# Chapter 4 - Arrangement of Electrons in Atoms by r5uhfBp

VIEWS: 27 PAGES: 12

• pg 1
```									                           Unit 7.1 - Arrangement of Electrons in Atoms

The “Puzzle” of the nucleus:
• Protons and electrons are attracted to each other because of opposite charges
• Electrically charged particles moving in a curved path give off energy
• Despite these facts, atoms don’t collapse

7.1-1 The Development of a New Atomic Model

I.        Properties of Light
1. Many types of EM waves a. visible
light
b. x-rays
c. ultraviolet light d.
infrared light
2. EM radiation are forms of energy which move through space as waves a. Move at speed
of light
8
(1). 3.00 x 10 m/s
b. Speed is equal to the frequency times the wavelength                    c = νλ
(1). Freqency (ν) is the number of waves passing a given point in one
second
(2). Wavelength (λ) is the distance between peaks of adjacent
waves

c. Speed of light is a constant, so νλ is also a constant
(1) ν and λ must be inversely proportional

Equations:                 E = hν    c=λν

1.    Determine the frequency of a quantum of blue light that has a wavelength of 492 nm.
2. Calculate the speed of a photon of red light that has a wavelength of 725 nm and a frequency of
4.08 x 1014 Hz.

3. Calculate the wavelength (in nm) of a photon of green light that has a frequency of 5.5 x 10 14 Hz.

4. Calculate the energy of a quantum of ultraviolet light that has a frequency of 1.2 x 1015 Hz.

5. Determine the energy of a photon of light that has a wavelength of 435 nm.

B. Light and Energy - The Photoelectric Effect
1. The Photoelectric Effect
a. Electrons are emitted from a metal when light shines on the metal
2. Radiant energy is transferred in units (or quanta) of energy called photons
a. A photon is a particle of energy having a rest mass of zero and carrying a
quantum of energy
b. A quantum is the minimum amount of energy that can be lost or gained by an
atom
3. Energy of a photon is directly proportional to the frequency of radiation a. E = hν
-27
(h is Planck’s constant, 6.62554 x 10        erg sec)

-12             -10            -8               -7        -4         -2                     2        4
10              10             10        4 to 7x10        10         10        1            10       10
gamma          xrays          UV          visible         IR        micro           Radio waves

FM       short    AM

Wavelength         increases
Frequency         decreases
Energy            decreases

4. Wave-Particle Duality
a. Energy travels through space as waves, but can be thought of as a stream of
particles (Einstein)

II.      The Hydrogen Line Spectrum
A. Ground State
1. The lowest energy state of an atom
B. Excited State
1. A state in which an atom has a higher potential energy than in its ground state
C. Bright line spectrum
1. Light is given off by excited atoms as they return to lower energy states
2. Light is given off in very definite wavelengths
3. A spectroscope reveals lines of particular colors

410nm     434nm      486nm                                               656nm
III.    The Bohr Model of the Atom

Niels Bohr (1913): Danish physicist. Bohr modified Rutherford’s model by suggesting that
electrons can only possess certain amounts of energy.
What does this mean in terms of the location of electrons?
 they can only be at certain distances from the nucleus

   Bohr received the Nobel Prize in 1922 for his Bohr model, or planetary model.
Bohr’s work was the forerunner for the work of many other individuals who, by the
1930’s and 1940’s, had modified Bohr’s model into the charge-cloud model, or
quantum mechanical model.

The quantum mechanical model of the atom is the currently-accepted model. It falls within the field
of physics called Quantum Mechanics which is the idea that energy is quantized = energy has
only certain allowable values; other values are NOT allowed

In an atom, where are the electrons, according to the quantum mechanical model?
we cannot say for sure, but the equations of Quantum Mechanics can tell us the probability that we
will find an electron at a certain distance from the nucleus

A. Electron Orbits, or Energy Levels
1. Electrons can circle the nucleus only in allowed paths or orbits
2. The energy of the electron is greater when it is in orbits farther from the
nucleus
3. The atom achieves the ground state when atoms occupy the closest
possible positions around the nucleus
4. Electromagnetic radiation is emitted when electrons move closer to the
nucleus
B. Energy transitions
1. Energies of atoms are fixed and definite quantities
2. Energy transitions occur in jumps of discrete amounts of energy
3. Electrons only lose energy when they move to a lower energy state

C. Shortcomings of the Bohr Model
1. Doesn't work for atoms larger than hydrogen (more than one electron)
2. Doesn't explain chemical behavior

7.1-2 The Quantum Model of the Atom
I.      Electrons as Waves and Particles
A. Louis deBroglie (1924)
1. Electrons have wavelike properties
2. Consider the electron as a wave confined to a space that can have only
certain frequencies
B. The Heisenbery Uncertainty Principle (Werner Heisenberg - 1927)
1. "It is impossible to determine simultaneously both the position and
velocity of an electron or any other particle
a. Electrons are located by their interactions with photons b.
Electrons and photons have similar energies
c. Interaction between a photon and an electron knocks the
electron off of its course
C. The Schroedinger Wave Equation
1. Proved quantization of electron energies and is the basis for Quantum
Theory
a. Quantum theory describes mathematically the wave properties of
electrons and other very small particles
2. Electrons do not move around the nucleus in "planetary orbits"
3. Electrons exist in regions called orbitals
a. An orbital is a three-dimensional region around the nucleus that
indicates the probable location of an electron.
*Schrödinger equation (YOU DO NOT NEED TO KNOW THIS
EQUATION) for probability of a single electron being found along a
single axis (x-axis)

II.       Atomic Orbitals and Quantum Numbers

A Closer Look at Electrons: Where are they in the Atom?

Electrons are located within energy levels, which range from 1 to 7. The higher the energy
level the electron is in…

1. the farther the electron is from the nucleus

2. the more energy the electron has

Within each energy level, there exist sublevels, which differ from each other by slight
differences in energy. In each sublevel there are “paths”, called orbitals, that an electron can
travel on.

orbital = a region of an atom in which there is a high probability of finding electrons
Quantum Numbers specify the properties of atomic orbitals and the properties of the
electrons in orbitals

A. Principal Quantum Number (n)
1. Indicates the main energy levels occupied by the electron
2. Values of n are positive integers
a. n=1 is closest to the nucleus, and lowest in energy
3. The number of orbitals possible per energy level (or "shell") is equal to 2
n
B. Angular Momentum Quantum Number (l)
1. Indicates the shape of the orbital
2. Number of orbital shapes = n
a. Shapes are designated s, p, d, f
C. Magnetic Quantum Number (m)
1. The orientation of the orbital around the nucleus
a. s orbitals have only one possible orientation
m= 0
b. p orbitals have three, d have five and f have 7 possible
orientations

s orbital               px orbital             py orbital            pz orbital

2 2                 2
dxy orbital        dxz orbital         dyz orbital      dx - orbital        dz orbital
Each orbital can hold a maximum of ____ electrons.
    In every s sublevel, there is ____ orbital, which holds a total of ___ electrons
    In every p sublevel, there are ____ orbitals, which hold a total of ___ electrons
    In every d sublevel, there are ____ orbitals, which hold a total of ___ electrons
    In every f sublevel, there are ____ orbitals, which hold a total of ___ electrons

Principal               Sublevels     in   Number      of    Number of           Number of
Quantum                 main     energy    orbitals   per    electrons per       electrons per
Number (n)              level              sublevel          sublevel            main energy
2
(n sublevels)                                            level (2n )
1                          s                 1                  2                 2
2                          s                 1                  2                 8
p                 3                  6
3                          s                 1                   2               18
p                 3                   6
d                 5                  10
4                          s                 1                   2               32
p                 3                   6
d                 5                  10
7                  14
f

D. Spin Quantum Number
1. Indicates the fundamental spin states of an electron in an orbital
2. Two possible values for spin, +1/2, -1/2
3. A single orbital can contain only two electrons, which must have opposite
spins

7-3 Electron Configurations
I.      Writing Electrons Configurations
The question is: Where are the electrons in the atom?
2
The format for the electron configuration is, for example: 1 s
1 = the energy level
s = the sublevel, or orbital
2 = the number of electrons in that sublevel

A. How to Write an Electron Configuration
1. Locate the element on the periodic table.
2. Fill the orbitals in the proper order.
3. Check that the total number of electrons you have equals the atomic number for that element.

Examples: Write the electron configurations for the following elements.
carbon (C)

lithium (Li)

sodium (Na)
chlorine (Cl)

potassium (K)

iron (Fe)

B. Using Shorthand Notation for the Electron Configuration
Put the noble gas that precedes the element in brackets, then continue filling the rest
of the orbitals in order, as usual.
Examples:
sodium (Na)

chlorine (Cl)

potassium (K)

iron (Fe)

C. Rules that apply to orbital notation
1. Aufbau Principle
a. An electron occupies the lowest-energy orbital that can receive it
2. Pauli Exclusion Principle
a. No two electrons in the same atom can have the same set of
four quantum numbers
3. Hund's Rule
a. Orbitals of equal energy are each occupied by one electron
before any orbital is occupied by a second electron, and all
electrons in singly occupied orbitals must have the same spin

2p                                    2p                             2p

4. Orbital diagrams
a). Unoccupied orbitals are represented by a line,
i. Lines are labeled with the principal quantum number and the
sublevel letter
b). Arrows are used to represent electrons
i. Arrows pointing up and down indicate opposite spins

Examples: Write the orbital diagrams for the following elements.
carbon (C)

lithium (Li)

sodium (Na)
chlorine (Cl)

potassium (K)

iron (Fe)

B. Elements of the Fourth Period
1. Irregularity of Chromium
2 2 6 2 6 2 4
a. Expected: 1s 2s 2p 3s 3p 4s 3d
2 2 6 2 6 1 5
b. Actual: 1s 2s 2p 3s 3p 4s 3d
2. Several transition and rare-earth elements borrow from smaller sublevels in
order to half fill larger sublevels

Unit 7.2 - The Periodic Law

7.2-1 History of the Periodic Table
I.      Mendeleev's
Periodic Table
A. Organization
1. Vertical columns in atomic weight order
a. Mendeleev made some exceptions to place elements in rows
with similar properties
(1) Tellurium and iodine's places were switched
2. Horizontal rows have similar chemical properties
B. Missing Elements
1. Gaps existed in Mendeleev’s table
a. Mendeleev predicted the properties of the “yet to be discovered”
elements
(1) Scandium, germanium and gallium agreed with predictions
1. Why didn't some elements fit in order of increasing atomic mass?
2. Why did elements exhibit periodic behavior?

II.     Moseley and the Periodic Table
(1911) A. Protons and Atomic Number
1. Xray experiments revealed a way to determine the number of protons in the
nucleus of an atom
2. The periodic table was found to be in atomic number order, not atomic mass
order
a. The tellurium-iodine anomaly was explained
B. The Periodic Law
1. The physical and chemical properties of the elements are periodic functions
of their atomic numbers
***Moseley was killed in battle in 1915, during WWI. He was 28 years old

III.        The Modern
Periodic Table
A. Discovery of the Noble Gases
1. 1868 - Helium discovered as a component of the sun, based on the
emission spectrum of sunlight
2. 1894 - William Ramsay discovers argon
3. 1895 - Ramsay finds helium on Earth
4. 1898 - Ramsay discovers krypton and xenon
5. 1900 - Freidrich Dorn discovers radon
B. The Lanthanides
1. Early 1900's the elements from cerium (#58) to lutetium (#71) are separated
and identified
C. The Actinides
1. Discovery (or synthesis) of elements 90 to 103

D. Periodicity
1. Elements with similar properties are found at regular intervals within the
"periodic" table

7.2-2 Electron Configuration and the Periodic Table
I.      Periods and the Blocks of the Periodic Table
A. Periods
1. Horizontal rows on the periodic table
2. Period number corresponds to the highest principal quantum number of the
elements in the period
B. Sublevel Blocks
1. Periodic table can be broken into blocks corresponding to s, p, d, f sublevels

II.     Blocks and Groups
A. s-Block, Groups 1 and 2
1. Group 1 - The alkali metals
a. One s electron in outer shell
b. Soft, silvery metals of low density and low melting
points c. Highly reactive, never found pure in nature
2. Group 2 - The alkaline earth
metals a. Two s
electrons in outer shell
b. Denser, harder, stronger, less reactive than
Group 1 c. Too reactive to be found pure in nature
B. d-Block, Groups 3 - 12
1. Metals with typical metallic properties
2. Referred to as "transition" metals
3. Group number = sum of outermost s and d electrons
C. p-Block elements, Groups 13 - 18
1. Properties
vary
greatly
a.
Metals
(1) softer and less dense than d-block metals
(2) harder and more dense than s-block
metals b. Metalloids
(1) Brittle solids with some metallic and some
nonmetallic properties
(2)
Semiconductors
c. Nonmetals
(1) Halogens (Group 17) are most reactive of the nonmetals
D. f-Block, Lanthanides and Actinides
1. Lanthanides are shiny metals similar in reactivity to the Group 2 metals
2. Actinides
b. Plutonium (94) through Lawrencium (103) are man-made
7.2-3 Electron Configuration and Periodic Properties
1. One half the distance between nuclei of identical atoms that are bonded
together
B. Trends
1. Atomic radius tends to decrease across a period due to increasing positive
nuclear charge
2. Atomic radii tend to increase down a group due to increasing number energy
levels (outer electrons are farther from the nucleus)

II.         Ionization Energy
A. Ion
1. An atom or a group of atoms that has a positive or negative charge
B. Ionization
1. Any process that results in the formation of an ion
C. Ionization Energy
1. The energy required to remove one electron from a neutral atom of an
element, measured in kilojoules/mole (kJ/mol)
+      -
A + energy →            A    + e
D. Trends
1. Ionization energy of main-group elements tends to increase across each
period
a. Atoms are getting smaller, electrons are closer to the nucleus
2. Ionization energy of main-group elements tends to decrease as atomic
number increases in a group
a. Atoms are getting larger, electrons are farther from the nucleus
b. Outer electrons become increasingly more shielded from the
nucleus by inner electrons
3. Metals have a characteristic low ionization energy
4. Nonmetals have a high ionization energy
5. Noble gases have a very high ionization energy

+             -
Na + 496 kJ/mol→Na           + e

+                  ++    -
Na + 4562 kJ/mol → Na   + e

++                    +++    -
Na   + 6912 kJ/mol → Na     + e

1. Ionization energy increases for each successive electron
2. Each electron removed experiences a stronger effective nuclear charge
3. The greatest increase in ionization energy comes when trying to remove an
electron from a stable, noble gas configuration

III.    Electron Affinity
A. Electron Affinity
1. The energy change that occurs when an electron is acquired by a
neutral atom, measured in kJ/mol
a. Most atoms release energy when they acquire an electron

-      -
A + e    → A       + energy (exothermic)

b. Some atoms must be forced to gain an electron
-                            -
A    + e   +       energy →        A endothermic)
B. Trends
1. Halogens have the highest electron affinities
2. Metals have characteristically low electron affinities
3. Affinity tends to increase across a period
a. Irregularities are due to the extra stability of half-filled and
filled sublevels
4. Electron affinity tends to decrease down a group
5. Second electron affinities are always positive (endothermic)

A. Cations
1. Positive ions
2. Smaller than the
corresponding atom a.
Protons outnumber
electrons b. Less
shielding of electrons
B. Anions
1. Negative ions
2. Larger than the
corresponding atoms
a. Electrons
outnumber protons
b. Greater electron-electron repulsion
C. Trends
1. Ion size tends to increase downward within a group

V.Valence Electrons
A. Valence Electrons
1. The electrons available to be lost, gained, or shared in the
formation of chemical compounds
2. Main group element valence electrons are outermost energy level s and p
sublevels

Group #                                    1         2        13         14         15         16     17   18
Number of valence Electrons                1         2         3          4          5          6      7    8

VI.     Electronegativity
A. Electronegativity
1. A measure of the ability of an atom in a chemical compound to attract
electrons
2. Elements that do not form compounds are not assigned electronegativities
B. Trends
1. Nonmetals have characteristically high
electronegativity a. Highest in the upper right
corner
2. Metals have characteristically low
electronegativity a. Lowest in the lower left
corner of the table
3. Electronegativity tends to increase across a period
4. Electronegativity tends to decrease down a group of main-group elements
VII.   Periodic Properties of the d- and f- Block Elements
1. Smaller decrease in radius across a period within the d- Block than within
the main-group elements
a. Added electrons are partially shielded from the increasing positive
nuclear charge
b. Slight increase at the end of the d-Block is due to electron-electron
repulsion
2. Little change occurs in radius across an f-block of elements
B. Ionization Energy
1. Tends to increase across d- and f-Blocks
C. Ion Formation and Ionic Radii
1. Electrons are removed from the outermost energy level s-sublevel first a.
Most d-block elements form 2+ ions (losing 2 s electrons)
2. Ions of d- and f-Blocks are cations, smaller than the corresponding atoms
D. Electronegativity
1. Characteristically low electronegativity of metals
2. Electronegativity increases as atomic radius decreases

```
To top