Download Chapter7 - FSU Chemistry

!Hcomb of benzoic acid (C6H5CO2H) is -3227 kJ/mol. When
a 2.442-g sample of benzoic acid is burned in a
calorimeter, all of the resulting heat is transfered to 2.5 kg
of water (c=4.184 J/g.K). What is the temperature change?
Calculate !Hrxn for
2 NOCl
N2 + O2 + Cl2
given the following set of reactions
N2 + O2
2NO + Cl2
2NO
2NOCl
!H = 180.6 kJ
!H = -77.2 kJ
Combustion of octane C8H18 produces carbon dioxide and water.
C8H18 + O2
CO2 + H2O
(a) Balance this reaction and calculate !Hocomb for 1 mol of octane, assuming water
forms as a gas;
!Hof = -208.45 kJ/mol for octane
!Hof = -393.5 kJ/mol for carbon dioxide
!Hof = -241.8 kJ/mol for H2O (g)
(b) How many liters of CO2 (at 25 oC and 1 atm, R = 0.0821 atm.L/mol.K) are formed
when 1 kg of octane is burned? How much work is done by the expanding CO2 as 1
kg of octane is burned (again, at 25 oC and 1 atm). (Hint, 1 J = 9.87.10-3 atm.L).
What is !E for the reaction? (Hint, the definition of H = E + PV, assume PV is done
by carbon dioxide only (ignore water and octane))
Atomic Structure: Quantum Theory
nuclear model from
Ratherford’s experiments
Problem with Classic Structure
accelerating charged
particle loses energy
The Nature of Light
! frequency (")
! amplitude
#= c
"
Electromagnetic Spectrum
Diffraction and Interference
diffraction
interference
here, light behaves as a wave
Blackbody Radiation
However:
0K
1000 K
1500 K
Max Planck, 1900
E = nh"
n is a positive whole number, quantum number
h is a constant
" is the frequency of light emitted
Photoelectric Effect
• threshold frequency
• no time lag
Einstein, ~1905
Ephoton = h"
> 2000 K
Photoelectric Effect
7.96. Electric power is typically stated in units of watts (1W
= 1J/s). About 95% of the power output of an incandescent
bulb is converted to heat and 5% to light. If 10% of that light
shines on your chemistry text, how many photons per
second shine on the book from a 75-W bulb? (assume a
wavelenght of 550 nm)
The Nature of Light
• electromagnetic radiation travels in waves
• at the speed of light (in vacuum, c = "#)
• intensity is either amplitude or number of photons/second
• energy of light is quantized, E = h"
• the energy of atom is also quantized, E = nh"
h = 6.626$10-34 J•s
Atomic Spectra
H
He
Na
sunlight
Ba
K
Atomic Spectra (Neon)
Atomic Spectra (H)
Rydberg equation
1 =R
!
1
22
_
1
n2
The Bohr Model of the
Hydrogen Atom
Niels Bohr proposed these postulates:
1. The H atom has fixed energy levels: stationary states
2. The atom does not radiate energy while in one of
those states
3. Change to another state: !E = !Ephoton = h"
The Bohr Model of the
Hydrogen Atom
Niels Bohr proposed these postulates:
E = h!
ground state: the
lowest energy
The Bohr Model of the
Hydrogen Atom
E = h!
the charge of the nucleus
-18
E = -2.18x10 J
Z2
n2
energy levels for the H atom
De Broglie: electrons are
waves
matter is wavelike:
#=
h
mv
electrons within atoms behave like waves
The Quantum-Mechanical
Model
Schrödinger equation:
d2%
d2%
d2%
8&2me
[E - V(x,y,z)]%(x,y,z)=0
+
+
+
dx2
dy2
dz2
h2
or
H% = E%
%2 is the probability that the electron is in a certain
region of space
The Quantum-Mechanical
Model
Schrodinger equation:
H% = E%
solutions (energy states) are associated with a given %,
called an atomic orbital
Hydrogen Atom
“point” probability
radial probability
Hydrogen Atom
32%
74%
99%
Atomic Orbitals
different energy states are associated with atomic orbitals
atomic orbitals are characterized by 3 quantum numbers
n (1,2,3,
etc.) is the size (level)
principal quantum number
l (fromangular
0 to n-1) is the shape (sublevel)
momentum quantum number
ml (from
-l to +l) is the orientation
magnetic quantum numbber
Atomic Orbitals
Sample Problem 7.5
Specify possible orbitals (sets of n, l, m) for n=3.
Atomic Orbitals
n (level)
1,2,3, etc
l (sublevel, shape)
s (l=0), p (l=1), d (l=3), f(l=4), etc
S orbital - m equals 0
Atomic Orbitals (s)
1 s orbital
3 s orbital
Atomic Orbitals: s
energy increases
1s
n=1
l=0
2s
n=2
l=0
3s
n=3
l=0
4s
n=4
l=0
S orbitals means l=0
(darker color indicates highest radial probability %2)
Atomic Orbitals: p
energy increases
1p
n=1
l=1
2p
n=2
l=1
3p
n=3
l=1
P orbitals means l=1
2p
Atomic Orbitals: p
2px
2py
n=2
l=1
m = -1
n=2
l=1
m=0
2pz
n=2
l=1
m = +1
Atomic Orbitals: d
d orbitals: start with n = 3; l = 2
Energy
Energy Levels: Hydrogen
3s
3p
2s
2p
1s
3d
Energy
Energy Levels: Multielectron
Systems
3s
3p
2s
2p
3d
1s
Practice! Problems
7.49. How many orbitals in an atom can have each of the
following designations:
(a) 1s
(b) 4d
(c) 3p
(d) 2f
(e) n = 3
Do 7.50
Practice! Problems
7.73. Carbon-carbon bonds form the backbone of nearly
every organic and biological molecule. The average
bond energy of the C-C bond is 347 kJ/mol. Calculate
the frequency and wavelength of the least energetic
photon that can break this bond. In what region of
electromagnetic spectrum is this radiation?
Practice! Problems
*7.76. A ground state H atom absorbs a photon of wavelength
94.91 nm, and its electron attains a higher energy level.
The atom then emits two photons: one of wavelength 1281
nm to reach an intermediate level, and a second to return
to the ground state.
(a) What level did the electron reach?
(b) What intermediate level did the electron reach?
(c) What was the wavelength of the second photon emitted?
Practice! Problems
*7.76. A ground state H atom absorbs a photon of wavelength
94.91 nm, and its electron attains a higher energy level.
The atom then emits two photons: one of wavelength 1281
nm to reach an intermediate level, and a second to return
to the ground state.
(a) What level did the electron reach?
(b) What intermediate level did the electron reach?
(c) What was the wavelength of the second photon emitted?
Bohr and Schrödinger: !E = -2.18x10-18J
1
n
2
final
1
n2initial