6.626 x J.s / x 1014 / kJ_____ / s / J 4.356 x kJ / 6.02 x 1023atoms Atom / mole 262 kJ/mol

6 QUANTUM NUMBERS Most probable place an electron can be found
Ψ amplitude (height) of the electron wave Ψ 2 directly proportional to the probability of finding the electron Quantum numbers Principal energy level n = 1, 2, 3, etc. Sublevels s, p, d, and f l = 0, 1, 2, 3, etc.

7 CONTINUE Orbitals or orientations m = l to -l l = 0 then m = 0
Spin s = +1/2 , -1/2 Pauli Exclusion Principle No two electrons can have the same four quantum numbers Capacities Each energy level (n) has n sublevels Each sublevel (l) has 2l + 1 orbitals Each orbital (m) has 2 electrons Total number of orbitals in an energy level is n2 and total number of electrons in an energy level is 2n2 How many electrons can fit in principal energy level 3?

8 CONTINUE Orbital does not mean the electron is circling the nucleus
We do not know how the electron moves Heisenberg principle state that there are limitations in knowing what the position and momentum are at any given time This uncertainty does not mean much to a large particle (ball or you) but with an electron this means a lot Probability distribution intensity of color -electron density map Most probable place to find a hydrogen electron is .529Ǻ from nucleus 1Ǻ = m Radius of 1s orbital is to encompass 90% of the probability for that electron

9 ELECTRON CONFIGURATION
Electron notation Usually a sublevel is filled to capacity before entering the next one Do the electron configuration of the following Li 1s22s1 B 1s22s22p1 N 1s22s22p3 O 1s22s22p4 What are the electron configuration of the chlorine atom and iron atom? What are the abbreviated electron configurations for the last two atoms?

10 QUANTUM NUMBERS Pages 296 to 298 Polyelectronic atoms
Treat electrons as if they have nuclear attraction and average repulsion from other electrons Effective nuclear charge Zeff = Zactual - (electron repulsion) Z = atomic number Periodic table can predict the filling of the sublevels Elements in groups 1 and 2 fill s Elements in groups fill p Transition metals in groups fill d The two sets of 14 elements at the bottom of the table fill f Write the electron configuration of tellurium.

11 MENDELEEV In 1869 Mendeleev introduced an arrangement of the elements based on the fact that their chemical properties vary periodically with their atomic mass He left gaps in his table and predicted some detailed physical and chemical properties for three elements that had not be found yet. i.e. germanium (ekasilicon) mass 73 the observed mass is 72.3 Some atoms have different electron configurations than would be predicted Cr predicted 1s22s22p63s23p64s23d4 Actual 1s22s22p63s23p64s13d5 Cu predicted 1s22s22p63s23p64s23d9 Actual 1s22s22p63s23p64s13d10 Aufbau principle (building up) - add one electron to lowest energy orbit until it is filled Valence electrons = outer most electrons Inner electrons = core electrons

12 ORBITAL DIAGRAM B C Hund’s rule
Within a given sublevel, the order of filling is such that there is the maximum number of half-filled orbitals Write the orbital diagram for silicon atom and cobalt atom Monoatomic ion electron configuration Ions with noble gas configuration Elements close to a noble gas will form an ion that has the same configuration as the noble gas 19K --> K+ + 1e Ca --> Ca2+ + 2e- 17Cl-+ 1e- --> Cl S2- + 2e- --> S Anions are formed from atoms with 1, 2, or 3 less electrons than the next noble gas Cations are formed from atoms with 1, 2, or 3 more electrons than the preceding noble gas

13 TRANSITION METAL CATIONS
When transition metal atoms form positive ions, the outer s electrons are lost first Mn atom 1s22s22p63s23p64s23d5 Mn2+ ion 1s22s22p63s23p63d5 Cu atom 1s22s22p63s23p64s13d10 Cu+ ion 1s22s22p63s23p63d10 Cu2+ ion 1s22s22p63s23p63d9 Give the abbreviated electron configuration of Ti atom and Ti2+ ion S atom and S2- ion Ni atom and Ni2+ ion

14 PERIODIC TRENDS Atomic radius In general atomic radii
Decreases as you move left to right Increases as you move down Effective nuclear charge Zeffective = Z - S Z = number or protons S = number of core electrons that are shielding the outer electrons from the nucleus

15 IONIC RADIUS Ionic radius
Ionic radius increases as you move down the group Cations decrease left to right Anions decrease left to right Negative ions are larger than the nonmetal atoms from which they are formed Positive ions are smaller than the metal atoms from which they are formed. Using only the periodic table, arrange each of the following sets of atoms and ions in order of increasing size Na, Al, Sr O, Cl-, Se2- V, V2+, V3+

16 IONIZATION ENERGY Energy required to remove an electron
First ionization energy Na(g) ---> Na+(g) + 1e- Mg(g) ---> Mg+(g) + 1e- Second ionization energy Mg+(g) ---> Mg2+(g) + 1e- Increase as you move across left to right Decrease as you move down Inverse correlation to atomic radius Consider three elements C,N, P. using the periodic table, predict which of the elements has The largest atomic radius. The smallest The largest ionization energy. The smallest

17 CONTINUED Energy required to remove an electron from a gaseous atom
E = (Z2eff/n2)(1310)kJ/mol n = Zeff = 10.3 n = Zeff = 1.84 3s has 10 electrons repulsing the 11th electron. This reduces the pull of the electron toward it (shielding) Shielding is not totally effective, because an outer electron could be found by the nucleus (though probability is low) Page 318 fig Core electrons provide effective shielding for the outer electrons Page 321

18 ELECTRONEGATIVITY Energy released from an atom as it acquires another electron Page 324 Atomic radius Decreases going across because the valence electrons are being pulled in, due to an increase of proton attraction. Shielding remains constant across the period Electronegativity How much an atom wants an electron Which element has the highest electronegativity? Fluorine

19 METALS Metals Metallic luster, ductile, malleable, good conductor of heat and electricity Melting point increases with an increase in atomic number as you go into the center of the periodic table Metals tend to lose electrons when they undergo chemical reactions 2Mg(s) + O2(g) ----> 2MgO(s) (oxidation number?) Compounds formed from metals and nonmetals tend to be ionic Most metal oxides are basic oxides Metal oxide Water Metal hydroxide MgO(s) H2O(l) > Mg(OH)2(aq) Metal oxide Acid Salt Water Fe2O3(s) HCl(aq) > 2FeCl3(aq) + H2O(l)

20 METALS AND NONMETALS Nonmetals Metals + nonmetals ----> salts
No luster, poor conductors Nonmetals tend to gain electrons to become anions Metals + nonmetals ----> salts When compounds are made up chiefly of nonmetals they are molecular compounds Most nonmetal oxides are acidic oxides Nonmetal oxide water Acid CO2(g) H2O(l) ----> H2CO3(aq) Nonmetal oxide Base Salt Water SO3(g) KOH(aq) ----> K2SO4(aq) + H2O(l) Semimetals Some properties of both, brittle, semi-conductor Trends in metallic characteristics The more an element shows physical and chemical properties characteristic for metals, increase right to left and top to bottom

21 THERMODYNAMICS Three thermodynamic functions
ΔH; Change in enthalpy, negative values of ΔH tend toward spontaneous ΔS; change in entropy, positive values of ΔS tend toward spontaneous Nature spontaneously moves to the highest probability of occurrence Increase of entropy Ssolid < Sliquid <Sgas Which has more entropy Solid CO2 or gaseous CO2 N2 gas at 1 atm or N2 gas at 1.0 x 10-2 atm N2 gas at 1.0 x 10-2 atm would be right Predict the sign of entropy Solid sugar is added to water to form a solution Iodine vapor condenses on a cold surface to form crystals ΔG; change in free energy, a reaction at constant temperature and pressure will be spontaneous if ΔG is negative Does not mean it is fast Means it occurs without intervention Tells us which direction it will go

22 SECOND LAW OF THERMODYNAMICS
Any spontaneous process will increase entropy in the universe. Entropy will always increase Example CH4(g) + 2O2(g) ----> CO2(g) + 2H2O(g) Factors affecting it: Energy factor: at 25 oC, 1 atm, exothermic reactions are spontaneous (H < 0) Randomness factor; when other things are equal system tends to move from an ordered structure to a random one

23 ENTROPY ΔS = Sproduct - Sreactants
This is done the same way as enthalpy Measure of change in order As the number of moles increase, the ΔS will be positive (usually) 2SO3(g) ----> 2SO2(g) + O2(g): ΔS = 187 J ΔS = [2(248) + 1(205)] - [2(257)] N2(g) + 3H2(g) ----> 2NH3(g); what is the ΔS? Reactions where ΔSo is positive tend to be spontaneous at high temperatures

24 WATER H2O(l) ----> H2O(g) Δ Ssys is positive, entropy increases
Δ Ssurr determined by the flow of heat. Heat is going into the system, randomness decreases, Δ Ssurr is negative When Δ Ssys = +, and Δ Ssurr = -. Δ Suniv will be temperature dependent Above 100 oC, spontaneous the direction it is written Below 100 oC, occurs opposite direction Sign of Δ Ssurr depends on the direction of the heat flow - lowest energy Magnitude of ΔSsurr depends on the temperature ΔSsurr = +J/K exothermic ΔSsurr = -J/K endothermic ΔSsurr = -ΔH (kJ need to change to J) T(oC need to change to K)

25 PROBLEM In the metallurgy of antimony, the pure metal is recovered via different reactions, depending on the composition of the ore. For example Sb2S3(s) + 2Fe(s) -----> 2Sb(s) + 3FeS(s) ΔH = -125kJ Sb4O6(s) + 6C(s) > 4Sb(s) + 6CO(g) Δ H = 778 kJ Δ Ssurr = - Δ H (kJ need to change to J) T (oC need to change to K) _ x 105J 298 K 419 J/K ΔS positive so exothermic _ 7.78 x 105J -2.61 x 103J/K ΔS negative so endothermic Page 778

26 FREE ENERGY ΔG = ΔH - T ΔS At a constant temperature and pressure a process will be spontaneous if ΔG is negative, ΔG (free energy) is decreasing. Energy not used in breaking down or building up ΔSo = ΔSsys , (o) all substances are in their standard state Page 780 tables 16.4 and 16.5 What temperature is this process spontaneous at 1 atm? Br2(l) -----> Br2(g) What is the normal boiling point? (Hint at the boiling point Br2(l) and Br2(g) are at equilibrium so ΔGo = 0) 333K above this number ΔSo controls it, below this number ΔHo controls it.

27 THIRD LAW OF THERMODYNAMICS
At zero Kelvin the entropy of a perfect crystal would be zero (unattainable ideal) ΔS found at 298 K and 1 atm The more complex a molecule is, the higher ΔS value is ΔG = change in free energy occurs when reactants are in standard states forming products in standard state As the ΔG increases in negative value the further right it will go to reach equilibrium 2SO2(g) + O2(g) -----> 2SO3(g) Carried out at 25 oC and 1 atm. Calculate ΔH, ΔS, and ΔG ΔHf (kJ/mol) ΔS (J/K.mol) SO2(g) SO3(g) O2(g)

28 ΔG = ΔH - T ΔS ΔG is dependent on pressure, concentration, and temperature ΔG < 0 spontaneous ΔG > 0 nonspontaneous ΔG = 0 equilibrium Effect of ΔH and ΔS on spontaneity If ΔH > 0, ΔS < 0, ΔG > 0 at all temperatures it will be nonspontaneous If ΔH < 0, ΔS > 0, ΔG < 0 at all temperatures it will be spontaneous If ΔH > 0, ΔS > 0, then ΔG > 0 at low temperatures, and ΔG < 0 at high temperatures so it will be spontaneous at high temperatures If ΔH < 0, ΔS < 0, then ΔG < 0 at low temperatures, and ΔG > 0 at high temperatures so it will be nonspontaneous at high temperatures

29 SECOND METHOD TO CALCULATE ΔG
Hess’s law Using the following data (25oC) Cdiamond(s) + O2(g) ----> CO2(g) ΔG = -397 kJ Cgraphite(s) + O2(g) ----> CO2(g) ΔG = -394 kJ Calculate ΔG for the reaction Cdiamond(s) > Cgraphite(s) Third method to calculate ΔGf (formation) ΔGf = ∑ np ΔGf (product) - ∑nrΔGf (reactant) Methanol is a high-octane fuel. Calculate ΔG for the reaction 2CH3OH(g) + 3O2(g) ----> 2CO2(g) + 4H2O(g) ΔGf (kJ/mol) ΔGf (kJ/mol) CH3OH(g) CO2(g) O2(g) H2O(g)


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