AP Chemistry



AP Chemistry Chapter 5 Notes

(Student’s edition)

Chapter 5 problem set: 10-13, 32, 40-42, 45, 49, 51, 56, 77, 79, 85, 94, 103, 108, 117, 118, 123, 124

5-1 Fundamental Particles

|Subatomic particle |Location |Charge |Mass |Determines… |

|Proton (p or p+) |Nucleus |+ 1 |1.0073amu | |

|Neutron(n or no) |Nucleus |0 |1.0087amu | |

|Electron (e-) |Orbit nucleus |-1 |0.00054858amu | |

2. The Discovery of Electrons – Students should read this section

In the 1870’s, English physicist William Crookes studied the behavior of gases in

vacuum tubes. Crookes tubes - forerunner of picture tubes in TVs

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Crookes’ theory was that some kind of radiation or particles were traveling from

the cathode across the tube. He named them cathode rays.

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20 years later, J.J. Thomson (English) repeated those experiments and devised

new ones. In 1897, JJ Thomas discovered the electron.

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Thomson used a variety of materials, so he figured cathode ray particles must be

fundamental to all atoms. J.J. Thomson gets credit for discovering the electron.

Plum Pudding Model

Thomson and Milliken (oil drop experiment) worked together (their data, not

them) to discover the charge and mass of the electron

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Electron charge: 1.602 x 10-19 coulomb this is the smallest charge ever detected

Electron mass: 9.11 x 10-28 grams this weight is pretty insignificant

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3. Canal Rays and Protons – Students should read this section

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4. Rutherford and the Nuclear Atom – Students should read this section

The Rutherford Gold Foil Experiment: In 1911, Rutherford (New Zealand) …

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- The Experiment: Alpha particles from Polonium (in the lead box) were released towards a thin sheet of gold foil. Most of the particles went through and were seen on the detector screen. 1 in 20,000 alpha particles bounced back.

Concluded: 1 – the positive portion of the atom is in the middle

2 – most of the atom is empty

3 – most of the mass is in the middle

4 – electrons orbit the nucleus

Shortcomings of the Rutherford Model: According to gravity, electrons should move towards the nucleus eventually - they don’t. So.... more work needs to be done to understand the structure of the atom.

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5. Atomic Number – Students should read this section

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6. Neutrons – Students should read this section…Chadwick

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7. Mass Number and Isotopes

We look to the periodic table to give us information about the number of particles are in atoms.

Atomic Number:

Nucleons: particles that make up the nucleus.

Proton and Neutrons make up most of the mass of atoms.

Protons: 1 amu, 1.67 x 10 -24 grams, positive charge, determines identity of the atom.

Atomic # (Z): Always a whole number, smaller number on the periodic table. # of protons in the nucleus, also indicates the # of electrons if the

element is not charged

Neutron: neutral, determines the mass of the atom, mass is slightly more than 1 amu

Electrons: not a nucleon, negative charge, orbit nucleus, determine charge of an atom

Mass Number:

Originally it was thought that all atoms of the same element had the same mass (Dalton)

Hydrogen was observed to have the lowest mass (assigned a weight of “1”).

Original periodic table listed elements in order of their atomic weights. This is not true today (see Iodine).

Mass number: Represented by the letter “A” . It is the sum of the protons and

neutrons in a nucleus. This number is rounded from atomic mass due to the fact that there are isotopes.

# neutrons = A - Z

Example - # of neutrons in Li = 6.941 - 3 = 3.941 rounds to 4

Atomic Mass:

The average mass of all of the isotopes of an element – is a number with a decimal – is always the larger number on the periodic table.

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Not all atoms of the same element are identical…

Isotopes: atoms of the same element with different masses because they have different

numbers of neutrons.

|Hydrogen Isotopes |Protium |Deuterium |Tritium (artificial and |

| | | |radioactive) |

|Protons |1 |1 |1 |

|Neutrons |0 |1 |2 |

|Electrons |1 |1 |1 |

|Mass |1 amu |2 amu |3 amu |

|% Abundance |99.85% |.15% |0% |

Some isotopes occur naturally. Most isotopes are produced artificially.

Ion : a charged atom. Atoms become charged by gaining electrons (become a negative

charge) or losing electrons (become a positive charge)

Counting protons, electrons, and neutrons:

Mass # = protons + neutrons

Neutrons = mass # - protons

Protons = mass # - neutrons

Atomic # = protons

Electrons = protons (unless the atom is charged)

+ Charge indicates the removal of an electron

- Charge indicates the addition of an electron

|Isotope |Protons |Electrons |Neutrons |Atomic # |Mass # |

|40K+1 | | | | |40 |

| |12 | |12 | | |

| | |36 |53 | |91 |

|-1 | |10 | | |19 |

|14C | | | | | |

|S-2 | | | | | |

|Na+1 | | | | | |

|Br | | | | | |

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8. Mass Spectrometry and Isotopic Abundance

The mass spec (chemist lingo) is used to detect, analyze, and identify unknown chemicals. Samples are vaporized, bombarded with electrons (in order to create + ions [positively charged particles due to a loss of 1 electron]), and placed in electrical and magnetic fields.

Due to differences in mass ( # of neutrons) or charge, the paths of the molecules curve based on their individual mass/charge. Heavier particles curve less as do smaller charge particles. This change in curvature causes the particles to land on different places on a detector.

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The mass spectrometer was invented in 1912. By 1922 it was discovered that there were 300 naturally occurring isotopes that existed out of the 92 elements known at that time. Used for identifying components of mixtures, analyzing pollution, and dating works of art.

9. The Atomic Weight Scale and Atomic Weights

Atomic weight is the average weight of all of the natural isotopes.

Originally H was the basis of all atomic masses and was given

the mass of 1.0. Later, chemists changed the standard to oxygen being

16.000 (which left H = 1.008). In 1961, chemists agreed that 12C is the

standard upon which all other masses are based.

Atomic weight is based on 12C = 12.0000 amu, so 1 amu = 1/12 mass of one 12C atom.

Atomic Mass:

Mass of Cl thought to be 35.5 times that of hydrogen.

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Today we know this isn’t true. It’s the weighted average of 2 isotopes:

75% 35 Cl and 25% 37 Cl

Analogy for weighted averages:

If your homework grade is 80.0 and your test grade is 95.0, then what is your average? Note: homework is worth 50.0% of your grade.

80.0 + 95.0 = 175.0 ( (175.0/2) = 87.5

or

[(80.0)(.500)] + [(95.0)(.500)] = 87.5

If your homework grade is 80.0 and your test grade is 95.0, then what is your average? Note: homework is worth 20.0% of your grade.

[(80.0)(.200)] + [(95.0)(.800)] =

16.0 + 76.0 = 92.0

Average atomic mass = [(%)(mass of 1st isotope)] + [(%)(mass of 2nd isotope)]......

Sample problem: find the average atomic mass of B

B11 = 80.20% B10 = 19.80%

[(11)(.8020)] + [(10)(.1980)] = 10.802

Sample problem: find the %’s of 2 isotopes of Carbon given the following

information:

average atomic mass = 12.0111 isotope 1 = 12 C , isotope 2 = 13 C

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Another example: Antimony consists of two naturally occurring isotopes: Antimony-121 (57.2500% abundance and an exact mass of 120.9038) and some other isotope. Calculate the percent and mass of the other isotope.

10. Electromagnetic Radiation

Pass out AP test equation sheet and show all applicable formulas

c = 3.00 x 108 m/s, h = 6.626 x 10-34 Jsec,

Useful conversion - 1 angstrom = 1 x 10-10 m

Electromagnetic Radiation: energy that travels in waves. Light is an example.

The electromagnetic spectrum:

gamma rays(x-rays(ultraviolet( visible (infrared(radar(microwaves(radio waves

1 x 10-16 m 1 x 10-6 m 1 x 104 m

The speed of light in a vacuum: 3.00 x 108 m/sec (186,000 miles/sec). It is a little

slower in air, but that measurement is still accurate to 3 sig figs.

Sir Isaac Newton: (English - 1600’s) suggests light is particles.

Christian Huygens: (Dutch - same time) suggests light is waves of energy.

Max Plank: (German - early 1900’s) revives the particle theory. Plank studied

light given off by hot bodies. He proposed light is given off in bundles of energy which he called quanta (photons). The amount of energy depends on the color of light.

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The modern theory is that light behaves both as particles and waves.

Continuous Spectrum:

Red versus Violet:

red light:

↑ λ = ↓ f = ↓ E

violet light:

↓ λ = ↑ f = ↑ E

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Frequency (f): the number of peaks that pass per second

measured in cycles per second

note: cycles per second = the Hertz (Hz) ( s-1 or 1/sec

Wave velocity (v): the distance a wave moves in one second

( = wavelength

So, red light has a longer wavelength and a lower frequency compared to violet.

Light as Energy: Planck derived a formula that expresses the energy of a photon

at any given frequency

Formula 1: E = hf E is the energy of a photon measured in joules

h is Plank’s constant (6.6 x 10-34 J/Hz) ( J∙s

f is frequency measured in Hertz ( s-1

Formula 2: v = f( f is frequency measured in Hertz

( = wavelength measured in meters

1,000,000,000 nm = 1 m

10,000,000,000 angstroms = 1 m

v = velocity (the speed of light (c) = 3.00 x 108 m/sec)

Combine formulas 1 and 2: E = ch

(

Ex1. The yellow light given off by sodium vapor street lights has a wavelength of 589 nm

(589 x 10-9 m or 5890 Angstroms). Calculate the frequency. Answer = 5.09 x 1014 s-1 (Hz)

Ex2. Lasers used in eye surgery have a frequency of 4.69 x 1014 s-1. What is the wavelength? Answer = 6.40 x 10-7 m

Ex3. Calculate the quantum energy of yellow light when the wavelength is 589 nm.

Answer = 3.36 x 10-19 J

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5-11 The Photoelectric Effect – Students should read this section

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11. Atomic Spectra and the Bohr Atom – Draw white light going through a prism (continuous spectrum). Glowing hot gas light passing through a prism ( bright line spectrum)

ROYGBIV: colors of the visible spectrum.

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Bright line spectrum (bls): frequencies of light give off by certain substances

when energy is added to them. Elements can appear to give off the same color

light, but each will have its own bls.

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The spectrum of an object is the variation in the intensity of its radiation at different wavelengths.

Objects with different temperatures and compositions emit different types of spectra. By observing an object's spectrum, then, astronomers can deduce its temperature, composition and physical conditions, among other things.

• A hot solid, liquid or gas, under high pressure, gives off a continuous spectrum.

• A hot gas under low pressure produces a bright-line or emission line spectrum.

• A dark line or absorption line spectrum is seen when a source of a continuous spectrum is viewed behind a cool gas under pressure.

The wavelength of the emission or absorption lines depends on what atoms are molecules are found in the object under study. The atoms or molecules exist depend on temperature and chemical composition..

Each atom or molecule exhibits a different pattern of lines (rather like a fingerprint or DNA signature).

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bls is used to determine the composition of stars and planets.

bls validates Bohr’s idea that electrons jump to different energy levels and give

off different wavelengths of light.

Ex. Hydrogen gas emits a green line with a wavelength of 486 nm. Calculate the energy of a single photon. Answer = 4.07 x 10-19J/photon

The Atomic spectra – relates to drawing “Bohr” atoms, show high energy and low energy jumps…

In 1913, Neils Bohr (Danish), stated that electrons occupy fixed paths or orbitals without giving off energy.

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Electrons far from the nucleus have higher energy.

This theory was deduced from flame tests and gas tubes.

Spectroscopy – the study of substances from the light they give off.

Spectroscope – the instrument used to break light into its component colors.

1 2 3 4 5 6 7

K L M N O P Q

Different levels: 1 through 7, K - Q

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Draw Lyman series – where all electrons return to level 1 (UV region) 1216 A, 1026 A, 973 A, 950 A)…see page 199 in text

Balmer series – where electrons fall to n = 2 (visible region) 6563 A, 4861 A, 4340 A…see page 199

To determine the energy of a moving object, use the de Broglie equation λ = h/mv (v = velocity) remind h has to change unit to kg(m2)/s and masses must correspond in unit

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To determine the energy of a moving electron, use the Rydberg equation:

1/ λ = R (1/n12 - 1/n22) where R = 1.097 x 107 m-1

or (even easier) AP test formulas sheet gives En = -2.178 x 10-18/n2 (unit is Joules)

Ex. Calculate the wavelength of light emitted (nm) when an electron falls from n = 4 to n = 2 in a

hydrogen atom. Answer = 485 nm

Option 1: 1/ λ = R (1/n12 - 1/n22) where R = 1.097 x 107 m-1 , n1 = 2, n2 = 4

Option 2: En = -2.178 x 10-18/n2

Option 3: See page 199…

13. The Wave Nature of the Electron – students should read this section

Classical mechanics: laws of motion - developed by Isaac Newton.

These laws apply to the macroscopic world brilliantly.

These laws also apply to the motion of atoms and molecules.

Einstein: theory of relativity shows Newton’s equations don’t work when objects approach the speed of light.

Bohr: electron “jumps” don’t fit into Newtonian physics as well. It only worked

for Hydrogen. Wave mechanics developed to help Bohr’s theory.

de Broglie (France) and Planck advance the theories of wave mechanics.

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14. The Quantum Mechanical Nature of the Atom

A few points before moving on to predicting quantum numbers.

1. Atoms and molecules can only exist in certain .

2. A change in energy state requires that an atom emit or absorb energy at a given

.

3. The allowed energy states of electrons can be described by a set of .

Electrons can be treated as small particles of matter, but more often it is more productive to treat them as waves – this leads to quantum mechanics.

Heisenberg uncertainty principle:

In 1927, Heisenberg (German) creates the Heisenberg uncertainty principle.

HUP: It is impossible to know both the speed and location of an electron at the same time. When light hits a particle, when you observe it, it changes the speed and location.

Quantum mechanics (equations) allows scientists to determine the probability of finding particles in certain places. This leads to the charge cloud model.

5-16 Atomic Orbitals 5-17 Electron Configurations 5-18 Paramagnetism and Diamagnetism and 5-19 The Periodic Table and Electron Configurations

Energy levels of the wave-mechanical model

Energy levels: principle: what shell, level, etc. n = 1,2,3...7

energy sub level - s, p, d, f

theoretical - g, h, i...

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Orbitals and Shapes of orbitals

|Sublevel |# of orbitals |e- per orbital |Max # of e- per sublevel |Shape of orbital |

| | | | | |

|s |1 | | |O |

| | | | |[pic] |

|p |3 | | |[pic][pic][pic] [pic] |

| |px, py, pz | | | |

| | | | | |

|d |5 | | | |

| |dxy, dyz, dxz, | | | |

| |dx2 – y2, dz2 | | | |

| | | | | |

|f |7 | | | |

| |(very complicated) | | | |

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- energy levels, sublevels, and total number of electrons per shell

|Energy Level (shell) |Sublevel(s) and Electrons |Total # of Electrons per Shell |

| | |2n2 |

|1 |s = 2 e- |2 |

|2 |s = 2 e- |8 |

| |p = 6 e- | |

|3 |s = 2 e- | |

| |p = 6 e- |18 |

| |d = 10 e- | |

|4 |s = 2 e- | |

| |p = 6 e- |32 |

| |d = 10 e- | |

| |f = 14 e- | |

|5 |s = 2 e- | |

| |p = 6 e- | |

| |d = 10 e- |50 |

| |f = 14 e- | |

| |etc… | |

|6 |s = 2 e- | |

| |p = 6 e- |72 |

| |d = 10 e- | |

| |etc… | |

|7 |s = 2 e- | |

| |p = 6 e- |98 |

| |etc… | |

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Bohr Model Examples (1 to 20):

| H | | |

|+ ) | | |

|1 | | |

| | | |

| |+ ) ) | |

|Oxygen |2 2 |1s22s22p6 |

| |4 | |

| |6 | |

| | | |

| |+ ) ) ) ) | |

|Potassium |2 2 2 1 |1s22s22p63s23p64s2 |

| |6 6 | |

| |8 8 | |

| | | |

| | | |

|Copper | | |

| | | |

| | | |

| | | |

| | | |

| | | |

| | | |

|Silver | | |

| | | |

| | | |

| | | |

| | | |

Draw diagonal diagram here:

1s2

2s2 2p6

3s2 3p6 3d10

4s2 4p6 4d10 4f14

5s2 5p6 5d10 5f14

6s2 6p6 6d10

7s2 7p6

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- Kernal Electrons: inner shell electrons + the nucleus

- Valence Electrons: outer shell electrons

- example Radium (Bohr model, electron configuration, orbits:

+ ) ) ) ) ) ) )

2 O 2 O 2 O 2 O 2 O 2 O 2 O

8 8 8 8 8

6 ∞ 6 ∞ 6 ∞ 6 ∞ 6 ∞

8 8

10 10 10

18 18

14

32

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2

Pauli Exclusion principle - electrons occupying the same orbital must have

opposite spin.

electron spin:

spin - clockwise or counterclockwise

Hund’s rule ( better known as the Stinky Bus Rule) - before any second

electron can be placed in a sub level, all the orbitals of that sub level must

contain at least one electron.

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Orbital notation is another way to represent electron arrangement in atoms.

Examples: Bohr Model, Electron Configuration, and Orbital Notation:

|Element |Bohr Model |Electron Configuration |Orbital Notation |

| | | | |

| |+ ) ) ) | | |

|Argon |2 2 2 |1s22s22p63s23p6 | |

| |6 6 | |1s 2s 2p 3s |

| |8 8 | |3p |

| | | | |

| | | | |

|Manganese | | | |

| | | | |

| | | | |

| | | | |

| | | | |

| | | | |

- example Iodine (Bohr model, electron configuration, orbits, and orbital notation:

+ ) ) ) ) )

2 O 2 O 2 O 2 O 2 O

8 8 8 8

6 ∞ 6 ∞ 6 ∞ 5 ∞

8 7

10 10

18 18

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p

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- Exceptions to the Aufbau Principle

- Cr is 4s1 3d5 not 4s2 3d4

- Other exceptions: Mo, Ag, Au, Cu, and Cr

- students should be able to identify these elements simply based on how many

electrons they have.

Do the orbital notation for Copper:

Significance of electron configurations:

- Valence shell electrons - outermost electrons involved with bonding

- for n = 5, pattern is very complicated - no atom has more than 8 valence electrons

- Noble gases - 8 valence electrons - least reactive of all elements

- kernel - part of the atom exclusive of valence electrons (includes the nucleus)

- example - sodium kernel = nucleus + 1s22s22p6

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Examples: (find the outershell electron configuration, shorthand notation, and Lewis dot diagram – N,N,E,S,W,E,S,W

Do outershell (s, f, d, p) for: Na, S, V, Ti, and Bi

|Element |Outershell Electron Configuration |Shorthand Notation |Lewis Dot Diagram |

| | | |(main column #) |

| | | | . . |

|Oxygen |2s22p4 |[He] 2s22p4 |.O: |

| | | |. |

| | | | . . |

|Chlorine |3s23p5 |[Ne] 3s23p5 |.Cl: |

| | | |. . |

| | | |. . |

|Iron |4s23d6 |[Ar] 4s23d6 |Fe |

| | | | . . |

|Cobalt |4s23d7 |[Ar] 4s23d7 |Co |

| | | |. |

|Potassium |4s1 |[Ar] 4s1 |K |

| | | | |

|Strontium | | | |

| | | | |

|Silver | | | |

| | | | |

|K+1 | | | |

| | | | |

| | | | |

|O-2 | | | |

| | | | |

| | | | |

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electron configurations for elements in the excited state:

- all previous examples have been for ground state electrons

- example: oxygen ( ground state) - 1s22s22p4

oxygen (excited state) - 1s22s22p35s1 or 1s22s12p5 , etc....

Paramagnetism/Diamagnetism – proof of Hund’s rule

Paramagnetic – sample is weakly attracted by a magnetic field – increases with increased # of unpaired electrons. More prevalent in transition metals.

Diamagnetic – sample is very weakly repelled. Happens with Helium, Magnesium, etc.

Ferromagnetism – permanent magnetism* associated with Fe, Co, Ni where unpaired electrons align their spin.

*If raised above a certain temperature (Curie Temperature), these substances can be demagnetized.

One would think Oxygen would be diamagnetic – it’s actually paramagnetic – we’ll see why in chapter 9.

15. Quantum Numbers

Quantum numbers: used to indicate the location of an electron

Pauli exclusion principle - 2 electrons in the same orbital must have opposite spins.

Or… No two electrons in an atom can have identical sets of four quantum numbers.

|Quantum # |Symbol |Identifies the . . . |Value |

|Principle |n |Energy Level | |

|Azimuthal |l |Energy Sublevel | |

| | |(s,p,d,f) | |

|Magnetic |m |Specific Orbital | |

| | | | |

| | | | |

|Spin |s |Direction of Spin | |

| | | | |

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sample problem - “what are the possible quantum numbers for the electrons of Carbon?”

|Level |n |l |m |s |notes |

|1 s2 |1 |0 |0 |+1/2 | |

| |1 |0 |0 |-1/2 | |

|2 s2 |2 |0 |0 |+1/2 | |

| |2 |0 |0 |-1/2 | |

|2 p2 |2 |1 |+1 |+1/2 |X |

| |2 |1 |+1 |-1/2 | |

| |2 |1 |0 |+1/2 |Y |

| |2 |1 |0 |-1/2 | |

| |2 |1 |-1 |+1/2 |Z |

| |2 |1 |-1 |-1/2 | |

sample problem - “write the possible quantum numbers for one of the d electrons of Iron”

sample problem - “which of the following are wrong quantum # combinations and why?”

1 1 0 +1/2

3 3 -1 -1/2

2 1 +2 +1/2

Notation for electron configurations:

- Electron configuration - a representation of the arrangement of electrons in an

atom

- example- 2py2

2 = principle quantum #

p = azimuthal

y = magnetic

2 stands for the number of electrons in that orbital

example - oxygen 1s22s22p4 or 1s22s22px22py12pz1

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