Electrons in Atoms

Ch 5 (Chem I/IH)

Light: Electromagnetic Spectrum• Energy can travel in waves. • There are high energy and low energy waves.• The ones we can see are called “the visible

spectrum.” ROY G BIV• Red is the low energy end: violet is the high

energy end.

Properties of Waves1. Wavelength (λ): distance between crests

of a wave. Ex: radio waves = 102 m

Properties, cont.

2. Frequency(ν): number of wave cycles to pass a point per second (wps).• wps = hertz (Hz)= s-1

• Ex: microwaves = 3 x 1010 - 3 x 10 12 Hz

Properties, cont.(skip this slide)

3. Amplitude: wave height from zero to crest

0→

crest→

Electromagnetic Spectrum 1Don’t copy this slide. Instead copy Fig 5.10 from p 139 of your text.

Electromagnetic Spectrum 2 (same instructions!)

Speed of EM Radiation

• All EM radiation travels at the speed of light, c.• c = 2.998 x 108 m/s

Relationship between wavelength & frequency

• c = λν• It is a constant relationship• The product of the 2 variables = the

speed of light• If λ increases, ν decreases• If ν increases, λ decreases

Sample Problem 5.1(p 140)(skip this slide)

• Calculate the wavelength of the yellow light emitted by the lamp shown if the frequency of the radiation is 5.10 x 1014 Hz

Solution (skip this slide)

1. Analyze: Knowns:• ν (frequency)= 5.10 x 10 14 Hz• c = 2.998 x 10 8 m/sUnknown:Wavelength (λ) = ?m

Solution(skip this slide)2. Calculate• Solve the equation c = ν λ for λ• Substitute the known values and

solve.• λ =2.998 x 10 8 m/s 5.10 x 10 14 Hz=

Solution (Cont.)

3. Evaluate: Does the result make sense?

The magnitude of the frequency is much larger than the numerical value of the speed of light, so the answer should be much less than 1. Is it?

Developments in the Atomic Model (skip this slide. Make sure you have this info in your group presentation notes)

• In 1913, we had the Rutherford model of the atom.• electrons thought to occupy the

area outside the nucleus.

Research at the time (1913)

• Scientists knew elements release light when they are excited (by electricity or other energy sources.)• Different elements released different

colors.

Bohr’s Model of the Atom (skip this slide)

Bohr theorized that e-s could only exist at certain distances from the nucleus in energy levels (E.L.’s):

Light, Energy, and Electrons• e-s are arranged in energy levels

(e.l.’s), at different distances from nucleus• Close to nucleus = low energy• Far from nucleus = high energy

Light, Energy, & Electrons, cont.

• e-s in highest occupied level are “valence e-s”• Only so many e-’s can fit in e.l.’s• e-s fill lower e.l.’s before being located in

higher e.l.’s* • Ground state is the lowest energy arrangement

of e-s.

* There are exceptions we will learn later!)

Light, Energy, and Electrons

• e-s can jump to higher energy levels if they absorb energy.• They can’t keep the energy so they

lose it and “fall back” to lower levels.• When they do this, they release the

energy they absorbed in the form of light.

Light, Energy, and Electrons

• (See p 129 of text ChemI/IH) Electron energy levels are like rungs of a ladder. • Ladder– To climb to a higher level, you can’t put your foot

at any level, – you must place it on a rung

• Electron energy levels– e-s must also move to higher or lower e.l.’s in

specific intervals

Niels Bohr (don’t copy this slide)• "The opposite of a correct statement is a

false statement. But the opposite of a profound truth may well be another profound truth." Neils Bohr

• Neils Bohr studied w/Rutherford

Bohr Model of the Atom (don’t copy this slide)

• Interactive Bohr Model

Light, Energy, and Electrons

• Quantum-the amount of energy required to move an electron from one E.L. to another.

Atomic Emission Spectrum (A.E.S)• Each element emits a

color when its excited e-s “fall back.”

• Pass this light thru a prism, it separates into specific lines of color.

• You can identify an element by its emission spectrum! (no 2 elements have the same AES)

Emission Spectra of H, He, Ne (don’t copy this slide)

Atomic Emission Spectrum (cont.)(Look at p 143, but don’t copy slide)

• See p143 of text (ChemI/IH)• H has 4 spectral lines (4 colored lines)• Mercury (Hg) has 11 lines! • Ne has 20+ lines!

Problem: there are more lines than you would expect if there are only a few energy levels.

Hypothesis: There must be many sublevels in an energy level

Quantum Mechanical Model of Atom(don’t copy this slide)

• Bohr’s Model only adequately explained behavior of H• This new model (QMM) explains why

so many emission spectrum lines

QMM, cont. (Don’t copy if you have this in your group presentation notes. Just highlight these items.)

1. Says that waves can behave like particles (Einstein, 1905)

2. Says that particles can behave like waves (de Broglie, 1923)

3. Gives us the allowed energies of e-s (Schrodinger, 1926)

4. & the likelihood of finding e-s at various locations around the nucleus (Schrodinger, 1926)

QMM, cont. (don’t copy this slide)

• Albert Einstein (1905) proposed that light behaves like particles(matter) b/c it has packets of energy called photons• These photons correspond to quanta

of energy

QMM, cont. (don’t copy this slide)

• Louis de Broglie (1924) proposed that particles (matter) can also behave like waves.

• Confirmed in 1927 by Clinton Davisson who bombarded metals with e- beams.– He observed reflection patterns very much like X-

rays (EM radiation)– e-s were behaving like waves!

Use of e- waves (don’t copy this slide)

• Electron microscope magnifies tiny objects b/c e- wavelength much smaller than visible light

snowflake

Heisenburg Uncertainty Principle

• Def: if you want to locate something, you can shine light on it

• When you do this to an electron, the photons send the e- off in an unpredictable direction

• (def):Therefore, you can never know BOTH the position and velocity of an e- at the same time

Electron SublevelsEach electron has an “address,” where it can be

considered to be located in the atom.• Main energy level (principal quantum #)

= “hotel”• Sublevel = “floor”• Orbital = “room” – Regions of space outside the nucleus– All orbitals in a sublevel have the same energy– 2 electrons max can fit in an orbital

Sublevels in Atoms

• See Fig 7.5 on p 235

Main energy level

Types of sublevels

# of orbitals # of electrons

1 s 1

2 s p

13 (4 total)

3 s p d

1 3 5 (9 total)

4-7 s p d f

1 3 5 7 (16 total)

Orbitals• s orbitals are spherical– There is only 1 orbital

• p orbitals are dumbbell shaped– There are 3 orbitals, all with = energy– Each is oriented on either x, y, or z axis– They overlap

• d orbitals have varying shapes– There are 5 orbitals, all with = energy

• f orbitals have varying shapes– There are 7 orbitals, all with = energy

Electron Configurations (don’t have to copy. Info in prior slide)

• Electrons are always arranged in the most stable (lowest energy) way

• This is called“electron configuration” or “ground state”

The Periodic Table & Atomic Structure

• Shape of p. table is based on the order in which sublevels are filled

REGIONS OF THE P. TABLE (see p 244 of book)• s REGION (“block”) - Groups 1 & 2• p REGION (block) - Groups 13-18• d REGION (block)- Groups 3-12 (Transition

Elements)• f REGION (block)- (Inner Transition Elements)

Regions or “Blocks” of the P. Table(don’t need to copy)

Writing e- Configurations for Elements Using the P. Table

1. Always start with Period 1-go from L to R.2. Go to Period 2-from L to R3. Go to Period 3- from L to R4. Continue w/Periods #4-7, L to R, until you arrive at

the element you are writing e- configuration for. • Exception: elements in d block are 1 main E.L lower

than the period where they are located• Exception: elements in f block are 2 main E.L.s

lower than the period where they are located

Correct Order of Sublevels (lowest to highest energy)

• 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

e- configurations

1. Use the P. Table to write the sublevels in increasing order.

2. Add a superscript next to each sublevel that shows how many e-s are in the sublevel

3. Ex: Hydrogen: 1s1

Helium: 1s2

Lithium: 1s22s1

Oxygen: 1s22s22p4

Identifying Valence e-s• Valence e-s are the electrons in the highest

occupied main energy level. (don’t copy. In prior slide)

• Identify them by finding the “biggest big number” in your e- configuration.Ex: Oxygen: 1s22s22p4

• There are 6 valence e-s in the 2nd main energy level (valence level)

Why are d & f block elements’ sublevels out of order?

• When you get to the higher main E.L.’s, the sublevels begin to overlap.

Exceptions: Some Transition Elements (don’t need to copy)

• Titanium - 22 electrons NORMAL• 1s22s22p63s23p64s23d2

• Vanadium - 23 electrons NORMAL• 1s22s22p63s23p64s23d3

• Chromium - 24 electrons EXCEPTION• 1s22s22p63s23p6 4s2 3d4is expected• But this is wrong!!

Chromium is actually… (copy this!)

• 1s22s22p63s23p63d54s1

• 3d54s1Instead of 4s2 3d4 • There is less repulsion (lower

energy) in the 2nd arrangement

4s 3d

Noble Gas Notation

• Short-cut way of showing e- configuration• A Noble Gas is a Group 18 element.1. Identify the noble gas in the period above your

element of interest. Write this symbol in brackets.

2.Write the e- configuration for any additional e-s that your element of interest has, but the noble gas doesn’t have.

Ex: Nitrogen: 1s22s22p5 becomes [He] 2s22p5

Arrow Orbital Diagram-Used to show e- configuration.

SYMBOLS:• A box represents an orbital– Label each box with the sublevel :1s 2s 2p

2p 2p

• An arrow represents an electron– 2 arrows (e-s) in the same orbital face opposite

directions.– Example: oxygen, see above

↑ ↓ ↑ ↓ ↑ ↓ ↑ ↑

Arrow Orbital Diagram-Used to show e- configuration, cont.

INSTRUCTIONS:• Fill electrons from lowest to highest sublevel.• Never place 2 e-s in the same orbital of a

sublevel until you have placed one in each of the orbitals (Hund’s Rule)