Chapter 10- Part One Modern Atomic Theory Objectives: Review history… (10.1) Describe electromagnetic radiation (10.2) Describe the Bohr atom (10.3) Explain.

Chapter 10- Part OneModern Atomic Theory

Objectives:

•Review history… (10.1)•Describe electromagnetic radiation (10.2)•Describe the Bohr atom (10.3)•Explain energy levels of electrons and diagram atomic structures for elements (10.4 & 10.5)

Review…

• Dalton• Thomson• Rutherford

– Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

Electromagnetic Radiation

• Light travels in waves, similar to waves caused by a moving boat or a pebble tossed in a pond

• Light is a form of Electromagnetic Radiation– Form of energy that exhibits wavelike

behavior as it travels through space

Electromagnetic Radiation

• All waves can be described in 3 ways:– Amplitude – the height of the wave,

results in the brightness or intensity of the light

– Wavelength distance between consecutive peaks in a wave

– Frequency number of waves that pass a given point in a second

Electromagnetic Radiation

• Speed of light in air Electromagnetic radiation moves through a vacuum at speed of 3.00 x 108 m/s

• Since light moves at constant speed there is a relationship between wavelength and frequency:

c = Wavelength and frequency are inversely

proportional

Electromagnetic Spectrum

Quantum TheoryQuantum Theory

• Wave theory does not explain– Heated iron gives off heat

•1st red glow yellow glow white glow

– How elements such as barium and strontium give rise to green and red colors when heated

Quantum TheoryQuantum Theory

• Max Planck (1858-1947)– Proposed that there is a fundamental

restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy a quantum.

•Energy is released in Quanta

Quantum TheoryQuantum Theory

• A quantum is a finite quantity of energy that can be gained or lost by an atom

E = hE = energy

v = frequencyh = 6.626 x 10-34 J/s

• This constant, h, is the same for all electromagnetic radiation

Photoelectric EffectPhotoelectric Effect

• The emission of electrons by certain metals when light shines on them– Albert Einstein (1905) used Planck’s

equation to explain this phenomenon;• proposed that light consists of quanta of

energy that behave like tiny particles of light

• Photon = individual quantum light (also known as a particle of radiation)

Photoelectric EffectPhotoelectric Effect

• He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore the energy of each photon is too low to dislodge an electron.

• Analogy:– 70 cents placed in soda machine: no soda– 30 cents more and you will get your soda

Now…

• Light can be described as both particles and waves

• Dual Wave-Particle Nature of Light was accepted

• What does this mean for the atom???

LineLineSpectrumSpectrum• Elements

in gaseous states give

off colored light

– High temperature or high voltage– Always the same– Each element is unique

• http://home.achilles.net/~jtalbot/data/elements/

Line SpectrumLine Spectrum• Ground state

– Lowest energy level available

• Excited state– State in which electron has a higher

potential energy than in its ground state

– Farther from nucleus– Higher potential energy

Line SpectrumLine Spectrum• Electron falls from higher energy

level to lower one…emits light at a specific frequency

• Color of light emitted depends on difference between excited state and ground state

– See figure 10.5 page 201

Line SpectrumLine Spectrum• Each band of color is produced by

light of a different wavelength• Each particular wavelength has a

definite frequency and has definite energy

• Each line must therefore be produced by emission of photons with certain energies

Line SpectrumLine Spectrum

Line SpectrumLine Spectrum• Whenever an excited electron

drops from such a specific excited state to its ground state (or lower excited state) it emits a photon

• The energy of this photon is equal to the difference in energy between the initial state and the final state.

Niels Henrik David Niels Henrik David BohrBohr

• 1885-1962• Physicist• Worked with

Rutherford– 1912

• Studying line spectra– of hydrogen

Niels Henrik David Niels Henrik David BohrBohr

• 1913 – proposed new atomic structure– Electrons exist in

specific regions away from the nucleus

– Electrons revolve around nucleus like planets around the sun

The Bohr AtomThe Bohr Atom• Nucleus with protons and neutrons• Electrons move in “stationary states”

which are stable (paths or orbits)• When an electron moves from one state

to another the energy lost or gained is done is ONLY very specific amounts

• Each line in a spectrum is produced when an electron moves from one stationary state to another

The Bohr AtomThe Bohr Atom

•Model didn’t seem to work with atoms with more than one electron

•Did not explain chemical behavior of the atoms

Wave Matters…Wave Matters…

•Louis de Broglie (1924)

•Proposed that electrons might have a wave-particle nature

•Used observations of normal wave activity

Wave Matters…Wave Matters…

•Erwin Schrodinger (1926)

•Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves

•Quantum Theory

Quantum TheoryQuantum Theory

• Describes mathematically the wave properties of electrons and other very small particles

• Applies to all elements (not just H)

Energy Levels of ElectronsEnergy Levels of Electrons

• Principal energy levels– Designated by letter n– Each level divided into sublevels

• 1st energy level has 1 sublevel• 2nd energy level has 2 sublevels• Etc.

Energy Levels of ElectronsEnergy Levels of Electrons

OrbitalsOrbitals• Electrons don’t actually

orbit like planets• Orbital:

region in space where there is a high probability of finding a given electron– Each orbital sublevel can hold 2

electrons

OrbitalsOrbitalsEach sublevel (orbital) has a specific shape

http://daugerresearch.com/orbitals/

OrbitalsOrbitals•Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons which must have opposite spins

•Electrons can only spin in two directions•Shown with arrows

Rules for Orbital FillingRules for Orbital Filling

• Pauli’s Exclusion Rule– No two electrons have the same set of

quantum numbers• Hund’s Rule

– Electrons will remain unpaired in a given orbital until all orbitals of the same sublevel have at least one electron

1s 2s 2p 3s 3p

Rules for Orbital FillingRules for Orbital Filling

• Diagonal Rule– The order of filling

once the d & f sublevels are being filled

– Due to energy levels

Rules for Orbital FillingRules for Orbital Filling

Quantum NumbersQuantum Numbers

• Numbers that specify the properties of atomic orbitals and their electrons

• Principle Quantum Numbers:– Symbolized by n, indicates the main

energy levels surrounding a nucleus, which indicates the distance from the nucleus (shells or levels)

Quantum NumbersQuantum Numbers• Orbital Quantum Number:

– Indicates the shape of an orbital– (subshell or sublevels)– s, p, d, f Principal Quantum # Orbital Quantum #

1 1s2 2s, 2p3 3s, 3p, 3d4 4s, 4p, 4d, 4f

Quantum NumbersQuantum Numbers

• Magnetic Quantum Number:– Indicates the orientation of an orbital

about the nucleus– Orbital position with respect to the 3-

dimensional x, y, and z axes

Quantum NumbersQuantum Numbers

• Spin Quantum Number:– Indicates two possible states of an electron

in an orbitalType of Orbital Number of Orbitals

s 1 ( )p 3 (x, y, z) ( , , ,)d 5 ( , , , , )f 7

Each orbital holds a maximum of 2 electrons

Application of Quantum Application of Quantum NumbersNumbers

• Several ways of writing the address or location of an electron

• Lowest energy levels are filled first• Electron Configuration: using the

diagonal rule, the principal quantum number (n), and the sublevel write out the location of all electrons

12C: 32S:

1s22s22p63s23p4

1s22s22p2

Application of Quantum Application of Quantum NumbersNumbers

• Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons

• See examples on whiteboard

Chapter 10 – Part TwoThe Periodic Table

Objectives:•Understand the arrangement of the Periodic Table (10.6)•Identify connections between electron configuration and placement on the periodic table

The Periodic TableThe Periodic Table• 1869 – arrangement proposed by

Dmitri Mendeleev– And Lothar Meyer (different layout)– Still similar today– Based on increasing atomic masses

and other characteristics– Was able to predict properties of

elements not yet discovered….and was correct!

The Periodic TableThe Periodic Table• Horizontal rows

– Periods– Corresponds to outermost energy

level

• Vertical Columns– Groups or families– Similar properties; reactions

The Periodic TableThe Periodic Table• Several systems for naming groups

– Left to right, 1-18– Roman numerals and A and B

• Used in this book• Group A: Representative Elements

– Noble Gases– IA – Alkali Metals– IIA – Alkaline Earth Metals– VIIA - Halogens

• Group B: Transition Elements

The Periodic TableThe Periodic Table• Chemical behavior and properties of

elements in a particular family similar– Have the same outer shell electron

configuration– Figure 10.15 page 211

• Noble gas configuration (shortcut)– Use previous noble gas in square brackets– Finish with valence electrons

The Periodic TableThe Periodic Table• Examples:

– K is 1s22s22p63s23p64s1 or [Ar]4s1

– Ca is 1s22s22p63s23p64s2 or [Ar]4s2

• Write abbreviated configuration for the following elements:– Fr– Y

The Periodic TableThe Periodic Table• Arrangement of Periodic Table also

means that elements filling similar orbitals are grouped– s block– p block– d block– f block

• Know these blocks…

The Periodic Table - The Periodic Table - HighlightsHighlights

• The number of the period corresponds to the highest energy level occupied by electrons in that period

• The group numbers for the representative elements are equal to the total number of valence electrons in that group

The Periodic Table - The Periodic Table - HighlightsHighlights

• The elements of a family have the same outermost electron configuration– (just different energy levels)

• The elements within each of the s, p, d, and f blocks are filling the corresponding orbitals

• There are some discrepancies with order of filling– (not covered in this book)