Honors Chem Chapter 4The Tiny but Mighty Electron

Electron: What do we already know?

• It has a negative charge (Thomson) • It is small in mass: approximately 0.0005 amu

or 9.11x10-28 grams (Milikan)• Previously, it was thought of traveling in orbits

around the nucleus (Bohr) • Currently, it is thought of as moving rapidly

outside of the nucleus in orbitals

Orbital

• A probable space outside of the nucleus where an electron is likely to be found

• Electrons are organized in their orbitals according to relative energy

• Low energy: Closest to the nucleus• High energy: Farthest from the nucleus

What do orbitals look like?

“In a science that we cannot see, a lot is left to the imagination…” –Correspondent from NOVA NOW

Level 1

• In energy level 1, there is one orbital shape. Its called an s orbital and looks like a sphere.

Level 2

• In energy level 2, there are two orbital shapes. First, an s orbital and then a p orbital. The p orbital is shaped like dumbbell. There are 3 of these shapes. Each one is a subshell.

Level 3 • In energy level 3, there are three orbital

shapes. One s, Three p’s, and Five d’s.

Level 4

• In energy level 4, there are four orbital shapes. One s, Three p’s, Five d’s, and Seven f’s

Heisenberg Uncertainty Principle

• We cannot know the exact location and speed of an electron at the exact time. We can know one or the other precisely, but never both at the same time.

• WHY????

Finding An Electron: Quantum Numbers

• We can assign an electron a series of 4 numbers to “guesstimate” where it can be found in an atom.

• 4 parts: n, l, m, s

Quantum Numbers: Principle Energy Level

• Represented by the letter “n”

• Whole number: 1, 2, 3, 4, etc…• Represents the size of the orbital. The bigger

the number, the larger the orbital and also the further it is from the nucleus

Quantum Numbers: Angular

• Represented by the letter “l” • Whole number: 0, 1, 2, and 3• Represents the shape of the orbital – For an “s” shape: l = 0– For a “p” shape: l = 1– For a “d” shape: l = 2– For an “f” shape: l = 3

Quantum Numbers: Magnetic

• Represented by the letter “m” • Can be a range of numbers from the negative

integer to the positive integer• Represents the orientation (or number of

subshells) – For “s” shape: m = 0 (only one orientation)– For “p” shape: m= -1, 0, 1 (3 orientations)– For “d” shape: m = -2, -1, 0, 1, 2 ( 5 orientations)– For “f” shape: m = -3, -2, -1, 0, 1, 2, 3 (7 orientations)

Quantum Numbers: Spin

• Represented by the letter “s”

• Can be either +1/2 or -1/2 • Represents the spin direction of the electron

Why can’t electrons stay in one place?

• Electrons are “hit” by ambient radiation sources and when they are given more energy they are “promoted” to a higher energy level

• This is borrowed energy so what goes up must come down.

• This release of energy comes in the form of light and/or heat!

Dual Nature: Particles and Waves

• Light is packets or bundles of energy, known as photons, that travel in waves (Einstein)

Electrons and Energy

• Photoelectric effect: electrons are emitted from samples of matter when they are exposed to radiation energy—LIGHT!

• Electrons have energy, absorb energy, and release energy

• Electrons and Light exhibit properties of both particles and waves

Why can’t electrons stay in one place?

• Electrons are “hit” by ambient radiation sources and when they are given more energy they are “promoted” to a higher energy level

• This is borrowed energy so what goes up must come down.

• This release of energy comes in the form of light and/or heat!

Sources of Energy

• Light/Heat • Electricity • Chemical Reaction• Nuclear Reaction

Ground State

• The ground state of an electron is its lowest-energy state.

Excited State

• An excited state is when an electron has been promoted to a higher energy level.

• As the electron returns to its ground state, it releases the specific gained energy in the form of light

Bright Line Spectrum

Bright Line Spectrum

Understanding Waves: Wavelength

• Wave = repetitive transfer of energy

• Wavelength (λ) = distance over which the wave’s shape repeats

• Generally measured in nanometers (1 x 109 nm = 1 m)

Understanding Waves: Frequency

• Frequency (ν) = number of occurrences of a repeating event per unit time

• Measured in Hertz (Hz) OR “waves per second” (1/s = s-1)

Frequency and Wavelength Related

• Frequency and Wavelength are inversely related.

• One variable increases while the other decreases.

The Light Equation

•c = λν• The speed of light (c) is equal to 3.0 x108 m/s• The wavelength (λ) is in meters • The frequency (ν) is in Hertz

Wave Practice

An orange light has a wavelength of 492nm. What is the frequency of this light?

Energy and Frequency

• High frequency (short wavelength) is a high energy wave.

• Frequency and energy are directly related • Both variables increase together

The Electromagnetic Spectrum

The Visible Spectrum• ROY G BIV: Red, Orange, Yellow, Green, Blue,

Indigo, and Violet

Increasing energy

Increasing frequency

decreasing wavelength

Bright Line Spectrum

• Light is a combination of multiple colors• Each atom has its own unique pattern of

electrons that absorb energy differently • No two atoms have the same bright line

spectrum: used for identification• Each line in the spectrum represents electrons

releasing a specific amount of visible energy—translating to a specific color.