Download Chemistry Chapter 4 - Harding Charter Preparatory High School

Chapter 7
Atomic Structure
Models of the Atom
• Rutherford's model of an atom could not explain why
metals or compounds of metals give off characteristic
colors when heated in a flame, or why objects-when
heated to higher and higher temperatures-first glow dull
red, then yellow, then white
• Rutherford’s atomic model could not explain the
chemical properties of elements
Electromagnetic Radiation
• Light is electromagnetic radiation
– Radiant energy that exhibits wavelike behavior and travels
through space at the speed of light in a vacuum
Wave Length (λ)
Crest
Amplitude
Trough
𝐹𝑟𝑒𝑞𝑢𝑒𝑛𝑐𝑦 ν = 𝐶𝑦𝑐𝑙𝑒𝑠 𝑆𝑒𝑐𝑜𝑛𝑑
1 Second
4 Cycles
Electromagnetic Radiation
http://upload.wikimedia.org/wikipedia/commons/c/cf/EM_Spectrum_Properties_edit.svg
Electromagnetic Radiation
𝐶 = λν
– “C” is the speed of light (𝑪 = 𝟑. 𝟎𝟎 × 𝟏𝟎𝟖 𝒎 𝒔 )
– “λ” is wave length of the light in meters
– “ν” is the frequency of the light
Practice Problems
• Calculate the frequency of light with a
wavelength of 6.50 × 102 𝑛𝑚?
The Nature of Matter
• Matter can absorb or emit energy only in whole-number
multiples of the quantity hv where h is a constant called
Planck’s constant
𝟔. 𝟔𝟐𝟔 × 𝟏𝟎−𝟑𝟒 𝑱 ∙ 𝑺
• Change in energy for system ΔE is represented by the
equation:
∆𝑬 = 𝒉𝒗
• h is Planck’s constant
• v is frequency
– Energy is quantized and occurs only in discrete units of size hv.
Practice Problem
• What is the increment of energy (the
quantum or energy of a photon) that is
emitted at 4.50 × 102 𝑛𝑚 of light?
The Nature of Matter
E is energy
h is Planck’s constant
ν is frequency
λ is wave length
C is the speed of light
m is mass
𝑣 is velocity
𝐸𝑝ℎ𝑜𝑡𝑜𝑛
𝐸 = 𝑚𝑐 2
𝐸
𝑚= 2
𝑐
𝑚=
𝑚=
ℎ
𝜆𝑐
ℎ𝑐
= ℎν =
λ
𝐸 ℎ𝑐 λ
ℎ
=
=
𝑐2
𝑐2
λ𝑐
or if you consider a particle with velocity 𝑚 =
𝒉
𝝀=
𝒎𝒗
•
ℎ
solved for λ,
𝜆𝑣
de Broglie’s equation allows one to calculate the wavelength for all
particles
 All matter exhibits both particulate and wave properties
Practice Problem
• Compare the wavelength for an electron
(𝑚 = 9.11 × 10−31 𝑘𝑔) traveling at a speed
of 1.0 × 107 𝑚 𝑠 with that for a ball (𝑚 =
0.10𝑘𝑔) traveling at 35 m/s?
Atomic Spectrum of Hydrogen
• When H2 molecules absorb energy, some H-H
bonds are broken and the resulting hydrogen
atoms are excited
– When excited, the hydrogen atoms contain excess
energy that is released in the form of light at specific
wavelength producing an special emission spectrum
called a line spectrum
– The spectrum indicated that only certain energies are
allowed for the electron in the hydrogen atom
• In other words, the energies were quantized
Atomic Spectra
• When atoms absorb energy, electrons move into higher
energy levels, and these electrons lose energy by
emitting light when they return to lower energy levels
• Light emitted by an element separates into discrete lines
to give an atomic emission spectrum of the element
• Each discrete line in an emission spectra corresponds to
one exact frequency of light emitted by the atom
• The emission spectrum of each element is like a
person’s finger print and can be used to identify each
element
The Bohr Model
• Niels Bohr proposed that an electron is found
only in specific circular paths, or orbitals, around
the nucleus
– Each possible electron orbit in Bohr’s model has a
fixed energy.
• The fixed energies an electron can have are called energy
levels
• To move from one energy level to another, an electron must
gain or lose just the right amount of energy called a quantum
– Thus, the energy of electron is said to be quantized
– The energy levels are not equally spread, the higher energy
levels are closer together
The Bohr Model
• The electron in a hydrogen atom moves
around the nucleus only in certain allowed
circular orbits
n=5
n=4
n=3
n=2
 The orbits are known as principle energy
levels aka shells
 When not excited, the electron in the
hydrogen atom resides in the n = 1 energy
level.
 A certain quantum of energy is required for
an electron to move to a higher shell
(energy level)
n=1
The energy of the hydrogen electron in any energy level can be
calculated as shown below
Line
Spectrum
Wavelength
−𝟐. 𝟏𝟕𝟖 × 𝟏𝟎−𝟏𝟖 𝑱
𝑬=
𝒏𝟐
The change in the energy of an electron moving between
different energy levels can be calculated as shown below:
∆𝑬 = −𝟐. 𝟏𝟕𝟖 × 𝟏𝟎−𝟏𝟖 𝑱
𝟏
𝒏𝒇𝒊𝒏𝒂𝒍
𝟐−
𝟏
𝒏𝒊𝒏𝒊𝒕𝒊𝒂𝒍 𝟐
Practice Problem
• Calculate the energy required to excite the
hydrogen electron from level n = 1 to level
n = 2. Also calculate the wavelength of
light that must be absorbed by a hydrogen
atom in its ground state.
The Quantum Mechanical Model
Schrodinger
• Like the Bohr model, electrons are restricted to certain
energy levels
• Unlike the Bohr model, the quantum mechanical model
does not involve an exact path the electron takes around
the nucleus
– Electron paths are not circular
• The quantum mechanical model determines the allowed
energies an electron can have and how likely it is to find
the electron in various locations around the nucleus
Quantum Mechanical Model
• Heisenberg uncertainty principle
– There is a fundamental limitation to how
precisely both the position and momentum of
a particle can be known at a given time
Atomic Orbitals
• Atomic orbitals (wave functions) are often thought of as a
region of space in which there is a high probability of
finding an electron
– Each orbital is characterized by a series of numbers called
quantum numbers, which describe various properties of the
orbital:
• Energy levels of electrons in the quantum mechanical model are
called principle quantum numbers (n) and are assigned the
numbers 1,2,3,4, and so forth
Atomic Orbitals
• For each principle energy level (shell), there may be
several orbitals (subshells) with different shapes and
different energy levels that constitute energy sublevels
– The energy sublevels are represented by angular momentum
quantum numbers (l)
– l has integral values from 0 to n – 1
– The value of l for a particular orbital is commonly assigned a
letter:
The Angular Momentum Quantum Numbers and Corresponding Letters
Used to Designate Atomic Orbitals
Value of l
Letter Used
0
1
2
3
4
s
p
d
f
g
Atomic Orbitals
Atomic Orbitals
•
•
For each orbital (subshells) there are magnetic quantum numbers (ml) that
describe the orientation of the orbital in space relative to the other orbitals in
the atom
The (ml) has integral values between l and l -1 including zero
Quantum Number for the First Four Levels of Orbitals in the Hydrogen Atom
n
l
Orbital
Designation
1
0
1s
0
1
2
0
1
2s
2p
0
-1, 0, +1
1
3
3
0
1
2
3s
3p
3d
0
-1, 0, +1
-2, -1, 0, +1, +2
1
3
5
4
0
1
2
3
4s
4p
4d
4f
0
-1, 0, +1
-2, -1, 0, +1, +2
-3, -2, -1, 0, +1, +2, +3
1
3
5
7
ml
Number of Orbitals
Atomic Orbitals
Atomic Orbitals
• For each magnetic quantum number (ml) that
describe the orientation of an orbital in space
relative to another orbital in the atom, there are
two electron spin quantum numbers (ms)
• The (ms) can have only one of two values,
1
and −
o
o
1
+
2
1
−
2
2
is represented with a up arrow (↑)
is represented with a down arrow (↓)
1
+
2
Electron Arrangements in Atoms
• In atoms, electrons and the nucleus
interact to make the most stable
arrangement possible
• Three rules you need to know about
electron arrangements are: the Afbau
principle, the Pauli exclusion principle, and
Hund’s rule that tell how to determine the
electron arrangement of atoms
Electron Arrangements in Atoms
• Aufbau principle – electrons occupy the orbitals of lowest energy
first
– The range of some energy levels within a principle energy level can
overlap the energy levels of another principle level
• Pauli exclusion principle – an atomic orbital may describe at most
two electrons
– To occupy the same orbital, the two electrons must have opposite spin
represented with an up or down arrow ↑↓
• Hund’s rule – electrons occupy orbitals of the same energy in a way
that makes the number of electrons with the same spin direction as
large as possible
– One electron enters each orbital until all the orbitals contain one
electron with the same spin direction
• Some actual electron configurations differ from those assigned using
the aufbau principle because half-filled sublevels are not as stable
as filled sublevels, but they are more stable than other
configurations
Electron Arrangements in Atoms
1s2
All quantum values for each electron in nitrogen
2s2
3s2
4s2
5s2
6s2
n
l
ml
1s
1
0
0
+
1
2
1s
1
0
0
−
1
2
2s
2
0
0
+
1
2
2s
2
0
0
−
1
2
2p
2
1
-1
+
1
2
2p
2
1
0
+
1
2
2p
2
1
+1
+
1
2
2p6
3p6
4p6
5p6
6p6
1s
2s
0
0
3d10
4d10
4f14
5d10
5f14
6d10
6f14
2p
-1
ms
0
3s
+1
0
3p
-1
0
4s
+1
0
3d
-2
-1
0
+1 +2
• Find the electron
configuration of Zinc
1s2
1. Find Zinc’s atomic number
• Zinc has an atomic number
of 30.
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
2. Follow the arrows adding the
superscripts until you reach 30
Answer 1s2 2s2 2p6 3s2 3p6 4s2 3d10
Notice 2 + 2 + 6 + 2 + 6 + 2 + 10 = 30
3. Rearrange electron
configuration in order of
increasing energy level
Answer 1s2 2s2 2p6 3s2 3p6 3d10 4s2
• Find the electron configuration
of Ruthenium (Ru)
1s2
1. Find Ruthenium’s atomic number
•
2s2
2p6
3s2
3p6
Ruthenium has an atomic number of 44.
2. Follow the arrows adding the super scripts
until you reach 44
3d10
Answer 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d6
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
Notice 2 + 2 + 6 + 2 + 6 + 2 + 10 + 6 + 2 + 6 = 44
Notice Only 6 electrons were needed for the 4d
sublevel
3. Rearrange electron configuration in order of
increasing energy level
Answer 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d6 5s2
Valence Electrons
• Valence electrons are the electrons in the
outermost principal quantum level of an atom
– The inner electrons are known as core electrons
• Valence electrons are the most important in
determining the chemical properties of an
element
• The elements in the same group have the same
valence electron configuration
Aufbau Principle and the Periodic Table