CHAPTER 5: ELECTRONS IN ATOMS

• Rutherford’s model of the atom did not explain how the electrons occupy the space surrounding the nucleus.

Recall:

• In the early 1900’s, investigations using the relationships that exist between light and electrons will reveal a more accurate model of the atom.

Model ofThe Atom

Opposites attract….right??

Emission Spectrum

• atoms can exist in low energy states called ground states or at high energy states called excited states.

• when excited atoms return to ground states, it gives off energy in the form of electromagnetic radiation.

• all elements have a characteristic emission spectra that corresponds to the amount of energy given off.

EmissionSpectra

Ground &ExcitedStates

Probability of Finding an Electron in a Hydrogen Atom.

Niels Bohr linked photon emission with an atom’s electrons.

• electrons exist in fixed orbits and never in between.

• his model only explains the emission of Hydrogen atoms.

The Quantum Model of the Atom

Louis deBroglie suggested that electrons be considered waves confined to the space around the nucleus.

Flame TestEmission

Flame TestExplanation

NielsBohr

The Bohr Model ---it’s reallynot boring

• investigations also confirmed that electrons can be bent or diffracted.

• electrons, like waves, are capable of interference.

Wave PropertiesOf Electrons

Heisenberg Uncertainty Principle: it is impossible to determine simultaneously both the position and velocity of an electron.

Erwin Schrodinger’s wave equation and Heisenberg’s Uncertainty Principle laid the foundation for the quantum model of the atom.

• electrons do not travel in specified orbits.

• instead electrons exist in regions called orbitals.

Quantum Modelof the Atom

Heisenberg

Electron Configurations

“The arrangement of electrons in an atom”

Rules Governing Electron Configurations

1. Aufbau Principle: an electron occupies the lowest-energy orbital that can receive it.

2. Pauli Exclusion Principle: no two electrons in the same atom can have the same set of four quantum numbers.

QuantumSubshells

3. Hund’s Rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron.

• all electrons in singly occupied orbitals must have the same spin.

Orbital Notation

• a way to illustrate what orbitals the electrons occupy in an atom.

ie: Nitrogen – 7 electrons

Electron Configurations

• electrons configurations (shorthand method) are a method of noting how many electrons are in an atom and where they are located.

Order of Filling

Maximum amount of electrons:

s-orbital 2 electrons

p-orbital 6 electrons

d-orbital 10 electrons

f-orbital 14 electrons

QuantumSubshells

Example:

Determine the electron configuration for the following elements:

Lithium 3 electrons

1s22s1

Nitrogen 7 electrons

1s22s22p3

Neon 10 electrons

1s22s22p6

In order to write a noble gas notation simply include the previous noble gas (in brackets) and list the electron configuration from the next element from the noble gas to the element you are determining the configuration for.

Write the noble gas notation for Sodium

[Ne]3s¹

Write the noble gas notation for Sulfur

[Ne]3s²3p⁴

Noble Gas Notation

• in order to write long configurations much easier we can use a noble gas configuration.

Properties of Light

• Light behaves both as a wave and a particle.

Light’s Wavelike Properties

• visible light is one kind of electromagnetic radiation.

Light asa Wave

Electromagnetic Spectrum

•Prism

• all electromagnetic radiation moves at the speed of light (3.00 x 108 m/s)

Properties of Waves

1. Wavelength: () the distance between corresponding points on adjacent waves.

2. Frequency: () the number of waves that pass a given point in one second.

• Frequency is measured in Hertz (Hz).

Speed of Light Equation

What is the wavelength of radiation with a frequency of1.5 x 1013 Hz?

Wavelength x 1.50 x 1013Hz = 3.00 x 108m/s

Wavelength = 3.00 x 108m/s / 1.50 x 1013Hz

Wavelength = 2.00 x 10-5m

The Photoelectric Effect

• a phenomena involved with the emission of electrons from a metal when light shines on a metal.

• light of only a certain frequency enabled this event to occur light behaves like streams of particles!

PhotoelectricEffectMax Planck

• suggested that energy could only exist in discrete lumps called quantum.

E = h

h = 6.626 x 10-34 JsPlanck andBlackbody Radiation

Einstein also proposed that light delivers its energy in chunks called photons.

Therefore, in order for an electron to be ejected (photoelectric effect) , the electron must be struck by a single photon, of certain energy, to knock the electron loose.

One more question……….Are You Ready?

The threshold photoelectric effect in tungsten is produced by light of a wavelength 260 nm. Give the energy of a photon of this light in joules.

Find frequency!

λ x ν = C 260 nm x 1m/109 nm = 2.6 x 10-7 m

2.6 x 10-7 m x ν = 3.0 x 108 m/s

v = 1.15 x 1015/sec

E = h x ν E = 6.626 x 10-34Js x 1.15 x 1015/sec

E = 7.6 x 10-19 J