Transcript Chapter 11

Modern Atomic Theory
11.1-11.2
Electromagnetic Radiation
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Electromagnetic radiation – forms of radiant energy (light in
all its varied forms)
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Electromagnetic spectrum – a continuous range of
wavelengths and frequencies of all forms of electromagnetic
radiation
11.2 Electromagnetic Radiation
Radiation energy – has
wavelike properties
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Frequency (υ, Greek nu) – the
number of peaks (maxima) that
pass by a fixed point per unit time
(s-1 or Hz)
Wavelength (λ, Greek lambda) –
the length from one wave
maximum to the next
Amplitude – the height
measured from the middle point
between peak and trough
(maximum and minimum)
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Intensity of radiant energy is
proportional to amplitude
Electromagnetic Radiation
11.3 Emission of Energy by
Atom
 How does atom emit light?
 Atoms absorbs energy
 Atoms become excited
 Release energy
 Higher-energy photon –>shorter wavelength
 Lower-energy photon -> longer wavelength
11.4 The Energy Levels of Hydrogen
 Excited state: atom with excess energy
 Releasing energy by emitting a photon
 Different wavelengths of light carry different amounts of
energy
 Energy contained in the photon corresponds to the change in
energy that the atom experiences
 Ground state: the lower energy level of an atom
 The level of energy of hydrogen and all other atoms are
quantized
The energy level
11.5 The Bohr Model of the
Atom
 the Bohr model created by
Niels Bohr depicts the atom
as a small, positively charged
nucleus surrounded by
electrons that travel in circular
orbits around the nucleus
 similar in structure to the solar
system, but with electrostatic
forces providing attraction,
rather than gravity
 Describe the behavior of
electrons in an atom
11.6 The Wave Mechanical Model of
the Atom
Schröndinger’s quantum mechanical model of atomic
structure is frame in the form of a wave equation;
describe the motion of ordinary waves in fluids.
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i. Wave functions or orbitals (Greek, psi , the
mathematical tool that quantum mechanic uses to describe
any physical system
ii. 2 gives the probability of finding an electron within a
given region in space
iii. Contains information about an electron’s position in 3D space
defines a volume of space around the nucleus where there
is a high probability of finding an electron
say nothing about the electron’s path or movement
11.7 The Orbitals
Orbital: the probability map for hydrogen electrons
The principal quantum number (n): Shell
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a. describes the size and energy level of the orbital
a.
positive integer (n = 1, 2, 3 …..)
as the value of n increases
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the number of allowed orbital increases
size of the orbital increases
the energy of the electron in the orbital increases
The Orbitals
 As the value of n increases, the number of allowed orbitals
increases and the size of the orbitals become larger, thus
allowing an electron to be far from the nucleus, because it
takes energy to separate a negative charge from a positive
charge
 E.g n = 3  third shell (period #3)
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n = 5  fifth shell (period # 5)
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n = 2  2nd shell (period 2)
The Orbitals
Orbitals are grouping
in group according to
the angular-momentum
quantum number l is
called subshells.
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Types of orbitals
Notations: s, p, d, f
The shape of the orbitals
 S orbital
 P orbital
Summary
 Subslevels (type of orbitals) Present
1s (1)
2s (1)
2p (3)
3s (1)
3p (3)
3d (5)
4s (1)
4p (3) 4d (5) 4f (7)
11.8 The Wave Mechanical Model
 Pauli exclusion principle: an atomic orbital can hold a
maximum of two electrons, and those two electrons must
have opposite spins
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons
11.9 - Electron Arrangements in the First
Eighteen Atoms on the Periodic Table
 Recall: Atomic number (Z) =
# electrons = # protons
 Electron configuration:
describes the orbitals that are
occupied by the electrons in
an atom
 Orbital diagrams: describe the
orbitals with arrows
representing electrons
 a. Arrows are written up or
down to denote electron’s spin
Z=2
He = 1 s2
Example
 Write the full electron configuration and orbital filling
diagram for: O, Na, Si, Ar, Cr
Electrons Configuration
Shorthand version – give the symbol of the noble gas in
the previous row to indicate electrons in filled shells, and
then specify only those electrons in unfilled shells
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E.g Shorthand version of P: [Ne] 3s2 3p3
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The valence-shell electrons are the outer most shell of
electron
E.g
Valence electrons of P is 5
11.10 Electron Configurations and the
Periodic Table
 Write the full electron configuration
 short hand notation
 Determine the valence electrons
 F, Mg, As
11.11 Atomic Size
 A.
Periodicity is the
presence of regularly
repeating pattern found in
nature
 B. Atomic radius is
distance between the nuclei
of two atoms bonded
together
 C. Atomic radius increases
down a group, decreases
across a period
 i.
Larger n, larger size of
orbital
Examples
 In each of the following sets of elements, indicate which
element has the smallest atomic size
 Ba, Ca, Ra
 P, Si, Al
11.11 Ionization Energies
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Ionization energy (Ei) –
the amount of energy
required to remove the
outermost electron from an
isolated neutral atom in the
gaseous state
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Which has higher ionization
energy (Ei)?
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K or Br
S or Te