Download Topic 2 - Jensen Chemistry

TOPIC 2.1
Dalton’s Atomic Theory (1808)
John Dalton
 All matter is composed of extremely
small particles called atoms
 Atoms of a given element are
identical in size, mass, and other
properties; atoms of different
elements differ in size, mass, and
other properties
 Atoms cannot be subdivided, created, or destroyed
 Atoms of different elements combine in simple
whole-number ratios to form chemical compounds
 In chemical reactions, atoms are combined,
separated, or rearranged
Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said:
Atoms of a given element are identical in size, mass, and
other properties; atoms of different elements differ in size,
mass, and other properties
Modern theory states:
Atoms of an element have a characteristic average mass
which is unique to that element.
Modern Atomic Theory #2
Dalton said:
Atoms cannot be subdivided, created, or destroyed
Modern theory states:
Atoms cannot be subdivided, created, or destroyed in ordinary
chemical reactions. However, these changes CAN occur in nuclear
reactions
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to
deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
Thomson’s Atomic Model
J.J. Thomson
Thomson believed that the
electrons were like plums
embedded in a positively charged
“pudding,” thus it was called the
“plum pudding” model.
Mass of the Electron
1909 – Robert Millikan
determines the mass of the
electron.
The oil drop
apparatus
Mass of the
electron is
9.109 x 10-31 kg
Conclusions from the Study of the
Electron
 Cathode rays have identical properties
regardless of the element used to produce them.
All elements must contain identically charged
electrons.
Atoms are neutral, so there must be positive
particles in the atom to balance the negative
charge of the electrons
 Electrons have so little mass that atoms must
contain other particles that account for most of
the mass
Rutherford’s Gold Foil Experiment
 Alpha particles are helium nuclei
 Particles were fired at a thin sheet of gold foil
 Particle hits on the detecting screen (film) are recorded
Try it Yourself!
In the following pictures, there is a target hidden by a
cloud. To figure out the shape of the target, we shot
some beams into the cloud and recorded where the
beams came out. Can you figure out the shape of the
target?
The Answers
Target #1
Target #2
Rutherford’s Findings
 Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
“Like howitzer shells bouncing off of
tissue paper!”
Conclusions:
 The nucleus is small
 The nucleus is dense
 The nucleus is positively charged
Atomic Particles
Particle
Charge
Mass (kg)
Electron
-1
9.109 x 10-31
Proton
+1
1.673 x 10-27
1
Nucleus
0
1.675 x 10-27
1
Nucleus
Neutron
Relative mass
(amu)
0.0005
Location
Electron
cloud
The Atomic
Scale
Helium-4
 Most of the mass of the atom is
in the nucleus (protons and
neutrons)
 Electrons are found outside of
the nucleus (the electron cloud)
 Most of the volume of the atom
is empty space
Image: User Yzmo Wikimedia Commons.
Atomic Number
Atomic number (Z) of an element is
the number of protons in the nucleus
of each atom of that element. This
identifies the atom.
Element
Carbon
Phosphorus
Gold
# of protons Atomic # (Z)
6
6
15
15
79
79
Mass Number
Mass number is the number of
protons and neutrons in the
nucleus of an isotope.
Mass # = p+ + n0
Nuclide
p+
n0
e-
Mass #
Oxygen - 18
8
10
8
18
Arsenic - 75
Phosphorus- 31
33
42
33
75
15
16
15
31
Elec trons mass are so much smaller than a proton and neutron that they don’t
Contribute much to the overall mass of the atom ,therefore they are not
Counted in the mass number of the atom.
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Isotope
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Carbon = 12.011
Ions
■ When an atom gains or loses an electron, it is called an
ion.
■ Cation if positively charged or giving electrons away.
Electrons leave the highest shell first.
Anion if negatively charged or accepting electrons.
Purpose of Mass Spectrometry
 Produces spectra of masses from the molecules in a
sample of material, and fragments of the molecules.
 Used to determine
 the elemental composition of a sample
 the masses of particles and of molecules
 potential chemical structures of molecules by
analyzing the fragments
 the identity of unknown compounds by determining
mass and matching to known spectra
 the isotopic composition of elements in a molecule
Stages
The ionizer converts some
of the sample into ions.
Mass analyzers separate
the ions according to
their mass-to-charge
ratio.
The detector records
either the charge induced
or the current produced
when an ion passes by or
hits a surface
Interpreting Mass Spectra
The height of each peak is proportional to the amount of each
isotope present (i.e. it’s relative abundance).
The m/z ratio for each peak is found from the accelerating
voltage for each peak. Many ions have a +1 charge so that the
m/z ratio is numerically equal to mass of the ion.
Calculating the relative atomic mass
from mass spectrometry
1. Measure the height of each peak.
2. Calculate the percentage relative abundance
% abundance =
amount of isotope x 100
total amount of all isotopes
3. Calculate the average mass
mass of isotope A x abundance + mass of isotope B x abundance
100
TOPIC 2.2
Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the
electron as a particle.
His son, George Thomson won the Nobel prize for
describing the wave-like nature of the electron.
The
electron is
a particle!
The
electron is
an energy
wave!
The Wave-like Electron
The electron propagates through space
as an energy wave. To understand the
atom, one must understand the behavior
of electromagnetic waves.
Louis deBroglie
Electromagnetic radiation propagates through space
as a wave moving at the speed of light.
c = νλ
c = speed of light, a constant (3.00 x 108 m/s)
ν = frequency, in units of hertz (hz, sec-1)
λ = wavelength, in meters
sometimes waves are measure in nanometers: 1 x109 nm = 1 m
The energy (E ) of electromagnetic radiation is
directly proportional to the frequency (ν) of the
radiation.
E = hν
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x 10-34 J·s)
ν = frequency, in units of hertz (hz, sec-1)
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High
Frequency
=
High ENERGY
Electromagnetic Spectrum
■ (ultra red) ROY G BIV (ultra violet)
■Large wavelengths to short wavelengths
Light is a Particle (Quantum Theory)
• Planck:
Energy can be released or absorbed in packets or fixed amounts of
a standard size he called quanta (singular: quantum). (photon)
Joule (J) is used to express energy.
Particle theory can be compared to a piano where the wave
theory is compared to the violin.
High frequency waves have high energy.
Planck’s constant (h) = 6.63 x 10-34 J-s
Max Planck
(1858-1947)
E = hν =
hc
λ
32
Niels Bohr’s Atom
■ Electrons orbit the nucleus in orbits, like a solar system.
Planetary Model
Electrons cannot
exist between orbits
(energy is quantized)
Bohr’s atom
■ Electrons closest to the nucleus are
lowest in energy.
■ Ground state- electrons are in the
lowest energy level possible
■ If energy is put into the atom, the
electrons will jump up in energymove away from the nucleus
(excited state).
Bohr’s atom
■ Excited electrons naturally go back
to ground state. In order to do this,
energy must leave the atom.
Because energy is quantized in an
atom, the amount of energy that
leaves is the difference in energy
between orbits If this energy is in
the visible light range, we will see
certain colors (line emission
spectrums)
Calculating Energy Change, ∆E, for
Electron Transitions
Energy must be absorbed from a photon (+∆E) to
move an electron away from the nucleus
Energy (a photon) must be given off (-∆E) when an
electron moves toward the nucleus
Electron transitions
involve jumps of
definite amounts of
energy.
This produces bands
of light with definite
wavelengths.
Spectroscopic analysis of the hydrogen
spectrum…
…produces a “bright line” spectrum
Bohr’s Model of the H Atom (and only H!)
Atomic emission spectra:
– Most sources produce light that contains many
wavelengths at once.
– However, light emitted from pure substances may
contain only a few specific wavelengths of light called
a line spectrum (as opposed to a continuous
spectrum).
– Atomic emission spectra are inverses of atomic
Hydrogen:
containsspectra.
1 red, 1 blue and 1 violet.
absorption
Carbon:
39
Flame Tests
Many elements give off characteristic light which
can be used to help identify them.
strontium
sodium
lithium
potassium
copper
The Bohr Model of the Atom
I pictured electrons
orbiting the nucleus much
like planets orbiting the
sun.
But I was wrong! They’re
more like bees around a
hive.
Neils Bohr
Heisenberg Uncertainty Principle
“One cannot simultaneously determine both
the position and momentum of an electron.”
The more certain you are about where the
electron is, the less certain you can be
about where it is going.
Werner
Heisenberg
The more certain you are about where the
electron is going, the less certain you can
be about where it is.
Quantum Mechanical
Model of the Atom
Mathematical laws can identify the
regions outside of the nucleus where
electrons are most likely to be found.
These laws are beyond the scope of
this class…
Erwin Schrödinger 1887-1961
■ Born in Vienna, Austria
■ University of Berlin professor
■ Nobel Prize 1933
■ Based on waves of light and probability.
■ Quantum Mechanical Model –
Electron density distribution in H atom
45
Quantum numbers, Energy levels and Orbitals
■ 4 quantum numbers (like an address)
■ n, l, ml , ms
■ Quantum number
1. n= principle Energy level (period)
2. l = sublevel orbital
3. ml = orientation of orbital
4. ms= spin = ½, -½
Energy levels (shells)- n
■ n = 1, 2, 3,..
■ n= 1 is lowest in energy and closest to the nucleus
■ As n increases, the distance from the nucleus increases
■
)
)
) ))
n=1
2
3 45
Electron Energy Level (Shell)
Generally symbolized
by n, it denotes the
probable distance of
the electron from the
nucleus. “n” is also
known as the Principle
Quantum number
Number of electrons
that can fit in a shell:
2n2
Electron Orbitals
An orbital is a region within an energy level where
there is a probability of finding an electron.
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.
The angular momentum quantum number,
generally symbolized by l, denotes the orbital
(subshell) in which the electron is located.
Representations of Orbitals
s orbital
p orbitals
50
d orbitals
f orbitals
52
53
Energy Levels, Sublevels, Electrons
Energy
Level
(n)
Sublevels in
main energy
level
(n sublevels)
Number of
orbitals per
sublevel
Number of
Electrons
per sublevel
Number of
electrons per
main energy
level (2n2)
1
s
1
2
2
2
s
p
1
3
2
6
8
3
s
p
d
1
3
5
2
6
10
18
4
s
p
d
f
1
3
5
7
2
6
10
14
32
Electron Spin
The Spin Quantum Number describes
the behavior (direction of spin) of an
electron within a magnetic field.
Possibilities for electron spin:
1
+
2
1
−
2
Electron configuration
Tells where all the electrons are in an atom
– Orbital notation (uses arrows to represent electrons)
– Electron configuration (uses superscripts to represent electrons)
– Noble gas configuration (uses a noble gas to represent core electrons and
superscripts for electrons)
Pauli Exclusion Principle
Two electrons occupying
the same orbital must
have opposite spins
Wolfgang
Pauli
■ A maximum of 2 e- can be in an orbital
■ Number of shapes in a shell = n
■ Number of orbitals in a shell = n2
■ Number of electrons in a shell = 2n2
■ s has 1 orbital, p has 3, d has 5, f has 7
■ s can hold 2 e-, p:6, d: 10, f:14
Aufbau principle
■ Electrons want to have the lowest energy possible.
■ First energy level fills up before 2nd , then 3rd, etc.
■ Within a shell s orbital fills up, then p orbitals, d orbitals, and last f orbitals
■ 1s 2s 2p 3s 3p 4s 3d (refer to book)
■ 4s fills before 3d because it is actually lower in energy
Filling Order of Orbitals
1. Aufbau principle: e- enter orbitals of lowest
energy first
7p
7s
6s
5s
6p
5p
4p
4s
6d
5f
x7
5d
4f
x7
4d
3d
3p
3s
2s
2p
1s
• Relative stability & average distance of e- from nucleus
60
Hund’s rule
■ Electrons want to be in their own orbital before they pair up (without breaking
aufbau’s principle)
■ All electrons in the same subshell must have the same spin (be in their own orbital
before pairing up)
Periodic table and e- configuration
■ The periodic table is set up to show the electron filling order (lowest to highest
energy)
■ We can use the periodic table to help write electron configurations and noble gas
configurations
Periodic table
■ Groups 1-2: s block (2 columns)
■ Groups 13-18: p block (6 columns)
■ Groups 3-12: d block (10 columns)
■ Lanthanides and actinides: f block (14 columns)
■ Periods tell the highest shell that is filled
There are some exceptions, but you will not have to worry about them.
Orbital filling table
Electron configuration of the elements of
the first three series
Electronic Configurations
Element
Standard Configuration
Noble Gas
Shorthand
Nitrogen
1s22s22p3
[He] 2s22p3
Scandium
1s22s22p63s23p64s23d1
[Ar] 4s23d1
Gallium
1s22s22p63s23p64s23d104p1
[Ar] 4s23d104p1
66
Noble Gas
Shorthand
Element
Standard Configuration
Lanthanum
1s2 2s22p6 3s23p6 4s23d104p6
5s24d105p6 6s25d1
[Xe] 6s25d1
Cerium
1s2 2s22p6 3s23p6 4s23d104p6
5s24d105p6 6s25d14f1
[Xe] 6s25d14f1`
1s2 2s22p6 3s23p6 4s23d104p6
Praseodymium
5s24d105p6 6s24f3
[Xe] 6s24f3
67