quantum theory...quantum numbers electron distribution electron configurations many important...

Quantum Theory Structure of an atom

What we are going to learn this chapter

Properties of waves and electromagnetism

Bohr’s theory of a hydrogen atom and its relation to emission lines

Debroglie wavelengths

Quantum numbers

Electron distribution

Electron configurations

Many important theories and principles

Plank’s quantum theory

Einstein’s photoelectric effect

Heisenberg uncertainty principle

Schrodinger wave equation

The Birth of Quantum Mechanics Before 1900 atoms and molecules were regarded as rebounding balls.

This is the basis of theories such as the ideal gas law and works well to describe macroscopic phenomena

After 1900 Max Planck (nobel prize 1918) discovered that atoms or molecules only emit energy in discrete levels or “quanta”

Energy was NOT continuous and this turned physics upside down.

We have no right to assume that any physical laws exist, or if they have existed up to now, that will continue to exist in a similar manner in the future. -Max Planck

Development of Atomic Models

Thomson: The plum pudding model postively charged jellylike blob, with suspended electrons

Rutherford: nuclear model all positive charge in the nucleus negative charge surrounding it:

Bohr model: expansion of Rutherford’s ideas.

Electromagnetic radiation Electromagnetic radiation is a form of energy. That has wavelike properties.

Heat

xrays

sunlight

microwaves

What Else?

PresenterPresentation NotesWhat other examples are there that you can think of: Radio waves (AM, FM TV), gamma rays, cell phone signal, garage door openers, remote controls ect…..

Properties of Waves Examples of Transverse Waves

Examples of Longitudinal Waves

PresenterPresentation Noteshave class do the wave :-D

use slinky for longitudinal wave

Transverse waves- Frequency

Frequency (ν) (pronounced nu) is equal to the _______________________________________

In equation form ν=c/λ

Units of ν is _____________(pronounced ______)

Units of λ is distance, visible light is around nm (nanometers)

Electromagnetic Waves Electromagnetic waves have and electric component and a magnetic component

Two components have same wavelength and frequency, in perpendicular planes

Travel at 2.9989x108m/s (the speed of light c)

Be sure you can identify: Amplitude Wavelength (λ) Frequency

http://hyperphysics.phy-astr.gsu.edu/

Electromagnetic Spectrum

As wavelength increases what happens to frequency? does it: A) Increase B) Decrease

Decrease

Answer the Following: Which has greater frequency red light or blue light?

Which has greater energy?

Does blue light have greater speed than red light?

Blue

Blue

NO! all have same speed

PresenterPresentation Notesblue has higher frequency

Blue has higher energy

no all are the speed of light (c)

Example Exercise The wavelength of green light from a traffic signal is centered at

522 nm. What is the frequency of this radiation.

Known: c= 3.00x108m/s λ=522 nm ν=c/λ

First convert λ to meters

Then fill into equation Fix sig figs and check for units

Another Example The frequency of radiation from your microwave is 120 GHz. Find the wavelength of the microwaves.

Similar at home exercise: The wavelength of radiation when you get an xray is 10.0 nm. Find the wavelength of the xrays.

The frequency of red light from a traffic signal is centered at 4.11x1014Hz. What is the wavelength of this radiation?

In Class Quiz

A) 1.37x106 m

B) 1.37x106 nm

C) 730 nm

D) 7.30x10-7nm

Planck’s Quantum Theory When solids are heated they emit __________

Attempts to explain this up until Plank failed at either very high or low wavelengths.

Planck discovered molecules can only emit light at specialized energies or “_________”

A “__________” is the smallest quantity of energy and is equal to:

h=6.626x10-34 J*s

PresenterPresentation NotesThis radiation is related to the temperature

Don’t let this concept of quantization seem weird. Think about everything else in life that is “quantized”. You can think of our money system as being quantized. Is there anything smaller than a penny? What else is quantized?

Number of kittens, eggs laid

Example Exercise What is the energy of green traffic light in the previous

example. (λ=522nm)

Known: c= 3.00x10-8m/s λ=522 nm=5.22x10-7m

Another Example The frequency of radiation from your microwave is 120 GHz. Find the energy of the microwaves.

The frequency of radiation of an xray is 3x1016 Hz. Find the energy of the xrays.

The wavelength of red light from a traffic signal is centered at 730 nm. What is the energy of this radiation.

To do at home:

Black Body Radiation

Blackbody radiation depends only on temperature of object

2700-3300K

Given this information about a light bulb, why are incandescent bulbs inefficient light sources?

PresenterPresentation NotesGiven this information about a light bulb, why are incandescent bulbs inefficient light sources?

Example Exercise What is the wavelength of the brightest part of the light from our next closest star, Proxima Centauri? Proxima Centauri is a red dwarf star about 4.2 light years away from us with an average surface temperature of 3,042 Kelvin.

Light: A wave or a particle.

~300B.C Aristotle

~400B.C Democritus

~1000 A.D. Alhazen

~1600 A.D. R. Descartes

~1700 A.D. Sir Isaac Newton

~1700 A.D. Robert Hooke

~1700 A.D. C. Huygens

Albert Einstein ~1900

Wave

Particle

Both

~1800 A.D. Thomas Young

Photoelectric Effect Ejection of _________ from the surface of a metal or other material when shines on it.

Energy of _______________must be higher than “_______” energy or no photon will be ejected. This ________________is called the “________________”

Ex: Violet light can eject electrons from potassium, but red light can’t.

Use E=hν to determine threshold frequency.

Shows _______________________

PresenterPresentation NotesExample: Violet light can eject an electron from potassium metal, but no amount of red light can.

Use E=hν to determine this threshold energy.

Shows wave particle duality of light.

http://phet.colorado.edu/en/simulation/photoelectric

http://phet.colorado.edu/en/simulation/photoelectric

Photoelectric Effect

If the photon is exactly the energy of the work function the electrons will be knocked loose

If the energy of the photon is higher the electrons will acquire _________________________

_________________________

What do you think happens if the intensity of light is increased but the frequency of the electrons stay the same?

Photoelectric effect example problem The work function for calcium is 4.34x10-19 J. What is the minimum frequency of light for the photoelectric effect for calcium. Then calculate the kinetic energy of the ejected electron if the light of frequency 1.00x1015 s-1 is used for irradiating the metal.

Photoelectric Effect: A second look at the equation

Before threshold frequency ___________________is present

After threshold frequency the relationship is _____________

Ek=hν - Φ y= mx + b

Photo Electric Effect Extra Notes: Rephrasing of what I said just for you

The idea here is that given one photon you are transferring that energy into the electrons in the metal. If you put exactly the amount of energy in it needed to eject an electron (aka the work function, aka Φ or W) then there is no energy left over for it to move (aka have kinetic energy)

However if the photon has more energy than the work function, it transfers that in the form of movement (aka kinetic energy)

You Try Similar at home Exercise:

You Try: The work function for magnesium is 5.9x10-19J. Calculate the minimum frequency of light required to eject electrons from magnesium. Then calculate the kinetic energy of the ejected electron if the light of frequency 1.00x1015 s-1 is used for irradiating the metal

Wave Interference

Waves with same phase interfere ________________

They ______

Waves with different phase interfere ________________

They _________

Constructive

Destructive

Double Split Experiment: Young

Wave or Particle? Summary: Experiments show it has properties of both.

Photo electric effect: Particle like

Double split experiment wavelike

This is true of very small matter

Largest so far to have measurable wave-like properties is those with mass 1610 amu

DeBroglie Wavelength: Describing matter as a wave.

De Broglie expanded wave particle duality from light to matter.

Here is equal to the velocity of the particle, not the frequency.

Often you’ll find the velocity using the kinetic energy.

Example Problems: DeBroglie Wavelength Find the DeBroglie wavelength of an electron with a K.E. of 2.56x10-27J.

Why can’t we see the waves of macroscopic matter? (hint :A baseball weights 0.145kg and it moves around 27 m/s.)

Heisenberg Uncertainty Principle

It is impossible to know simultaneously both the momentum (p) and the position (x) of a particle with certainty.

The uncertainty values are related by: (momentum=mass*velocity)

If your uncertainty of momentum decreases your measurements of position will _______________ and vice versa.

What happens to the uncertainty in position if your uncertainty in velocity decreases?

PresenterPresentation NotesThis means that even if measurements were done perfectly, the best you possible measurement will give the uncertainties in position and momentum will be h/4π

Heisenberg went for a drive and got stopped by a traffic cop. The cop asked “Do you know how fast you were going?{ Heisenberg replied, “No, but I know exactly where I am”

Why are quantum physicists so poor at sex: Because when they find the position they can’t find the momentum and when they have the momentum they can’t find the position.

Heisenburg Uncertainty Principle Examples:

You are pulled over by a police officer for speeding, and try to escape a ticket by claiming if he knew where you are, he can’t possibly know how fast you were going. The officer says, “well, I knew where you were within 0.5 m, your car probably weighs around 1300kg, so I know your velocity within…….”

Wavefunctions, Energy levels, and particles in a box

Ψ Is the “______________” Describes the __________________.

Its just a symbol for a mathematical function

Different particles have different wavefunctions

Ψ2 is the “_______________” The probability of finding the ________________________

Technically “__________________” probability of finding a particle divided by the volume.

Some questions for you.

Ψ is a function, so at any point can we absolutely determine the sign. Is it-

Always Positive Always Negative Can be both

Ψ2 Is the square of that function so can we absolutely determine the sign. Is it-

Always Positive Always Negative Can be both

Wavefunctions, Energy levels, and particles in a box

Nodes A node occurs when Ψ is _____

When this happens, what is Ψ2 equal to?

__________

So what is the probability that the electron is at a node?

_____________

Number of nodes is equal to n-1

Um….what was that again:An example.

Lets take an example where Ψ is a number (to make it simple). We’ll say 0.44cm-3

So then Ψ2=0.20cm-3.

So the wavefunction is equal to 0.44cm-3

and the probability density is equal to 0.20cm-3 .

What is the probability of finding the particle within 3cm3?

0.2cm-3 * 3cm3=0.6 or 60%

Ψ

Ψ2

X 0cm3 3cm3

Schrödinger Equation

The “hat” on the H indicates it’s an ____________.

Can use to find _________________

We will use the results of this, not the actual equation.

Particle in a Box.

Set up: A single particle is in a box.

Only certain ______________ are allowed.

Energy is “__________”

Real life relation: String anchored at two points.

Particle in a Box: Solutions

L is the ______________.

n is the _______________ it is an integer.

Solution to Shrodinger equation:

Particle in a Box Example: Use the solutions to the Schrödinger equation for a particle in a box to estimate the value of the wavelength of an electron in a helium atom given that the approximate radius of a helium atom is 100. pm.

Bohr model of the atom Before this arose they knew there were electrons traveling around a proton filled nucleus.

Laws of physics said electrons should spiral into nucleus.

Bohr said energy levels are ________________

Only works for _____________

Energy difference is given by:

Rh=2.18x10-18J n= 1,2,3…… “_________________________”. Energy levels Z=atomic number

A Hydrogen Atom Two different ways of thinking about it.

Rydberg equation

Was discovered by analyzing line spectra

Schrodinger equation

Solving the particle in a box problem in three dimensions

Same results from either equation. Rydberg’s was determined first, then Shrodinger’s results agreed.

Also true for “hydrogen-like” atoms. 1 electron atoms.

We’ll start with Schrödinger Results from solving the Shrodinger equation for the hydrogen atom

Z is the atomic number

me is the mass of an electron

e is the charge of an electron or proton

h is Planks constant

r is atomic radius

Agrees with experimental results

Stair and Ball Analogy

Bohr Model Energy Levels

The energy of a given energy level is given as:

• So the difference in energy is

PresenterPresentation NotesNotice these energies are going to be negative. Since Rh is a constant that is positive, and Z and N are positive integers. This is because an electron in its free state is equal to zero, so by convention the energy of an electron attached to a nuclei is negative. 1 is the most negative and therefore lowest energy level.

Rydberg Equation If the energy of a given energy level is defined as En=-Rh(1/n2). We can find the difference in energy shells by Enf-Eni

Either of these are often shown in text books. Use whichever you prefer but be sure not to mix them up. One has a negative one does not.

Change= __________________

__________________________

Filling the above equations into Change= Final-initial

…and finally rearranging

Rydberg's constant is experimental value

Sign Conventions for Energy: *common error causes*

A free electron has ΔE=

Energy of levels are negative, most negative and therefore lowest energy is n=1

ΔE is positive if going from ___________ energy shell i.e. ground state to anything, 25 ect… Words such as “photon absorbed” will be used.

ΔE is negative if going from _____________ i.e. anything to ground state, 52 ect… Words such as “photon emitted” will be used.

E=hν=hc/λ This energy is the energy of __________________. This must always be ______________!

ΔE is +

ΔE is -

Example Problems Rydberg Equation A photon that is emitted when a hydrogen undergoes a transition to n=2 is of the frequency 6.17x1014 s-1. Find the initial energy level.

What is the wavelength of a photon emitted during a transition from the n=5 to n=2 state in a hydrogen atom?

Lasers Light Amplification by Stimulated Emission of Radiation.

http://www.infoplease.com/images/ESCI112LASERS003.gif

PresenterPresentation NotesExplain laser pulses

Example Problem Rydberg Equation A monochromatic beam of light with a total energy of 2.5J contains 8.56x10-4 mols of photons. What is the wavelength of the beam.

Emission Series

Emission Series: Another Look (Hydrogen)

What do we use Emission Spectra for? Emission Spectra is unique to individual elements

Spectra can be collected and matched to known emission spectra to determine the element present.

Flame tests

Hydrogen Emission Spectra

Fe Emission Spectra

Flame test

Emission/Absorption Spectra Lines in visible region shown above (Hydrogen shown)

Also some occur in other spectral regions and aren’t seen without detectors

Emission Spectra: Applications in Astronomy

•Spectra is collected •Compared to known ions •Temperatures known

PresenterPresentation Notesspectra is collected and separated, they are compared to known emission spectra. Elements making up the areas are identified. Because certain elements exist at certain temperatures the temperature of the suns, supernovae ect… can be figured out.

Upper left, xray picture. upper middle and right are ultraviolet, lower left is sun’s magnetic field, lower right is IR

Time to regress to our 5 yr old selves

Why is the grass green?

Why is the sky blue?

Why is the sun yellow?

The grass is green because…

What happens if you shine green light on the plants? What happens if you shine red light on the plants?

PresenterPresentation Notesabsorbs 665 and 465 nm which is blue and red. No absorbance at green. Reflectance that’s what we see. We see it green.

Why is the Sky Blue?

PresenterPresentation NotesLong wavelength Red and orange light travels relatively unperturbed through the atmosphere. Shorter wavelengths get scattered turning the sky blue. As it travels farther much of the blue light interferes with each other cancelling each other out and making the sky paler.

During sunset the light has to travel so far through the atmosphere that most of the blue light is completely scattered and interfered away. Leaving only the reds/oranges and yellows to reach your eyes.

Previous Slide Lecture Summary

Long wavelength Red and orange light travels relatively unperturbed through the atmosphere. Shorter wavelengths get scattered turning the sky blue. As it travels farther much of the blue light interferes with each other cancelling each other out and making the sky paler. During sunset the light has to travel so far through the atmosphere that most of the blue light is completely scattered and interfered away. Leaving only the reds/oranges and yellows to reach your eyes.

Other Transitions

PresenterPresentation NotesI want you to recognize fluorescence and phosphorescence

Schrodinger Equation

Ψ is the “__________”- complex function of space and time. Describes particle.

Ψ ____________________

|Ψ2| is the “_____________________”

These probability densities make up “____________” which describe ______________________________________________.

Technically this is only good for one electron atoms but there are good approximations for many electron systems.

Eψ=Hψ

PresenterPresentation Notespsi squared is what gives us our atomic orbitals. Within 90% probability.

Atomic Orbitals Exact solutions to the Schrödinger equation for hydrogen like atoms

Approximate solutions to the Schrodinger equation for atoms with more than one electron.

Shows where the ______________________________.

Technically extend to infinity, normally drawn where there is a 90% probability of finding the electron

Atomic Orbitals

Wavefunctions

You don’t need to know these wavefunction equations, won’t need to use them. Just FYI so you know what they look like.

1s 2s

1p

3s

3p

3d

s

px

py

pz

dxy

dyz dzx

dx2-y2

dy2

Quantum Numbers Describes the distribution of electrons.

The combination of them all specifies the wavefunction. You can think of it as an address to an electron

Analogy: Finding Person on Campus

Building number 403

Room number 1100

Row F

Chair 1

Male or Female

Quantum numbers

n

l

ml

ms

Quantum Number- Principle Quantum Number

______________________- “n”

Can have any positive integer.

We saw this in the ____________ as the energy levels. 1,2,3….∞

Quantum Number- Angular Momentum Quantum Number

Angular Momentum Quantum Number- “l”

Distinguishes the orbitals of different shape.

Values have letters associated with them.

0=s, 1=p, 2=d, 3=f, 4=g,…….. These are associated with the atomic orbitals we just discussed.

Can be any integer from 0 to n-1.

n

allowed values of l

1

2

3

4

Question: For the 4th energy shell, what are the orbitals present?

Quantum Numbers- Magnetic Quantum Number

Magnetic Quantum number- “ml” (read as m sub l)

Distinguishes orbitals that have same n and l (same energy and shape) but having a different orientation.

Allowed values are from –l…0…l

n l ml

1

2

3

4

Table of Allowed Values for 1-4 energy levels

Quantum numbers Lets fill in the numbers for the n=3 energy level.

l=0 ml=0

l=1

ml= -1 0 1

l=2

ml= 2 -1 0 1 2

s

p

d ***The quantum numbers ml here doesn’t necessarily relate to the exact orbital its under, don’t worry about memorizing which goes where.

Wavefunctions

1s 2s

1p

3s

3p

3d

s

px

py

pz

dxy

dyz dzx

dx2-y2

dy2

You don’t need to know these wavefunctions equations, won’t need to use them. Just FYI so you know what they look like.

Quantum numbers- spin quantum number Spin quantum number- “ms” (read this as m sub s)

Distinguishes spin axis of electron. (also shown as spin up or spin down on diagrams)

Allowed values for each orbital are +1/2 or -1/2.

Question- How many electrons are allowed per individual orbital?

State whether each of the sets of quantum numbers is permissible for an electron, if not explain why.

Example 1: n=1, l=1, ml=0, ms=+1/2 Example 2: n=3, l=1, ml=-2, ms=-1/2 Example 3: n=2, l=1, ml=0, ms=+1 Example 4: n=3, l=2, ml=-2, ms=-1/2

Shielding and Penetration Orbitals with radial probability closer to the nucleus are more penetrating.

The closer, more penetrating orbitals are shielding the further orbitals from the nucleus.

Energies of Orbitals

Ect….to infinity

Ect….to infinity

E

Hydrogen Everything other than Hydrogen

Questions: What is the difference between the two? Why do you think hydrogen is different? Here I have draw hydrogen with atomic orbitals higher than 1. Is this correct? Why or why not?

Energies in Relation to Periodic Table

Electron Configuration Pauli Exclusion Principle states that no two electrons can have all four quantum numbers be the same.

Fill in electrons in order of energy levels.

Each orbital holds 2 electrons.

Hund’s rule- Fill across degenerate energy levels before filling orbital.

Write electron configuration as follows: 1s22s22p2

Right

Diagram for Carbon

Example: Write electron configuration for Cobalt

Diamagnetism vs Paramagnetism

_____________ has paired electrons.

Slightly Repelled by magnets

Carbon

Neon

______________ has unpaired electrons

Drawn Toward magnets

Electron configurations of ions

Simply add or subtract electrons based on the charge of the ion

Example: Ca to Ca2+

1s22s22p63s23p64s2 or ___________

Ca



Example:Ca to Ca2+

_________________

1s22s22p63s23p6 or ________________

Ca2+



Example:Cl to Cl1-

_________________ Cl

1s22s22p63s23p5 or [Ne] 3s23p5



Example:Cl to Cl1-

_________________ Cl-

Electron Configurations Exceptions

Half filled and full filled orbitals are more stable. To get d orbitals to this stability promote electrons from s orbitals

When valence d orbitals have electrons and you need to make an positive ion, remove from valence s orbitals first.

Fs also have a bunch of exceptions, don’t worry about these. If I ask for an electron configuration from the F block, just follow the rules

Question: Why can we be so fluid with our exchange of electrons from the s and d orbitals?



Example: Co to Co2+

Take away two electrons

Where are we taking them from? 1s22s22p63s23p64s23d7 or ____________

Co


add or subtract electrons based on the charge of the ion

Example: Co to Co2+

Take away two electrons _________________

1s22s22p63s23p63d7 or [Ar] 3d7

Co2+

Electron configurations of ions Example:Cr to Cr1+

Electron configuration of neutral chromium requires promotion of electron

To give stable half shell

Cr


Electron configuration of neutral chromium requires promotion of electron to give stable half shell

Then take away electron from 4s shell

1s22s22p63s23p64s13d5 or _____________

Cr


Electron configuration of neutral chromium requires promotion of electron to give stable half shell


1s22s22p63s23p63d5 or _____________

Cr1+

You Try

Use what we just learned about Chromium to write the electron configuration of Cu and Cu1+.

Electron configurations of ions Example:Cu to Cu1+

Electron configuration of neutral copper requires promotion of electron to give stable full shell


Cu

Electron configurations of ions Example:Cu to Cu1+

Electron configuration of neutral copper requires promotion of electron to give stable half shell


Cu1+

Discussion Question Using the logic we just worked out why do Cr, Mo, W, Cu, Ag and Au all form +1 ions?

Switching Gears….. Periodic Trends

Atomic Radius

Ionic Radius

Ionization Energy

Electron Affinity

What we are going to learn!

Some history about the periodic table development

Trends: using the periodic table to predict Ionization energy, electron affinity, atomic radius, effective nuclear charge and electro negativity.

Instead of 1.21, directly, we’ll just talk about some interesting elements.

PresenterPresentation Noteselectronegativity is technically chapter 9, but I think it makes a lot of sense to learn about it here!

Development of the Periodic Table

PresenterPresentation NotesOriginally scientists just wrote down tables of elements and their masses. John Newlands noticed that every 8 elements the elements had similar properties….up through calcium. Then Mendeleev and Meyer designed a different periodic table closer to our own. This was able to predict the discovery of new elements with similar properties to those already known.

There were still problems with it though since it was arranged by mass, for instance Ar and K would have been switched which makes no sense. Finally Mosley figured out that it was related to atomic number.

Effective Nuclear Charge

Ele

ctro

nega

tivity

Electronegativity

Effective Nuclear Charge

Increases __________________

Vertical Trend doesn’t make a lot of sense to discuss here so we will ignore it

Nuclear charge felt by an electron when both the actual _____________________and the _______________ of other electrons are taken into account.

Core electrons shield the ______________.

Reason behind many of the other trends!

Atomic Radius

Increases down a group

Increasing amount of ___________ increases size.

Increases right to left across a period

Due to ________________________________, holds electrons closer

Examples Rank the following in order of increasing radius

Li, C, F

Li, K, Rb

Ba, Se, F

Would you be asked to rank, Be, Al, Ge? Or similar?

Ionic Radius Cations are smaller

More positive is smaller. _____ is smaller than ______

Anions are bigger More negative is bigger. _____is larger than ______

Higher effective nuclear charge when electrons are removed, lower when electrons are added.

Example: Draw an arrow from the smallest to largest species in the following isolectronic series

S2-, Cl-, Ar, K+, Ca2+

First: what is this “isoelectronic” word iso=_________ electronic: sounds sort of like electron so…… isoelectronic= ___________________

More examples Rank the following in order from smallest to largest:

Li+, Be2+, F-

Cu+, Cu2+, K+

Cl-, F-, Br-

and another:

Given the following pictures of ionic compounds match the pictures with the ions: Na+, Mg2+, Cl-, O2-

Purple

Pink Blue Green

First Ionization Energy Minimum energy required to remove an electron from its ground state

General Trends with exceptions

Exceptions come from getting half and fully filled subshells: You need to know these!

Be

B

C

N

O

F

Mg

Al

Cl

P

S

What is the electron configuration of Be, B, N, O? Use this to explain the higher ionization energy.

Be:

B:

N:

O:

Examples:

First some normal ones

He, Ne, Ar

B, Li, Ne

Rank the following in order from lowest to highest first ionization energy.

Now some exceptions:

Li, Be, B

C, N, O

Second and Third Ionization Energy

First is always the _________.

Nuclear charge** doesn’t change, so each electron feels more of a positive pull.

**Notice this says nuclear charge, not effective nuclear charge

1st

2nd

3rd

PresenterPresentation Notesget students to come up and do demo.

Electron Affinity e-

eg. Cl + e- Cl-

Electron Affinity General trends, but has a fair amount of exceptions

PresenterPresentation NotesNotice all the exceptions here: Point them out.

Also notice: it’s the same trend as effective nuclear charge. This is because the higher the effective nuclear charge is, the more likely that it can accept another electron.

Lets look at some values

Positive electron affinities means its exothermic

PresenterPresentation NotesMany of the noble gasses haven’t actually had their electron affinities measured and are assumed to be zero, it may be negative though.

Electron Affinity

Why do the noble gasses and not form stable atomic anions?

Ne: 1s22s22p6 Already has a filled shell. Already stable, adding an electron isn’t energetically favor.

N: 1s22s22p3 Have a filled half shell, adding another electron causes ______________________________.

Why do group 2A (alkaline earth) an 7A (halogens) form stable atomic anions? (aka -1 ions)

F: High effective nuclear charge means ______________. One more electron fills the shell, this adds __________

Na: Filled sub-shells are stable.

F-:

Na-:

Electronegativity (Chapter 2.12) Ability of an atom to attract electrons _______________________________________________

Not as many important exceptions, mostly involving the d block. We won’t worry about them.

Why aren’t there as many exceptions with these?

Electronegativity and Electron affinity

Same general trend

Both involve ability of an atom to _______________

Electronegativity is that ability while _________, attracts shared ______________

Electron affinity is that ability of an _______________.

e-

eg. Cl + e- Cl-

CO2

O atoms are electronegative

high electron affinity

Inert Pair Effect

Use electron configurations to explain.

Sb Sb: Sb3+:

Sb5+:

Pb Pb:

Pb2+:

Pb4+:

Do the rest at home for practice!

Tendency of heavier atoms to form ions with a difference in charge of two.

Diagonal Relationship Diagonal bands going down and right have similar properties.

Look at radii and ionization energies and explain.

General Trends in Chemical Properties

Understand why the groups in general form certain ions.

Understand why trends occur including reactivity in section 1.21

What follows is a smatterering of interesting aspects of elements I think is pertinent to real life.

Silicone as a basis for life?

Smallest element with same valence electron structure as carbon

Many similar properties to carbon

Probably not as likely as SciFi makes it out to be

Doesn’t bind with as many atoms Doesn’t make double or triple bonds, severely limiting chemistry

Si chains with H are unstable in water.

Si chains with O are more stable, but still not as stable as carbon.

Noble Gases He: refrigerant for super conducting magnets like in MRIs: used in scuba diving and blimbs.

Many uses where you need an “inert” atmosphere, chemistry, lightbulbs, storage ect…..

“Neon” lights, which are really many elements.

Interesting tidbit: Argon means “the lazy one” in greek.

He Ne Ar Kr Xe

PresenterPresentation NotesHelium walks into a bar, The bar tender says "We don't serve noble gasses in here." Helium doesn't react.

Quantum TheoryWhat we are going to learn this chapterThe Birth of Quantum MechanicsDevelopment of Atomic ModelsElectromagnetic radiationProperties of WavesTransverse waves- FrequencyElectromagnetic WavesElectromagnetic SpectrumExample ExerciseAnother ExampleIn Class QuizPlanck’s Quantum TheoryExample ExerciseAnother ExampleBlack Body RadiationExample ExerciseLight: A wave or a particle. Photoelectric EffectSlide Number 20Photoelectric EffectPhotoelectric effect example problemPhotoelectric Effect: A second look at the equationPhoto Electric Effect Extra Notes: Rephrasing of what I said just for youYou Try Similar at home Exercise:Wave InterferenceDouble Split Experiment: YoungWave or Particle? DeBroglie Wavelength: Describing matter as a wave. Example Problems: DeBroglie WavelengthHeisenberg Uncertainty PrincipleHeisenburg Uncertainty Principle Examples:Wavefunctions, Energy levels, and particles in a boxWavefunctions, Energy levels, and particles in a boxNodesUm….what was that again:An example. Schrödinger EquationParticle in a Box. Particle in a Box: SolutionsParticle in a Box Example:Bohr model of the atomA Hydrogen AtomWe’ll start with SchrödingerStair and Ball AnalogyBohr Model Energy LevelsRydberg EquationSign Conventions for Energy: �*common error causes*Example Problems Rydberg EquationLasersExample Problem Rydberg EquationEmission SeriesEmission Series: Another Look (Hydrogen)What do we use Emission Spectra for?Emission/Absorption SpectraEmission Spectra: Applications in AstronomyTime to regress to our 5 yr old selvesThe grass is green because…Why is the Sky Blue?Previous Slide Lecture SummaryOther TransitionsSchrodinger EquationAtomic OrbitalsAtomic OrbitalsWavefunctionsQuantum NumbersAnalogy: Finding Person on Campus Quantum numbersQuantum Number- Principle Quantum NumberQuantum Number- Angular Momentum Quantum NumberQuantum Numbers- Magnetic Quantum NumberQuantum numbersWavefunctionsQuantum numbers- spin quantum numberState whether each of the sets of quantum numbers is permissible for an electron, if not explain why.Shielding and PenetrationEnergies of OrbitalsEnergies in Relation to Periodic TableElectron ConfigurationExample:Diamagnetism vs ParamagnetismElectron configurations of ionsElectron configurations of ionsElectron configurations of ionsElectron configurations of ionsElectron Configurations ExceptionsSlide Number 88Electron configurations of ionsElectron configurations of ionsElectron configurations of ionsElectron configurations of ionsElectron configurations of ionsYou TryElectron configurations of ionsElectron configurations of ionsDiscussion QuestionSwitching Gears….. Periodic TrendsWhat we are going to learn!Development of the Periodic TableSlide Number 102Effective Nuclear ChargeAtomic RadiusExamplesIonic RadiusExample:More examplesand another:First Ionization EnergySlide Number 111Examples: Second and Third Ionization EnergyElectron AffinityElectron AffinityLets look at some valuesElectron AffinityElectronegativity (Chapter 2.12)Slide Number 119Electronegativity and Electron affinityInert Pair EffectDiagonal RelationshipGeneral Trends in Chemical Properties Silicone as a basis for life?Noble Gases

quantum theory...quantum numbers electron distribution electron configurations many important...

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