Chapter 11

11.1 Rutherford’s Atom: To describe Rutherford’s model of the atom.11.2 Energy and Light: To explore the nature of electromagnetic radiation.11.3: Emission of Energy by Atoms: To see how atoms emit light.

11.1 Rutherford’s Atom

Gold Foil experiment: small dense nucleus containing protons and neutrons

QUESTIONS:What are the electrons doing?How are the electrons arranged and how do

they move?

Rutherford suggested that electrons might revolve around the nucleus.

11.2: Energy and Light

Electromagnetic Radiation: (energy transmitted as light).

Figure 11.4 Rays X rays UV Visible Infrared Micro Radio

Wavelength 10-12 to 104

Wave: wavelength() frequency(and speed ( how fast)

11.2: Energy and LightLight can behave as waves

or photons (stream of tiney packets of energy).

Wave-particle nature of light.

Red light- lower energy. Longer the wavelength the lower the energy of its photons.

11.3: Emission of Energy by Atoms

Atoms emit light of different energy.Heat from the flame allows atoms to absorb

energy-we say that the atoms become excited. Some of this excess energy is released as light.

11.3: Emission of Energy by Atoms

Li+ : gives off a beautiful, deep-red color(emits photons of red light)Cu2+: burns greenNa+: burns yellow-orange color

Excited Li atom

Photon of redlight emitted

Li atom in lowerEnergy state

Reminder: High energy (short Wavelengths)Red light has long wavelength, lowerenergy

11.4: The Energy Levels of Hydrogen

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

Excited Li atom

Photon of redlight emitted

Li atom in lowerEnergy state

11.4: The Energy Levels of Hydrogen

An atom with excess energy is called the EXCITED STATE

The lowest possible enery state of an atom is called the GROUND STATE

ENERGY contained in the photon corresponds to the change in energy thatThe atom experiences in going from the excited state to the lower state.

Excited Li atom

Photon of redlight emitted

Li atom in lowerEnergy state

11.4: The Energy Levels of Hydrogen

An atom with excess energy is called the EXCITED STATE

The lowest possible enery state of an atom is called the GROUND STATE

ENERGY contained in the photon corresponds to the change in energy thatThe atom experiences in going from the excited state to the lower state.

11.4: The Energy Levels of Hydrogen

An atom with excess energy is called the EXCITED STATE

The lowest possible enery state of an atom is called the GROUND STATE

ENERGY contained in the photon corresponds to the change in energy thatThe atom experiences in going from the excited state to the lower state.

Energy

4 excited states

Ground state

H atoms only certain types of photonsare produced.

We only see certain colors.

This means H has discreteenergy levels (with specific wavelengthsNEVER EMIT photons w/ energiesin between

11.4: The Energy Levels of Hydrogen

Energy levels of hydrogen are QUANTIZEDThat is only certain values are allowed.Scientists have found that the energy levels of all atoms are quantized.

Energy

4 excited states

Ground state

H atoms only certain types of photonsare produced.

We only see certain colors.

This means H has discreteenergy levels (with specific wavelengthsNEVER EMIT photons w/ energiesin between

11.5: The Bohr Model of the Atom

Objective: To learn about Bohr’s model of the hydrogen atom.

11.5: The Bohr Model of the Atom

Neils Bohr: at the age of 25 earned his PhD in Physics.

Constructed a model of the hydrogen atom with quantized energy levels that agreed with the hydrogen emission results.

Current theory of atomic structureIs not the same as the Bohr model.Electrons do not move around theNucleus in circular orbits.

WE DO NOT KNOW EXACTLYHow the electrons move in an atom.

11.6: The Wave Mechanical Model of the Atom

Objective: To understand how the electron’s position is represented in the wave mechanical model.

11.6: The Wave Mechanical Model of the Atom

By the mid 1920’s, Bohr’s model was shownTo be incorrect.

Victor deBroglie (France)Erwin Schrodinger (Austria)

Suggested that because light seems to have both wave and particle characteristics.

Electrons also exhibit both of these characteristics.

11.6: The Wave Mechanical Model of the Atom

Wave Mechanical Model of the atom

-electron states are described by orbitals.

Orbitals are nothing like orbits.

H atoms can be pictured as a probability of

An electron being in a certain position.

It does not tell you when the electron occupies this space or how it moves.

11.7: The Hydrogen Orbitals

Objective: To learn about the shapes of orbitals designated by s, p and d.

11.7: The Hydrogen Orbitals

An orbital is a probability map for the location of an electron.

Chemists arbitrarily define its shape as 90% of the total electron probability

HYDROGEN 1s orbitalLowest energy state (the ground state)

11.7: The Hydrogen Orbitals

Principal Energy Levels

Sublevels

Principal energy level

Shape

Figure 11.25: The three 2p orbitals.

Figure 11.27: Relative sizes of the spherical 1s, 2s, and 3s

orbitals of hydrogen.

Figure 11.28: The shapes and labels of the five 3d orbitals.

Summary

1) The number tells the principal energy level.

2) The letter tells the shape. The letter s means a spherical orbital; the letter p means a two-lobed orbital. The x, y, or z subscript on a p orbital label tells along which of the coordinate axes the two lobes lie.

11.8: The Wave Mechanical Model: Further Development

Objective: To review the energy levels and orbitals of the wave mechanical model of the atom.

To learn about electron spin.

11.8: The Wave Mechanical Model: Further Development

The last property of electrons is spin.Pauli’s exclusion principle says that no

orbital can contain more than 2 electrons and those electrons must have opposite spins.

Summary of Wave Mechanical ModelElectrons are organized in 4 levelsThese are called the quantum numbersn, l , ml, ms

n: Principal quantum number(principal energy level)l: azimuthal quantum number (shape) is the sublevel

l=0 for s sublevel, l=1 for p sublevel, l=2 for d sublevel, and l=3 for f sublevel ml=magnetic quantum numberWhen l=0, ml=0 l=1, ml= -1, 0,+1 l=2, ml=-2, -1, 0, 1, 2 l=3, ml=-3, -2, -1, 0, 1, 2, 3ms=spin quantum number +1/2 , -1/2

Add n+l = the higher the number the higher the energy level

THREE RULES

• Aufbau Principle: “building up” fill lower energy levels first

• Hund’s Rule: every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

• Pauli exclusion principle: In each energy level, electrons go in as opposite spins

11.9: Electron Arrangements in the First 18 Atoms on the Periodic TableObjective: To understand how the principal

energy levels fill with electrons in atoms beyond hydrogen.

To learn about valence electrons and core electrons.

11.9: Electron Arrangements in the First 18 Atoms on the Periodic TableElectrons fill the orbitals in the following order:1s2s2p3s3p

Electron arrangement or Electron configuration 1s1 orbital diagram(box diagram)

1s

11.9: Electron Arrangements in the First 18 Atoms on the Periodic Table

Valence electrons: electrons in the outermost(highest) principal energy level of an atom.These electrons are the most important because they are involved in bonding.Inner electrons are called the core electrons.

Atoms of elements in the same group(vertical)have the same number of electrons in a giventype of orbital (sublevel) except that the orbitalsare in different principal energy levels.

11.9: Electron Arrangements in the First 18 Atoms on the Periodic Table

So another way of writing them are:

K 1s22s22p63s23p64s1, or [Ar]4s1.

Ca 1s22s22p63s23p64s2, or [Ar]4s2

11.9: Electron Arrangements in the First 18 Atoms on the Periodic TableAtoms of elements in the same group(vertical)have the same number of electrons in a giventype of orbital (sublevel) except that the orbitalsare in different principal energy levels.

Elements with the same valence electronarrangement show very similar chemicalbehavior.

11.10: Electron Configurations and the Periodic Table

Objective: To learn about the electron configuration of atoms with Z greater than 18.

11.10: Electron Configurations and the Periodic Table

1) In a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals

2) The lanthanides, fill the seven 4f orbitals.3) The actinide series fill the seven 5f orbitals.Except for helium, the group numbers indicate the

sum of electrons in the ns and np orbitals in the highest principal energy level that contains electrons.

The same type of elements occurperiodically so that groups ofelements show similar chemistry

indicate the total # of valence electronsFor the atoms in this group.Group 5 ns2 np3

Main group elementsOr representative elements

11.10: Electron Configurations and the Periodic Table

1) The group labels for Groups 1,2,3,4,5,6,7+8 indicate the total number of valence electrons for the atoms in these groups.

2) The elements in Groups 1,2,3,4,5,6,7,8 are often called the main-group elements, or representative elements.

11.11 Atomic Properties and the Periodic Table

Objective: To understand the general properties in the periodic table.

11.11 Atomic Properties and the Periodic Table

Objective: To understand the general properties in the periodic table.

METALS: lustrous appearance, the ability to change shape without breakingand excellent conductivity of heat and electricity

NONMETALS: TEND to GAIN ELECTRONS to formNEGATIVE IONS.

METALS:TEND to LOSEELECTRONS toForm POSITIVE IONS

H

Li

Na

K

Rb

Cs

Cs > Rb > K > Na > LiLoses Leastan likelyelectron to losemost aneasily electron

Most chemically active metals

Most chemically active nonmetals

Ionization Energies Increase

IONIZATIONENERGY

DECREASESIonization Energy of an atom is the energy requiredto remove an electron from an individual atom in thegas phase.

Metals haveLow ionizationenergies