chap 03b-modern atomic theory - los angeles mission college · 2015. 3. 13. · bohr’s model of...

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Chapter 3B

Modern Atomic

Theory

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CHAPTER OUTLINE��

§  Waves §  Electromagnetic Radiation §  Dual Nature of Light §  Bohr Model of Atom §  Quantum Mechanical Model of Atom §  Electron Configuration §  Electron Configuration & Periodic Table §  Abbreviated Electron Configuration

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wavelength (measured from

peak to peak) wavelength

(measured from trough to trough)

10.1

WAVES��

q  All waves are characterized by wavelength, frequency and speed.

q  Wavelength (λ) is the distance between any 2 successive crests or troughs.

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10.1

WAVES��

q  Frequency (nu,ν) is the number of waves produced per unit time.

q  Wavelength and frequency are inversely proportional.

q  Speed tells how fast waves travel through space.

As wavelength of a wave increases

its frequency decreases inversely

proportional

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ELECTROMAGNETIC��RADIATION��

q  Energy travels through space as electromagnetic radiation. This radiation takes many forms, such as sunlight, microwaves, radio waves, etc.

q  In vacuum, all electromagnetic waves travel at the speed of light (3.00 x 108 m/s), and differ from each other in their frequency and wavelength.

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q  The classification of electromagnetic waves according to their frequency is called electromagnetic spectrum.

q  These waves range from γ-rays (short λ, high f) to radio waves (long λ, low f). Short

wavelength High frequency

Long wavelength

Low frequency

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10.2

Visible light is a small part of the EM

spectrum X-rays have longer λ but lower υ than

γ-rays

Infrared waves have longer λ but lower υ

than visible light


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DUAL NATURE��OF LIGHT��

q  When white light is passed through a glass prism, it is dispersed into a spectrum of colors.

q  This is evidence of the wave nature of light.

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q  Scientists also have much evidence that light beams act as a stream of tiny particles, called photons.

A photon of red light

A photon of blue light

Red light has longer wavelength and less energy than blue light

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q  Scientists, therefore, use both the wave and particle models for explaining light. This is referred to as the wave-particle nature of light.

q  Flatland Video q  Scientists also discovered that when atoms are

energized at high temperatures or by high voltage, they can radiate light. Neon lights are an example of this property of atoms.

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ATOMIC LINE��SPECTRUM��

q  When the light from the atom is placed through a prism, a series of brightly colored lights, called a line spectrum is formed.

q  These lines indicate that light is formed only at certain wavelengths and frequencies that correspond to specific colors.

q  Each element possesses a unique line spectrum that can be used to identify it.

Each line represents a particular λ and

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BOHR MODEL��OF ATOM��

q  Neils Bohr, a Danish physicist, studied the hydrogen atom extensively, and developed a model for the atom that was able to explain the line spectrum.

q  Bohr’s model of the atom consisted of electrons orbiting the nucleus at different distances from the nucleus, called energy levels.

q  In this model, the electrons could only occupy particular energy levels, and could “jump” to higher levels by absorbing energy.

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q  The lowest energy level is called ground state, and the higher energy levels are called excited states.

q  When electrons absorb energy through heating or electricity, they move to higher energy levels and become excited.

energy

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q  When excited electrons return to the ground state, energy is emitted as a photon of light is released.

q  The color (wavelength) of the light emitted is determined by the difference in energy between the two states (excited and ground).

Lower energy transition

give off red light

Higher energy transition give off blue light

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q  The line spectrum is produced by many of these transitions between excited and ground states.

q  Bohr’s model was able to successfully explain the hydrogen atom, but could not be applied to larger atoms.

q  Quantum Mechanics & Structure of Atom

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QUANTUM MECHANICAL��MODEL OF ATOM��

q  In 1926 Erwin Shrödinger created a mathematical model that showed electrons as both particles and waves. This model was called the quantum mechanical model.

q  Double-Slit Experiment q  This model predicted electrons to be located in a

probability region called orbitals. q  An orbital is defined as a region around the

nucleus where there is a high probability of finding an electron.

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q  Based on this model, there are discrete principal energy levels within the atom.

q  Principal energy levels are designated by n.

q  The electrons in an atom can exist in any principal energy level.

As n increases, the energy of the

electron increases

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10.7, 10.8


q  Each principal energy level is subdivided into sublevels.

q  The sublevels are designated by the letters s, p, d and f.

q  As n increases, the number of sublevels increases.

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q  Within the sublevels, the electrons are located in orbitals. The orbitals are also designated by the letters s, p, d and f.

q  The number of orbitals within the sublevels vary with their type.

s sublevel = 1 orbital

p sublevel = 3 orbitals

d sublevel = 5 orbitals

f sublevel = 7 orbitals

An orbital can hold a maximum of 2 electrons

= 2 electrons

= 6 electrons

= 10 electrons

= 14 electrons

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ELECTRON��CONFIGURATION��

q  Similarities of behavior in the periodic table are due to the similarities in the electron arrangement of the atoms. This arrangement is called electron configuration.

q  The modern model of the atom describes the electron cloud consisting of separate energy levels, each containing a fixed number of electrons.

q  Each orbital can be occupied by no more than 2 electrons, each with opposite spins.

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q  The electrons occupy the orbitals from the lowest energy level to the highest level.

q  The energy of the orbitals on any level are in the following order: s < p < d < f.

q  Each orbital on a sublevel must be occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital (Hund’s Rule).

q  Visualizing Orbitals

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q  Electron configurations can be written as:

2 p6 Principal

energy level Type of orbital

Number of electrons in

orbitals

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q  Another notation, called the orbital notation is shown below:

1 s

Principal energy level

Type of orbital

Electrons in orbital with

opposing spins

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↑ 1s2

H ↑ 1s1

Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s.

He

Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins.

↓

1s

1s


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Li

1s22s2

The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital.

↑ ↓ ↑

↓

1s

↑

1s22s1

2s

Be ↑ ↓

The 2s orbital fills upon the addition of beryllium’s third and fourth electrons.

1s 2s


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B 1s22s22p1 ↑ ↓ ↑ 1s

2p

Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first.

C

The second p electron of carbon enters a different p orbital than the first p due to Hund’s Rule.

1s22s22p2 ↑

2s

↑ ↓

1s

↑ ↓ 2s

↑ ↓

2p

↑ ↓


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N 1s22s22p3 ↑ ↓ ↑ 1s

2p

The third p electron of nitrogen enters a different p orbital than its first two p electrons due to Hund’s Rule.

O

The last p electron of oxygen pairs opposite of another since each orbital has an electron in it and Hund’s Rule is satisfied.

1s22s22p4 ↑

2s

↑ ↓

1s

↑ ↓ 2s

↑ ↓

2p

↑ ↓


↑ ↑

↑

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F 1s22s22p5 ↑ ↓ ↑ 1s

2p

Two of the p electrons for fluorine pair up with other electrons in the p orbitals.

Ne

The last p electron for neon pairs up with the last lone electron and completely fills the 2nd energy level.

1s22s22p6 ↑

2s

↑ ↓

1s

↑ ↓ 2s

↑ ↓

2p

↑ ↓


↑ ↑

↑

↓ ↓

↓ ↓

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ELECTRON ��CONFIGURATION��

q  As electrons occupy the 3rd energy level and higher, some anomalies occur in the order of the energy of the orbitals.

q  Knowledge of these anomalies is important in order to determine the correct electron configuration for the atoms.

q  The following study aid is used by beginning students to remember these exceptions to the order of orbital energies.

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ELECTRON ��CONFIGURATION��

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s

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ELECTRON CONFIG. ��& PERIODIC TABLE��

q  The horizontal rows in the periodic table are called periods. The period number corresponds to the number of energy levels that are occupied in that atom.

q  The vertical columns in the periodic table are called groups or families. For the main-group elements, the group number corresponds to the number of electrons in the outermost filled energy level (valence electrons).

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One energy level

3 energy levels

4 energy levels

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1 valence electron

5 valence electrons

3 valence electrons

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10.15

q  The valence electrons configuration for the elements in periods 1-3 are shown below.

q  Note that elements in the same group have similar electron configurations.


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10.16


Arrangement of orbitals in the periodic table

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10.16


d orbital numbers are 1 less than the period number

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10.16


f orbital numbers are 2 less than the period number

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10.15

q  The electrons in an atom fill from the lowest to the highest orbitals.

q  The knowledge of the location of the orbitals on the periodic table can greatly help the writing of electron configurations for large atoms.

q  The energy order of the sublevels is shown next. Note that some anomalies occur in the energy level of “d” and “f” sublevels.


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Example 1:��Use the periodic table to write complete electron configuration for phosphorus.

P Z = 15

1s2 2s2 2p6 3s2 3p3

10 electrons used

5 electrons remaining

Core electrons

Valence electrons

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Example 2:��Draw an orbital notation diagram for the last in-complete level of chlorine and determine the number of unpaired electrons.

3s 3p

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Cl ↑ ↓ ↑ 3s 3p

↑ ↑

Example 2:��

↓ ↓

One unpaired electron

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ABBREVIATED��ELECTRON CONFIG. ��

q  When writing electron configurations for larger atoms, an abbreviated configuration is used.

q  In writing this configuration, the non-valence (core) electrons are summarized by writing the symbol of the noble gas prior to the element in brackets followed by configuration of the valence electrons.

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K Z = 19

1s22s22p63s23p6 4s1

core electrons

valence electron [Ar] 4s1

Previous noble gas

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Br Z = 35

1s22s22p63s23p6 4s2

core electrons

valence electrons [Ar] 4s23d104p5

3d10 4p5

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Example 3:��Write abbreviated electron configurations for each element listed below:

Fe Z = 26

4s2 3d6 [Ar]

18 electrons used


20 electrons used


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Example 3:��Write abbreviated electron configurations for each element listed below:

Sb Z = 51

5s2 4d10 [Kr]

36 electrons used


38 electrons used


48 electrons used


5p3

5 valence electrons

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TRENDS IN��PERIODIC PROPERTIES ��

q  The electron configuration of atoms are an important factor in the physical and chemical properties of the elements.

q  Some of these properties include: atomic size, ionization energy and metallic character.

q  These properties are commonly known as periodic properties and increase or decrease across a period or group, and are repeated in each successive period or group.

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ATOMIC SIZE ��

q  The size of the atom is determined by its atomic radius, which is the distance of the valence electron from the nucleus.

q  For each group of the representative elements, the atomic size increases going down the group, because the valence electrons from each energy level are further from the nucleus.

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ATOMIC SIZE ��

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ATOMIC SIZE ��

q  The atomic radius of the representative elements are affected by the number of protons in the nucleus (nuclear charge).

q  For elements going across a period, the atomic size decreases because the increased nuclear charge of each atom pulls the electrons closer to the nucleus, making it smaller.

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ATOMIC SIZE ��

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IONIZATION��ENERGY ��

q  The ionization energy is the energy required to remove a valence electron from the atom in a gaseous state.

q  When an electron is removed from an atom, a cation (+ ion) with a 1+ charge is formed.

Na (g) + IE Na+ + e-

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q  The ionization energy decreases going down a group, because less energy is required to remove an electron from the outer shell since it is further from the nucleus.

Larger atom Less IE

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q  Going across a period, the ionization energy increases because the increased nuclear charge of the atom holds the valence electrons more tightly and therefore it is more difficult to remove.

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q  In general, the ionization energy is low for metals and high for non-metals.

q  Review of ionization energies of elements in periods 2-4 indicate some anomalies to the general increasing trend.

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q  These anomalies are caused by more stable electron configurations of the atoms in groups 2 (complete “s” sublevel) and group 5 (half-filled “p” sublevels) that cause an increase in their ionization energy compared to the next element.

Be 1s2 2s2

B 1s2 2s2 2p1

More stable Higher IE

N 1s2 2s2 2p3

O 1s2 2s2 2p4

More stable (1/2 filled) Higher IE

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METALLIC��CHARACTER��

q  Metallic character is the ability of an atom to lose electrons easily.

q  This character is more prevalent in the elements on the left side of the periodic table (metals), and decreases going across a period and increases for elements going down a group.

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METALLIC��CHARACTER��

Most metallic elements

Least metallic elements

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Example 1:��Select the element in each pair with the larger atomic radius:

Li K or

Larger due to more

energy levels

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K Br or

Larger due to less nuclear

charge

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P Cl or

Larger due to less nuclear

charge

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Example 2:��Indicate the element in each set that has the higher ionization energy and explain your choice:

K Na or

Higher IE due to less

energy levels

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Example 2:��Indicate the element in each set that has the higher ionization energy and explain your choice:

Mg Cl or

Higher IE due to more

nuclear charge

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Example 2:��

F C or

Highest IE due to most

nuclear charge

Indicate the element in each set that has the higher ionization energy and explain your choice:

N

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THE END

chap 03b-modern atomic theory - los angeles mission college · 2015. 3. 13. · bohr’s model of...

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