Electron Arrangement in Atoms Chapter 4. Development of Atomic Model(s)
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Transcript of Electron Arrangement in Atoms Chapter 4. Development of Atomic Model(s)
Electron Arrangement in Atoms
Chapter 4
Development of Atomic Model(s)
• Light is thought to be a wave– Since the 1900’s
• Newton thought light consists of particles.– Interestingly, the current view of atoms, the
quantum mechanical model of the atom, came about from the study of light.
Electromagnetic Spectrum: all the forms of light - acting like waves
Particle: would cause 2 linesWaves: interference causes multiple lines
Light Waves
• Wavelength ()– distance between the crests. (unit =nm)
• Amplitude– wave’s height from zero to the crest.
• Frequency ()– Is the number of waves to pass a given point per unit of
time.– The unit of frequency is cycles/second which is called the
hertz (Hz)
• C= speed of light, a constant
• c = wavelegth x frequency or c = • The product of frequency and wavelength
always equals a CONSTANT (c), the speed of light.
As wavelength increases, frequency decreases
• ALL EMS waves travel at c.
Which wave has greatest wavelength Greatest amplitude? How do the frequencies compare?
A
B
C
D
E
Electromagnetic Spectrum• As wavelength decreases, frequency increases
– High frequency waves = gamma– Low frequency waves = radio
• As frequency increases, energy increasesQuantum: min amount of energy that can be lost by an electron
Visible light is a small part of the ES• Parts of visible light: ROYGBIV• When visible light passes through prism, it can be
separated
Infrared waves: Low frequency/high wavelengthUltraviolet waves: High frequency/low wavelength
So…infrared is LOWER in energy than UV
atomic emission spectrum: the discrete lines created when an element emits distinct frequencies of light- unique to each element (like a fingerprint) - light is emitted when electrons move energy levels
from an exited state to a ground state • When atoms absorb energy, e- move into higher energy levels…
but they don’t stay there…these e- then lose energy (a quanta) and emit light when they return to their lower energy levels.
Photon: quantum “particle” of energyQuantum: min amount of energy that can be lost by an electron
Check point• How does the electromagnetic spectrum differ
from the atomic emission spectrum?
– Shows spectrum of all light vs distinct light given off my certain atoms?
The Dual Nature of Light• Is light a PARTICLE or a WAVE?• Quantum of light are called photons and
behave like PARTICLES but have WAVE properties– Light has BOTH qualities (Thx Einstein)– Light is a particle that travels like a wave!– Not all scientists believed
this at first…now it is accepted as FACT
De Broglie won the Nobel Prize for his studies on the wave nature of matter, changing the way the world saw quantum physics.
Classical Physics: Deals with average size objects traveling at average speeds…
Quantum Mechanics/ Physics: deals with atomic size objects traveling at incredibly high frequencies (speeds)
• Watch this physicist trying to explain what De Broglie opened up with this theory– https://www.youtube.com/watch?v=JIGI-eXK0tg#t=15 start
at 1:12
Bohr Model• Each possible e- has a fixed energy called an
energy level …like steps on a ladder.– Switching levels requires a gain or loss of energy in
the form of photons.• Gain energy to go up a level (farther from nucleus)• Loss energy to go down a level
– High energy levels are furthest from the nucleus.
Drawing an atom with Bohr model:• Can abbreviate what’s in the nucleus• First orbital holds only 2 e-
• all other orbitals hold 8 e-
Draw model for boron:
how electrons can be excited and move to excited energy states which is higher than their original ground state
The Heisenberg Uncertainty Principle• it is impossible to know exactly both the velocity and the position of a particle at the same time.
• If you determine one, it affect and changes the other, so you can’t ever calculate both at the same time. Ex: measuring position changes its velocity making its location uncertain.
• Based on Schrödinger equation: the same Schrödinger who came up with the dead and alive cat paradox to explain the Copenhagen interpretation of quantum mechanics
https://www.youtube.com/watch?v=QisnPsu7_Uk
• Erwin Schrödinger (1926) used theoretical mathematical calculations to describe e- motion.
• This model used mathematical equations describing the behavior of the e- in a hydrogen atoms.
• Modern quantum mechanics deals with the laws of motion which govern the behavior of atomic and subatomic particles.
Quantum Mechanical Model• Location of e- shown by a cloud.• mathematical theory that predicts the probability of
electron in certain area
Check point
• How does the quantum mechanical model differ from the Bohr model of an atom?
Quantum Mechanical Model cont. Includes 4 quantum #s describing the position of electrons • The quantum mechanical model of the atom does NOT
involve an exact path the e- takes around the nucleus (like the Bohr model does).
• Rather, it is based on the likelihood or probability of where the e- will be in various locations around the nucleus.
Valence e-: electrons in outermost orbital• have the most
potential energy
Quantum Numbers SummaryPrinciple: n
Energy level
1-7 periods
Angular: l
Orbital type
s, p, d, f,
Magnetic: m
Orbital orientation
--------------
Spin Electron spin direction
+1/2 or -1/2
4 quantum numbers:1. Principle Quantum #: n
- Describes energy level occupied by e-
- represented by the periods 1-7 (horizontal rows)
the larger the number, the farther the valence e- from the nucleus, the more energy (potential)
Valence e-: electrons in outermost orbital• have the most potential energy
4 quantum numbers:2. Angular quantum #: l
- Indicates shape of orbital: s, p, d, f
4 quantum numbers:3. Magnetic Quantum #: m
- orientation of orbital around nucleus
4 quantum numbers:4. Spin Quantum #: +1/2 or -1/2
- Spin of an electron
Quantum Numbers SummaryPrinciple: n
Energy level
1-7 periods
Angular: l
Orbital type
s, p, d, f,
Magnetic: m
Orbital orientation
--------------
Spin Electron spin direction
+1/2 or -1/2
https://www.youtube.com/watch?v=accyCUzasa0
Electron configurations• Electron configurations are similar to postal
“zipcodes”.– Written distribution of electrons in their orbitals
• ExamplesHydrogen has 1 electron: 1s1
He has 2 electrons: 1s2
Li has 3 electrons: 1s2 2s1
Be has 4 electrons: 1s2 2s2
B has 5 electrons: 1s22s22p1
Sublevels: s, p, d, f • Within each energy level, there are sublevels
– s sublevel- consist of e- from groups 1 and 2– p sublevel- consist of e- from groups 13-18– d sublevel- consists of e- from groups 3-12– f sublevel- consists of e- from the inner transition
metals S
pd
f
Orbitals: hold a maximum of 2 electrons.• Each sublevel can be broken down into
orbitals:– s sublevel: 1 orbital : 2 e- total– p sublevel: 3 orbitals: 6 e- total– d sublevel: 5 orbitals: 10 e- total– f sublevel: 7 orbitals: 14 e- total
Rules for writing electron configurations1.The number of electrons in an atom is equal to
the atomic number; unless told otherwise. 2.Electrons fill the lowest energy level/sublevel
before moving to a higher energy level/sublevel. Order: s, p, d, f
3.All orbital must have one electron before any getting a second electron
4.An orbital cannot take more than 2 electrons, and they must have opposite spins.
P: 1s22s22p63s23p3
Orbital Diagram
Electron Configuration
Ex P: 1s22s22p63s23p3
RIGHTWRONG
General Rules• Hund’s Rule = “empty bus seat rule”
– Within a sublevel, place one electron per orbital before pairing them.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
S: 1 orbital : 2 e- totalp: 3 orbitals: 6 e- totald: 5 orbitals: 10 e- totalf: 7 orbitals: 14 e- total
O 8e-
• Orbital Diagram
• Electron Configuration
1s2 2s2 2p4
Notation
1s 2s 2p
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
O15.9994
8
S: 1 orbital : 2 e- totalp: 3 orbitals: 6 e- totald: 5 orbitals: 10 e- totalf: 7 orbitals: 14 e- total
Filling Rules for Electron Orbitals come from these principles:
Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.
Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.To occupy the same orbital, two electrons must spin in opposite directions.
Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.
*Aufbau is German for “building up”
Spin Quantum Number (ms): +1/2, -1/2
Practice:
1.Be
2.N
3.Ni
4.He
5.S
1.Be2.N3.Ni
4.He5.S
Practice: S
p
d
f
3d456
4f5f
Practice:
1.Be
2.N
3.Ni
4.He
5.S
1.1s2 2s2
2.1s2 2s2 2p3
3.1s2 2s2 2p6 3s2 3p6 4s2 3d8
(or 1s2 2s2 2p6 3s2 3p6 3d8
4s2)
4.1s2
5.1s2 2s2 2p6 3s2 3p4
Orbital Filling
Element 1s 2s 2px 2py 2pz 3s Configuration
Electron ConfigurationsElectron
H
He
Li
C
N
O
F
Ne
Na
1s1
1s22s22p63s1
1s22s22p6
1s22s22p5
1s22s22p4
1s22s22p3
1s22s22p2
1s22s1
1s2
• Shorthand Configuration
S 16e-
Valence ElectronsCore Electrons
S 16e- [Ne] 3s2 3p4
1s2 2s2 2p6 3s2 3p4
Notation
• Longhand Configuration
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
S32.066
16
Abbreviate configuration using noble gas notation: Ex: 1. Zn 2. Sn
1.Be2.N3.Ni
4.He5.S
Practice: S
p
d
f
3d456
4f5f
Practice:
1.Be
2.N
3.Ni
4.He
5.S
1.1s2 2s2
2.1s2 2s2 2p3 = [He] 2s2 2p3
3.1s2 2s2 2p6 3s2 3p6 4s2 3d8
[Ar] 4s2 3d8
4.1s2
5.1s2 2s2 2p6 3s2 3p4
abbreviated notation
Bozeman orbitals: start at 6:49 https://www.youtube.com/watch?v=2AFPfg0Como
This fills the valenceshell and tends to givethe atom the stabilityof the inert gasses.
The Octet RuleAtoms tend to gain, lose, or share electrons until they have eight valence electrons.
8
ONLY s- and p-orbitals are valence electrons.
• The groupA number equals the number of electrons in its outermost energy level
(this will be more important later on…)
• So all Group 1A elements have 1 electron in the outer shell
Stability: to gain stability, atoms will gain or lose electrons, forming an ion with the same # of e- as the closest noble gas
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
• Electron Configuration Exceptions to gain stability:
– Copper
EXPECT: [Ar] 4s2 3d9
ACTUALLY: [Ar] 4s1 3d10
– Copper gains stability with a full d-sublevel.
1.Be2.N3.Ni
4.He5.S
Practice: S
p
d
f
3d456
4f5f
1
2
3
4 5
6
7
Stability• Ion Formation
– Atoms gain or lose electrons to become more stable to be isoelectronic with the Noble Gases.
Group 1-2 lose elections = cations Group 3A to 6A gain electrons= for anions (except He!)
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
form cations! form anions!
METALS form CATIONS
• Heavy metal cats…ROCK ON!!
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
• Ion Electron Configuration– Write the e- configuration for the closest Noble
Gas• EX: Oxygen ion O2- Ne
O2- 10e- [He] 2s2 2p6
Gilbert N. Lewis• Lewis suggested that atoms with
fewer than 8 valence e- bond together to share e- and complete their valence shells.
• He noticed that many elements are more stable when they have eight electrons in their outer shell.
Lewis Dot Structure• Simple dot diagram of valence electrons
Drawing Lewis Structures
• Practice writing Lewis Dot Structures for the following elements:
1. K
2. Mg
3. Ga
4. C5. P6. Se7. Br8. Ar
1. K2. Mg3. Ga
Practice: S
p
d
f
4.C5.P
6.Se7.Br
8.Ar
Ch 5: Periodic Trends• Valence electron rules, electron configuration,
and cation and anion formation all are part of the trends that govern the periodic table!
Electrons are the Key!
Valence electrons: group elements have same number of valence electrons
• The reason elements in the same family(group) have the same chemical and physical properties is because they have the same number of electrons in their outer shell
Ch 5
Ch 5 Summary of Trends
Increasing Electronegativty
Increasing Ionization Energy
Decreasing Atomic RadiusD
ecre
asin
g I
on
izat
ion
En
erg
y
Dec
reas
ing
Ele
ctro
neg
ativ
ity
Incr
easi
ng
Ato
mic
Rad
ius
Atomic Radius (Size)• decrease as you go L to
R across period.– Same energy level but:– Add p+ and e-, increase
nuclear charge, pulls in orbitals closer to the nucleus
• increase as you go down a group– Electrons being added to
outer orbital (increasing principle energy level)
Ion Size
• Cations are smaller in size than the neutral element.
• Anions are larger in size than the neutral element.
Ionization Energy
• Ionization Energy (IE): The energy required to remove and electron from a gaseous atom.– Remove 1st electron = 1st IE– Remove 2nd electron = 2nd IE
Ionization Energy cont.• IE generally increases as you move L to
R across the period.– Harder to remove an electron as you go L to
R because of greater attraction to nucleus =Shielding effect
• IE generally decreases as you go down a group. (first IE only)– Atom gets bigger, outermost e- farthest from
nucleus, easy to be removed.
Electronegativity• Electronegativity: The tendency for
atoms to attract electrons when they are chemically combined. – Do they share electrons equally? (remember
polar water bonds1?)
Electronegativity Trends• increases as you go L to R across the period
– Elements want to be like noble gases!
• decreases as you go down a group.
The most electronegative element is Fluorine.
– Noble gases have no electronegativity as they already have a full valence shell.
Summary of Trends
Increasing Electronegativty
Increasing Ionization Energy
Decreasing Atomic RadiusD
ecre
asin
g I
on
izat
ion
En
erg
y
Dec
reas
ing
Ele
ctro
neg
ativ
ity
Incr
easi
ng
Ato
mic
Rad
ius
Is that it!