Chapter 4

Arrangement of Electrons in an Atom

4.1 Refinements of the atomic model

Models of the atom so far: Dalton – atoms are like little “bb’s” - then the

electron gets discovered Thomson – atom is like a charged “bb” Rutherford - Gold foil experiment – hollow charged

“bb” Bohr model of the atom (1913) – Neils Bohr – Danish

Physicist The Bohr model of the atom comes from the idea that

light is waves of energy

http://web.visionlearning.com/custom/chemistry/animations/CHE1.2-an-atoms.shtml

The Bohr Atom (1913)

All the positive charge was in the nucleus Electrons orbited the nucleus much like planets orbit the

sun (at fixed distances) The closer the electrons to the nucleus, the less energy it

has. The farther the electron is from the nucleus, the more

energy it has.

The Electromagnetic Spectrum

Visible light, x-rays, ultraviolet radiation, infrared radiation, microwaves and radio waves are all part of the electromagnetic spectrum

The Electromagnetic Spectrum

The spectrum consists of electromagnetic radiation – energy that travels like a wave

Waves can be described by the wave equation which includes velocity (c = speed of light), wavelength (λ) and frequency (ν).

Wavelength (definition) = the distance between peaks of a wave

Light through prism leads to high energy (violet) low energy (red)

The Electromagnetic Spectrum

ROYGBIV - colors of the visible spectrum Bright Line Spectrum (BLS) – caused by e- emitting

energy as they return to lower energy levels energy level.

heat sodium - yellow light 2 c heat lithium - red light elements can appear to give off the same color light, but

each will have its own BLS BLS - used to determine identity of an element BLS - validates Bohr’s idea that electrons jump to different

energy levels and give off different wavelengths of light

The Electromagnetic Spectrum

Light from the sun (white light) appears as a continuous spectrum of light.

Continuous Spectrum of Light (definition) = There are no discrete, individual wavelengths of light but rather all wavelengths appear, one after the other in a continuous fashion

Spectroscopy (definition) = the study of substances from the light they emit.

We will use spectroscopes (An instrument that splits light into its component colors) and flame tests to study elements because each element emits a different spectrum of light when exited .

Birght Line Spectrum

Bohr proposed that the energy possessed by an e- in a H- atom and the radius of the orbit are quantized (bls) Quantized (definition): a specific value (of energy)

The ramp is an example of a continuous situation in which any energy state is possible up the ramp

Like a set of stairs, the energy states of an electron is quantized – i.e. electrons are only found on a specific step

Bohr’s Energy Absorption Process Light or energy excites an e- from a lower

energy level (e- shell) to a higher energy level

These energy levels are “ quantized “ (the e- cannot be in between levels), the e- disappears from one shell and reappears in another

This absorption or excitation process is called a quantum leap or quantum jump

Bohr’s Energy Absorption Process Ground State Analogy = a spring and two balls

This is an energy emission process and

what we observe in the hydrogen line spectrum

Both the atom and e- now have higher

energyThe e- absorbs energy in the ground state and is

excited to a higher level

Bohr’s Energy Absorption Process When energy is added, the electron is found in the

“excited state.” The Excited State (definition) = an unstable, higher

energy state of an atom An illustration of Bohr’s Hydrogen atom (from ground to

excited state):

Bohr’s Energy Absorption Process The atomic line spectral lines - when an e- in an excited

state decays back to the ground state

The electron loses energy, light (colors) is emitted and the e- returns to the ground state

This is another illustration of bls.

The Bohr Model - Summary

1. When an atom absorbs energy, its electrons are promoted to a higher energy level. When the electron drops back down, energy is given off in the form of light.

2. Each distance fallen back is a specific energy, and therefore, a specific color.

3. Since electrons can fall from level 5 to 4, 5 to 3, etc., many colors are produced.

The Bohr Model - Summary

Bohr also predicted that since electrons would occupy specific energy levels and each level holds a specific number of electrons

The maximum capacity of the first (or innermost) electron shell is two e-.

Any element with more than two e-, the extra e- reside in additional electron shells.

The Bohr Model - Summary

Group IA

Lithium

VIA

Oxygen

VIIA

Fluorine

VIIIA

Neon

IA

Sodium

Electron Configurations for Selected Elements

The number of e- per shell = 2n2 (where n is the shell number)

Bohn Models

Draw Bohr Diagrams for the elements 1-10. Save room to draw them short hand also

Short Hand Bohr Model

Write the symbol of the element Use a ) to represent each shell Write the # of e- in each shell Ex. Element Short-Hand e-

Configuration

Hydrogen H )1e-

Lithium Li )2e- )1e-

Fluorine        F  )2e- )7e- Sodium Na )2e- )8e- )1e-

The Truth About Bohr Models At atomic # 19 (z = 19), there is a a break in the pattern. One

would expect that energy level #3 would continue to fill up. However, the next two electrons go into the next energy level. Look at K and Ca:

IA VIIIA

H+ ) 1 IIA IIIA IVA VA VIA VIIA

He+ ) 2

Li+ ) ) 2 1

Be+ ) ) 2 2

B+ ) ) 2 3

C+ ) ) 2 4

N+ ) ) 2 5

O+ ) ) 2 6

F+ ) ) 2 7

Ne+ ) ) 2 8

Na+ ) ) ) 2 8 1

Mg+ ) ) ) 2 8 2

Al+ ) ) ) 2 8 3

Si+ ) ) ) 2 8 4

P+ ) ) ) 2 8 5

S+ ) ) ) 2 8 6

Cl+ ) ) ) 2 8 7

Ar+ ) ) ) 2 8 8

K+ ) ) ) ) 2 8 8 1

Ca+ ) ) ) ) 2 8 8 2

The Truth Continued……

So, there is a relationship between the main column # and the number of outershell electrons.

Column # = the number of valence electrons And, there is a relationship between the row # and

the number of energy levels. Row # = the number of shells The Bohr model truly works well for the H atom only

for elements larger than H the model does not work.

Bohr Summed Up

Bohr made 2 huge contributions to the development of modern atom theory He explained the atomic line spectra in terms of electron

energies He introduced the idea of quantized electron energy

levels in the atom

The Bohr atom lasted for about 13 years and was quickly replaced by the quantum mechanical model of the atom. The Bohr model is a good starting point for understanding the quantum mechanical model of the atom

Do ws# 1, question 1- Use short-hand configuration

4.2 Quantum Numbers and Atomic Orbitals

4.3 Electron Configuration

Quantum Numbers & Atomic Orbitals

The Bohr model describes the atom as having definite orbitals occupied by electron particles.

Schrödinger (1926) introduced wave mechanics to describe electrons – proved Bohr’s Model to be a lie Based his idea that electrons behaved like light

(photons). Electrons show diffraction (interference)

properties like light. Treats electrons as waves that are found in

orbitals. Orbitals (definition) = clouds that show region of

probable location of a particular electron.

Wave Mechanical Model

The Bohr model really is the wave mechanical model

There are many types of orbitals – we can see them on the periodic table

Subatomic Orbitals

Type # sublevels Total # e Shape

s 1 2 sphere

p 3 6 peanut

d 5 10 dumbbell

f 7 14 flower

S P D

Quantum Numbers

An electron’s address principle (n): what shell, level, the e- is in n =

1,2,3...7

azimuthal (l): energy sub level - s, p, d, f

magnetic – orientation of orbital about the nucleus (s has only 1, p has 3, etc.)

spin - clockwise or counterclockwise (+1/2 or -1/2)

Label Your Periodic Tabel

On your periodic table, shade azimuthal s,p,d,f blocks different colors

Label the principal quantum numbers…1-7

Label the valence electrons across the top

Electron Configuration

Electron Configuration - a representation of the arrangement of electrons in an atom

Electron Configuration

Examples of electron Configuration 1. Li 1s22s1

2. C 1s22s22p6

principle azimuthal

# of e- in that shell

Electron Configuration

Take note that after 4s is filled, 3d is than filled before 4p.

…… 6s than 4f than 5d than 6p When writing out the electron

configuration, always write your numbers in numerical order Y 1s22s22p63s23p64s23d104p65s24d1 – NO!

Y 1s22s22p63s23p63d104s24p64d15s2

Electron Configuration

Examples: Be

O

Ca

Mn

Electron Configuration

Examples Pb

Os

Electron Configuration

Short Hand Write the name of the last noble gas Write the electron config. that follows

Ex. Fe [Ar]3d64s2

Exceptions Cr [Ar] 3d54s1

Cu, Ag, Au- all s’s donate 1 e- to make the d orbital full

Cu [Ar] 4s13d10

Orbital Notation Electrons enter orbitals in a set pattern. For the most

part, they follow these rules: 1) The Aufbau Principle - electrons must fill lower

energy levels before entering higher levels.

Orbital Notation

Orbitals are like "rooms" within which electrons "reside".

The s subshell has one s-orbital. The p subshell has

three p-orbitals. The d subshell has 5 and f has 7.

Each orbital can hold at most 2 electrons

Orbital Notation

2. Hund’s Rule (better known as the 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 – spread out the e- before pairing them up.

3. Pauli Exclusion Principle - electrons occupying the same orbital must have opposite spin.

See a good online illustration at http://www.avogadro.co.uk/light/aufbau/aufbau.htm

Orbital Notation

H1s

1s 2s 2p

F

Orbital Notation

Examples: Li F

Na Sc

Orbital Notation

We can also do shorthand orbital notation (outer shell only)

Ca N

Fe

Ag [Kr] 4d105s1

Ag [Kr]

4d 5s

Significance of Electron Configurations Valence shell electrons - outermost electrons involved

with bonding no atom has more than 8 valence electrons Noble gases - 8 valence electrons – least reactive of all

elements Lewis Dot structures: NSEW (cheating) also show

correct way, count to 8

Lewis Dot Structures