Atoms, Energy and Electron Configurations Objectives in Terahertz ... Write the full electron configuration of Neon and Sulfur ... Atoms, Energy and Electron Configurations

Chapter 11

Atoms, Energy and Electron Configurations

1.  To review Rutherford’s model of the atom 2.  To explore the nature of electromagnetic radiation 3.  To see how atoms emit light

Objectives

Chapter 11


A. Rutherford’s Atom

…….but there is a problem here!!

Chapter 11


Using Rutherford’s Model Atoms Should Collapse….

…..then how are they so stable?

Chapter 11


When were you last exposed to electromagnetic radiation?

Chapter 11


•  What is the nature of electromagnetic radiation? •  How are the types of electromagnetic radiation different?

B. Energy and Light

–  Light can be modeled as a wave, like a wave in water

Chapter 11


•  How are the types of light different?

B. Energy and Light

–  Wavelength, λ (Greek letter “lambda”) units - m –  Frequency, ν (Greek letter “nu”) units - s-1 or Hertz (Hz)

–  Amplitude, peak height – relates to energy in wave –  Speed, c : velocity of light in a vacuum is 3x108 m.s-1

c = ν.λ ν (s-1)

(m)

Chapter 11


•  Electromagnetic radiation

B. Energy and Light

Chapter 11


Wavelength and Frequency of Visible Light

Frequency in Terahertz (1THz = 1012 Hz) and Wavelength in Nanometers (1nm = 10-9 meters)

Chapter 11


Light Sometimes Behaves in Un-wavelike Ways •  The Photoelectric Effect

–  Light shining on a metal surface can cause electrons to be separated from their atoms

–  Below a threshold frequency no electrons are emitted however high the intensity

–  At the threshold frequency electrons start to be emitted –  At higher frequencies

electrons have additional kinetic energy

Photoelectric Effect Demo

Chapter 11


•  Dual nature of light – Two co-existing models

B. Energy and Light

–  Wave –  Photon – packet of energy

Chapter 11


•  Different photons (from light of different wavelengths) carry different amounts of energy.

B. Energy and Light

Energy of photon = h.ν (h is Planck’s Constant )

Chapter 11


C. Emission of Energy by Atoms

•  Atoms can give off light –  They first must receive energy and become excited. –  The energy is released in the form of a photon.

Chapter 11


Typical Colors From Flame Tests

Na+

Li+ Cu2+

K+

Which flame represents a higher energy transition?

Periodic Table of Fireworks

Chapter 11


1.  To understand how the emission spectrum of hydrogen demonstrates the quantized nature of energy

2.  To learn about Bohr’s model of the hydrogen atom 3.  To understand how the electron’s position is

represented in the wave mechanical model

Objectives

Chapter 11


A. The Energy Levels of Hydrogen

•  Only certain types of photons are produced when excited Hydrogen atoms release energy. Why?

Chapter 11



•  Atomic states –  Excited state – atom

with excess energy –  Ground state – atom in

the lowest possible state

•  When an H atom absorbs energy from an outside source it enters an excited state.

Chapter 11



•  Energy level diagram

•  Energy in the photon corresponds to the energy used by the atom to get to the excited state.

Chapter 11



•  Only certain types of photons are produced when H atoms release energy. Why?

Chapter 11



•  Quantized Energy Levels –  Since only certain energy changes occur the H atom

must contain discrete energy levels.

Chapter 11


B. The Niels Bohr Model of the Atom (1913-1915)

•  Bohr’s model of the atom –  Quantized energy levels –  Electron moves in a circular orbit –  Electrons jump between levels by absorbing or

emitting photons of a particular wavelength

Able to mathematically explain the emission spectrum of Hydrogen (compare with Rutherford)

Chapter 11


B. The Bohr Model of the Atom

•  Bohr’s model of the atom was not totally correct. –  Had difficulties with spectra of larger atoms –  Electrons do not move in a circular orbit and don’t

seem to behave like discrete particles all the time

Maybe small particles, such as electrons, can also behave like waves having a dual nature (just like photons)……?

Chapter 11


C. The Wave Mechanical Model of the Atom (de Broglie and Schroedinger – mid-1920’s)

•  Orbitals –  Nothing like orbits –  Probability of finding the electron within a certain space

Bohr-Schrodinger

Chapter 11


1.  To learn about the shapes of the s, p and d orbitals 2.  To review the energy levels and orbitals of the wave

mechanical model of the atom 3.  To learn about electron spin

Objectives

Chapter 11


A. The Hydrogen Orbitals

•  Orbitals do not have sharp boundaries.

90% boundary

Chapter 11



•  Hydrogen has discrete energy levels. •  Called principal energy

levels (electron shells) •  Labeled with whole numbers •  Energy is related to 1/n2

•  En = E1/n2 •  Energy levels are closer

together the further they are from the nucleus

Hydrogen Energy Levels

Chapter 11



•  Each principal energy level is divided into sublevels.


–  Labeled with numbers and letters –  Indicate the shape of the orbital

Chapter 11


Orbital Designations

Chapter 11


Orbitals Define the Periodic Table

“Orbitals” are “Energy Levels”

Chapter 11



•  The s and p types of sublevel


Chapter 11


Representation of s, p, d atomic orbitals

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•  Why does an H atom have so many orbitals and only 1 electron? –  An orbital is a potential space for an electron. –  Atoms can have many potential orbitals.

•  s, p, d, f orbitals named for sharp, principal, diffuse and fundamental lines on spectra. Further orbitals designated alphabetically

Hydrogen Orbitals

Chapter 11


g g g g

f f f

s p d d

Orbitals

Chapter 11


•  Close examination of spectra revealed doublets •  Need one more property to determine how electrons

are arranged •  Spin – electron modeled as a spinning like a top •  Spin is the basis of magnetism

B. The Wave Mechanical Model: Further Development

Electron Spin

Chapter 11


•  Pauli Exclusion Principle (Wolfgang Pauli 1925) - an atomic orbital can hold a maximum of 2 electrons and those 2 electrons must have opposite spins

•  When an orbital contains two electrons (of opposite spin) it is said to be full

B. The Wave Mechanical Model: Further Development

Pauli Exclusion Principle

What are the four descriptors that define an energy level / electron’s position in an atom?

Chapter 11


1.  To understand how the principal energy levels fill with electrons in atoms beyond hydrogen

2.  To learn about valence electrons and core electrons 3.  To learn about the electron configurations of atoms 4.  To understand the general trends in properties in the

periodic table

Objectives

Chapter 11


A. Electron Arrangements in the First 18 Atoms on the Periodic Table •  H atom

–  Electron configuration – electron arrangement – 1s1 –  Orbital diagram – orbital is represented as a box with a

designation according to its sublevel. Contains arrow(s) to represent electrons (spin)

Chapter 11


A. Electron Arrangements in the First 18 Atoms on the Periodic Table

•  He atom –  Electron configuration – 1s2 –  Orbital diagram

Chapter 11



•  Li atom –  Electron configuration– 1s2 2s1 –  Orbital diagram

Write the electron configuration and orbital diagrams for Boron, Nitrogen, Fluorine and Argon

Chapter 11



Write the full electron configuration of Neon and Sulfur

Draw an orbital diagram for Magnesium and Chlorine

Chapter 11


Order of Filling of Orbitals

Atoms fill their orbitals in the order of their energies:

Chapter 11


•  Orbital filling and the periodic table

B. Electron Configurations and the Periodic Table Dynamic Periodic Table

PT 2a PT 2b

Chapter 11


•  Electron configurations for K through Kr

B. Electron Configurations and the Periodic Table

Chapter 11


If there were more elements………. B. Electron Configurations and the Periodic Table

Chapter 11



Classifying Electrons

•  Valence electrons – electrons in the outermost (highest) principal energy level of an atom

•  Core electrons – inner electrons •  Elements with the same valence electron arrangement

(same group) show very similar chemical behavior.

Chapter 11


B. Electron Configurations and the Periodic Table

Chapter 11


Using a Noble Gas Shorthand

•  We can abbreviate electron configurations by using the configuration of the previous noble gas to cover the first part of the list of orbitals

•  Mg is 1s2 2s2 2p6 3s2 or [Ne] 3s2

•  The noble gas portion is the equivalent to the group of core electrons

•  Use the Noble Gas shorthand to show the electron configurations of Carbon, Chlorine and Zirconium

Chapter 11


C. Atomic Properties and the Periodic Table

Metals and Nonmetals •  Metals tend to lose electrons to form positive ions. •  Nonmetals tend to gain electrons to form negative ions.

Chapter 11



Atomic Size •  Size tends to increase down a column. •  Size tends to decrease across a row.

(close to scale) Larger

Effects of Shielding of outer electrons by inner orbitals

Chapter 11



Ionization Energies •  Ionization Energy – energy (ΔH) required to remove an

electron from an individual atom (gas)

–  Tends to decrease down a column –  Tends to increase across a row –  Changes in an opposite direction to atomic size

Chapter 11


Ionization Energies

Chapter 11


Electron Affinity

•  Electron Affinity is defined as ΔH for the process: X(g) + e- = X(g)

-

ΔH = Electron Affinity Larger

Chapter 11


Electronegativity

•  Ionization Energy and Electron Affinity are combined to give Electronegativity – a measure of how well atoms compete for electrons in a bond

Larger

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Atoms, Energy and Electron Configurations Objectives in Terahertz ... Write the full electron configuration of Neon and Sulfur ... Atoms, Energy and Electron Configurations