Chapter 5 – Models of the Atom

Friday, October 1

• Do Now – Current Events and Chemistry

• Objective – Chapter 4 Review 5.1 Models of the Atom

• Homework – Start Pg. 132 # 1-6 (Due Tuesday)

Supersonic Man

How does this relate to anything we have talked about already this year?

Write you idea in your notes

10/11/11

• Objective: to explore the electronic structure of the atom

• Do now: Take out HW questions (Pg 119 # 25-28, 30-33)

Review Questions – Pg 119 # 25-28, 30-33

25) Atoms of different elements contain different numbers of protons26) Mass # - Atomic # = # of neutrons27) Isotopes have different mass numbers and different numbers of neutrons28) For each isotope, multiply its atomic mass by its % abundance, and then add the products30) 194 is the mass number31) The atomic mass is a weighted average of the masses of its

isotopes32) Lithium-6, 3p+, 3 e-, 3 no; Lithium-7, 3p+, 3 e-, 4 no; Calcium-42,

20 p+, 20 e-, 22 no; Calcium-44, 20 p+, 20 e-, 24 no; Selenium-80, 34 p+, 34 e-, 46 no; Selenium-78, 34 p+, 34 e-, 44 no;

33) Beryllium, Magnesium, Strontium, Barium, Radium

Review what we know

Democritus

Dalton

Thompson

Rutherford

The Bohr Model • Rutherford model: electrons move around the nucleus

•Bohr: How do they move around the nucleus?

Bohr model

• Bohr knew that different elements could emit different colors of light when heated– Each element has its own specific color– Why does light come in different colors?

Iron Lithium

• 1) Rutherford’s atomic model needed to be replaced because it could not explain why elements give off characteristic colors

• 2) In the bohr model, an electron is found only in specific circular paths around the nucleus.

• 3) the QMM determines the allowed energy levels an electron can have and the likelihood of finding an electron in a given location

• 4) The sublevels have different shapes• 5) An electron can move from one energy level to another by

gaining or losing energy• 6) To move from one energy level to another requires emission

or absorption of an exact amount of energy.

Light!

•Light is composed of waves–Shorter wavelength = higher energy•Blue = short wavelength

–Long wavelength = lower energy•Red = long wavelength

Bohr model

• Bohr put this knowledge together to explain how electrons are organized in atoms– Electrons exist in fixed energy levels– They can jump from one level to another• But they are never in between two levels

Bohr Model

• Energy levels– All electrons in a given energy level have equal

energy– Electrons in higher energy levels have greater energy

Bohr Model

• Energy has to be added for an electron to level up

• Energy has to be released for an electron to level down– This energy is released as light

Bohr Model

• So why do different elements produce different colors of light? – Rydberg equation:

λ = wavelength R = 0.0110 nm−1 (Rydberg constant)m = ending energy level (1, 2, 3…)n = starting energy level ( n>m)

– Each element’s electrons are excited to higher energy levels in different patterns that produce different wavelengths of light

Emission Spectra

• Each element emits a combination of wavelengths that we interpret as a unique color

10/18/11

• Do now: How are electrons localized around the nucleus in the Bohr model?– What sort of regions do they move in?– How do they move?– How many electrons are in each region?– How does an electron’s position relate to its

energy?• Do later: Complete your orbital poster

Bohr Model v. Quantum Mechanical Model

• In the Bohr model, electrons travelled in circular orbits around the nucleus– 2 in first shell, 8 in second shell, 8 in third shell…– One shell for each row of the periodic table– All electrons in a shell have equal energy

Quantum Mechanical Model

• Our understanding of energy levels has improved since the Bohr model

• Electrons do not orbit around the nucleus in circles.– They exist in 3-D orbitals – regions of space where

you are likely to find an electron• Orbitals show the probability of where an electron is likely to be.• (this is where chemistry starts to get weird)

Where is the propeller?

The Quantum Mechanical Model

• The QMM is based on the Schrodinger equation– Predicts where electrons are likely to be in an atom

based on their energy– The solutions to this equation are orbitals

Atomic Orbitals

• Atomic Orbitals are regions of space around a nucleus where a given electron is likely (90%) to be. – 4 kinds of orbitals: s, p, d, f• Each has a different shape

– Each orbital can hold up to 2 electrons

Bohr Model v. Quantum Mechanical Model

• In the Bohr model, electrons existed in shells around the nucleus– 2 in first shell, 8 in second shell, 8 in third shell…– One shell for each row of the periodic table

Bohr Model v. Quantum Mechanical Model

• In the QMM, electrons are found in energy levels that are subdivided into orbitals– Each orbital can hold up to 2 electrons– 1st energy level has 1 orbital (2 electrons)– 2nd energy level has 4 orbitals (8 electrons)– 3rd energy level has 9 orbitals (18 electrons)

S Orbitals

• The first energy level only has 1 orbital– 1s orbital – Sphere of space where up to 2 electrons are likely

to be found

P orbitals

• The second energy level has 1 2s orbital and 3 2p orbitals– P orbitals are shaped like barbells– Each holds up to 2 electrons, 6 e-s total– 3 p orbitals oriented on the x, y, and z axes

D orbitals

• The third energy level has 1 3s orbital, 3 3p orbitals and 5 3d orbitals– 9 orbitals, 18 electrons total

F orbitals

• The fourth energy level has 1 4s orbital, 3 4p orbitals, 5 4d orbitals, and 7 4f orbitals– 16 orbitals, 32 electrons total

Review of 5.1

Principal Number ofEnergy Levels Sublevels Orbitals

1 1 1s (1 orbital)

2 2 2s (1 orbital), 2p (3 orbitals)

3 3 3s (1 orbital,) 3p (3 orbitals), 3d (5 orbitals)

4 4 4s (1 orbital,) 4p (3 orbitals), 4d (5 orbitals), 4f (7 orbitals)

Chlorine has 17 electrons 1s2 2s2 2p63s23p5

Energy of Orbitals

• Two trends– More complex higher energy• (f > d > p > s)

– Higher energy level higher energy• (p3 > p2 > p1)

Bohr Model v. Quantum Mechanical Model

• Like the Bohr Model, in the QMM:– Electrons always fill in a lower energy orbital

before a higher energy orbital– Higher energy orbitals are farther from the

nucleus

Practice time!

• Which orbital has higher energy electrons?– 2S or 1S– 2P or 2S– 3D or 2P– 3P or 4S

10/20

• Objective: To practice our knowledge of the QMM and to understand why orbitals take their specific shapes

• Do now: – Take out your orbital/energy level poster– Take 5 minutes to fill in the worksheet at the front desk

• Do later: Read section 5.3, questions 16-21

Quantum Mechanical Model Review

• Electrons are localized in 3-D regions called orbitals– 2 e- in each orbital

• 4 kinds of orbitals: s, p, d, f– 1st energy level has 1 s1 orbital

– 2nd energy level has 1 s2 & 3 p2 orbitals

– 3rd energy level has 1 s3, 3 p3, & 5 d3 orbitals

– 4th energy level has 1 s4, 3 p4, 5 d4, & 7 f4 orbitals

Quantum Mechanical Model Review

• All electrons in a given orbital have equal energy

Quantum Mechanical Model Review

• Within an energy level, more complex orbitals have higher energy electrons

Quantum Mechanical Model Review

• Electrons in the same kind of orbital will have greater energy at higher energy levels

Practice time!

• Which orbital has higher energy electrons?– 2S or 1S– 2P or 2S– 3D or 2P– 3P or 4S

Quantum Mechanical Model Review

• Higher energy orbitals are located farther out from the nucleus

Practice problem!

• If you were drawing 2p, 3p, and 4p orbitals, how would your drawings differ?

Orbital Shapes

• Why do orbitals take these funny shapes?– Honors topic– Physics detour!

10/21/11

• Objective: – To practice calculating wave measurements– To explore the wave-particle duality and its history

• Do now: Check your homework with a partner

• Do later: Interference worksheet

Homework• 16) Frequency and wavelength of light are inversely proportional

to each other. λ=c/ν, ν=c/λ• 17) Electrons in atoms absorb energy as they move to energy

levels, and then lose the energy by emitting it as light as they drop back

• 18) The light emitted in an electronic transition from higher to lower energy levels has a frequency that is directly proportional to the energy change of the electron

• 19)QM describes the motions of atoms and subatomic particles; classical mechanics describes the motion of larger bodies

• 20)Electron transitions from higher levels to n=1• 21) c, a, b

Waves

Amplitude

Wavelength (λ)

Frequency (v)

• Waves are predictable disturbances that travel through space, usually accompanied by the transfer of energy• Transverse and longitudinal

10/24/11

• Objective: – To review the properties of waves – To explore the experiments that led to quantum

mechanics• Do Now: Take 7 minutes to finish the waves

practice sheet. Use a calculator.• Do Later: Worksheet – Experimental

foundations of Quantum Mechanics

C = λ ν

• C = speed of light = 3 x 108 meters/second• λ = wavelength of light (meters)• ν = frequency (s-1, hz)

Practice problem!• Yellow-orange light has a wavelength of 600 nm. What

is its frequency?

• 72 waves crash on a beach in an hour. What is the frequency of these waves in hz?

• What is the wavelength of radiation with ν = 1.50 x 1013 Hz? Does this radiation have a longer or shorter wavelength than red light (700 nm)?

Interference

• Waves can add to each other arithmetically– Constructive interference: When 2 waves are in

sync, they add together

Interference

• Destructive interference: When 2 waves are out of sync, they can cancel each other out

Planck & Black Body Radiation

• Black body – absorbs/emits all colors of light• Equipartition theory

• Ultraviolet catastrophe

• E = hν– h = 6.63 x 10-34

J s * ν

Photoelectric Effect

10/27/11

• Objective: To explore the concept of wave-particle duality

• Do Now: Take out your homework sheet and open your binder to your notes from yesterday

• Do Later: Quantum Mechanics equations practice

E = hν

• Light is made of massless “particles” called photons– Each photon has an energy proportional to the

frequency of its light wave

Energy = planck constant * frequencyE = 6.63 x 10-34

J s * ν

E = 4.14 x 10-15 eV * ν

Practice Problem!

• The ionization energy of electrons in potassium is 2 electron volts. What is the longest wavelength of light (in nm) that can eject an electron from this potassium?– E = hν– C = λν– h = 4.14 x 10-15 eV

De Broglie Wavelength

• Planck/Einstein: Waves also behave like particles

• De Broglie: Do particles of matter behave like waves?

Derivation of the De Broglie Wavelength

• E = hν

De Broglie Wavelength

• λ = h/mv• Matter also acts like a particle and a wave– What does this mean?– Existence is wavy

De Broglie Wavelength

• What is the de Broglie wavelength of:– A person running at 1 m/s• Mass = 50 kg

– A baseball at 10 m/s• Mass = 0.15 kg

– An electron at 6 x 106 m/s • Mass = 9.1 x 10-31 kg

Wave interference

The Schrodinger Equation

• The QMM is based on the Schrodinger equation– Predicts where electrons are likely to be in an atom

based on their energy– The solutions to this equation are orbitals

Tuesday Oct. 5th

• Do Now – Check neighbor’s homework and give an effort grade of √-, √, √+

• Objective – Electron Configuration in Atoms

• HW - Pg 136 # 10-13

Pg 132 # 1-6 1. It couldn’t explain why metals give of characteristic colors when heated or explain the chemical properties of elements.2. An electron is found only in a specific path or orbital around the nucleus3. It determines allowed energies and where it is likely to be located. 4. The have different shapes. 5. By gaining or losing a quantum of energy.6. In an atom, electrons have certain fixed energy levels. To move to a different level requires the emission or absorption of an exact amount of energy or quantum. The energy of the electron is said to be quantized.

11/1/11

• Objective: To practice QM equations and finish (?) quantum explorations of matter

• Do now: Take out homework and check with a partner– When you’re done, volunteer to do a problem on the

board• Do Later: Read Ch. 5.2– pg 136 # 10, 12– Pg 150 # 50, 57, 63, 64

11/2/11

• Objective: to finish QM and understand how electrons fill the orbitals of an atom

• Do now: Take out QM equations practice, look over with a partner (5 min) and volunteer to write a question on the board!

11/3/11

• Objective: To learn and practice the principles of electron configurations

• Do Now: Check your homework against the board– Questions?

• Do Later: – Pg. 149, # 24, 29-32, 37, 41, 44, 55, 59

Homework 11/2

10) Aufbau principle, Pauli exclusion principle, Hund’s rule12) 3d, 4s, 3p, 3s, 2p50) Argon, Ruthenium, Gadolinium57) Sodium/Na, Nitrogen/N, Silicon/Si, Oxygen/O, Potassium/K, Titanium/Ti63) a, 3p to 3s64) a, 1s22s22p63s2

Electron Configuration

• How atoms are arranged to get the most stable atom possible.

Three Rules 1. Aufbau Principle 2. Pauli Exclusion Principle 3. Hund’s Rule

Aufbau Principle

Electrons will occupy the orbitals of the lower energy levels first.

Pauli Exclusion Principle

An orbital may describe 2 electrons, in other words, orbitals can hold 2 electrons If two electrons occupy the same orbital they must have opposite spins ↑↓

Hund’s Rule

If multiple orbitals have the same energy, each will fill with one electron before any gets two electrons

Orbitals with equal energyare called degenerate

orbitals

Practice Time!

• Draw in the electron configuration of silicon (atomic number = 14)

The Periodic Table

11/5/11

• Do now (5 min): – Draw an aufbau diagram in your notes – then list an atom’s orbitals from least energy to

highest (1s, 2s, 2p… 7s)– New thing! You get a small

participation grade daily for completing each do now well and on time.

Hw 11/4• 24) Rutherford: negatively charged electrons surround a dense, positively charged

nucleus. In bohr’s model, the electrons are assigned to concentric circular orbits• 29) a.1 b.2 c.3 d.4• 30) a.2 b.1 c.3 d.6• 31) Electrons occupy the lowest possible energy levels. An atomic orbital can hold

two electrons. One electron occupies each degenerate orbital before any electron pairing occurs

• 32) 1s22s2 2p63s23p3 , 1s22s2 2p63s2, 1s22s2 2p5

, 1s22s2 2p63s23p6

• 37) 2s, 3p, 4s, 3d• 41) Frequency is the number of oscillations that a wave makes in a second. • 44) Both UV light and microwaves travel at the same speed. UV has short

wavelength and high frequency, microwaves have long wavelengths and low frequency.

• 55) 4.36 x 10-5 cm, visible spectrum, 6.88 x 1014 hz• 59) It is not possible to know both the position and velocity of a particle at the

same time.

Exceptions to the Configuration Rules

What do you think would happen if an atom were to differ from the configuration rules?

There are small differences between the 3d and 4s energy levels and even smaller differences between the 5f and 6d levels

Friday, October 8th

• Do Now – The Electric Pickle

• Obj – Physics and the Quantum Mechanical Model

• HW – Pg. 146 # 16 – 19 Lab Handout Part 1

C = λ v

The speed of light is the product of the wavelength and the frequency.

Light is electromagnetic radiation

Atomic Spectra

• When atoms absorb energy, the electrons move to a higher energy level. These electrons lose energy by emitting light when they return to a lower energy level.

The frequencies emitted are unique for different elements and combinations of elements

Spectrum

Atomic Spectrum

Quantum Mechanics

Heisenberg Uncertainty Principle

It is impossible to know exactly both the velocity and position of a particle (electron) at the same time.