ARRANGEMENT OF ELECTRONS IN ATOMS

Chapter 4

THE DEVELOPMENT OF A NEW ATOMIC MODEL

Light has characteristics of both particles and waves

Electromagnetic radiation – a form of energy that exhibits wave-like behavior as it moves through space

Electromagnetic spectrum

ELECTROMAGNETIC SPECTRUM = wavelength (units of meters) = frequency (units of Hertz,

Hz, ) = speed of light (3.00 x 108 m/s) = Planck’s constant (6.626 x 10-34 Js)

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EMISSION AND ABSORPTION Ground state – the lowest energy state of an atom

Excited state – when an atom has a higher potential energy than it does at its ground state

Balmer series – represents visible light region

Lyman series – represents the UV region

Paschen series – represents the infrared series

EMISSION LINE SPECTRA When a narrow beam of light is passes through a prism,

it separates into specific colors in the visible spectrum. Each element has a signature spectrum.

EMISSION SPECTRA In addition to indicators in the visible spectrum,

energy can be detected in the UV region as well as the infrared region. This has been scientifically observed.

What wavelengths represent indicators in the UV range? What about the IR range?

EMISSION SPECTRAScientists had predicted that emission

spectra would be on a continuous spectrum.

Is this what was scientifically observed?

BOHR’S MODEL AND TRANSITION STATES Bohr’s Model helped to explain the quantum energy levels

of the atom

When a photon is absorbed, the electron gains enough energy to move to an outer energy level.

When an electron loses energy (in the form of a photon), energy is released.

ELECTRON ENERGY TRANSITIONS

PHOTOELECTRIC EFFECT Recap – what is the relationship

between energy and frequency of a wave?

What is the visible region of light? Which colors in the visible spectrum carry more energy?

http://www.youtube.com/watch?v=0qKrOF-gJZ4

HOMEWORK P.97, #1-6 (2 days to complete)

QUANTUM MODEL OF THE ATOM Electrons have

wave-like properties Investigations from

the photoelectric effect and hydrogen’s emission line spectra determined that light acts as both a wave and a particle.

ELECTRONS Electrons have interference patterns

Constructive interference Destructive interference

DOUBLE SLIT EXPERIMENT http://www.youtube.com/watch?v=DfPe

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HEISENBERG Heisenberg’s Uncertainty Principle – you

cannot know the instantaneous position and velocity of an electron (or any other particle)

When you have extremely precise information about x, your uncertainty for p goes through the roof.

x – represents positionp – represents momentum (velocity multiplied by mass) - represents a constantΔ – in this case, delta represents the uncertainty.

ORBITALS Orbitals indicate probable electron locations

Schrodinger’s equation!This formed the foundation of modern quantum theory

Quantum theory –

Quantum numbers –

PRINCIPLE QUANTUM NUMBER Symbolized by n

Indicates the main energy level occupied by an electron. True or false – these can only be integers

The total number of orbitals in a given shell is equal to n2!

ANGULAR MOMENTUM QUANTUM NUMBER

Different orbital shapes can exist for each principle quantum number.

Represented by l What orbital shapes could you find at

n=3? What about when n=3 and l=2?

MAGNETIC QUANTUM NUMBER Represented by “m”

Ranges from –l to l Indicates the orientation of an orbital

around the nucleus

SPIN QUANTUM NUMBER Represented by “s”

Has only two possible states

A single orbital can hold a maximum of two different electrons

SMARTBOARD DEMONSTRATION

LET’S PRACTICE! CLASS WORK How many electrons could be represented by n=2?

How many electrons could be represented with a principle quantum number of 3 and an angular momentum quantum number of 0?

How many electrons could be represented with a principle quantum number of 1 and a spin number of ½?

How many electrons could be represented with a principle quantum number of 3 and a spin number of ½?

How many electrons could be represented by l=3?