Chapter 4

Chemistry chapter 4 1

Chapter 4

Arrangement of Electrons in Atoms


Light

• A kind of electromagnetic radiation• A form of energy that exhibits

wavelike behavior as it travels through space


Electromagnetic Spectrum


Frequency

• The number of wave peaks that occur in a given time period

• Represented by the letter f or the greek letter nu (n).

• One wave per second is 1 Hertz (Hz).


Wavelength

• The distance between peaks• Measured in length units such as m

or nm.• Represented by the letter l

(lambda)


Frequency and wavelength


The speed of light

• c is the speed of light (and all electromagnetic waves)• It’s value is 3.0 x 108 m/s

• Frequency and wavelength are related by the equation• c = lf


Photoelectric effect

• The emission of electrons from a metal when light shines on the metal

• Only works when the light is above a certain frequency.


Planck’s Hypothesis

• Quantum – the smallest amount of energy that can be lost or gained by an atom.

• Quanta of radiant energy are called photons.

• E=hf• h= 6.63 x 10-34 J∙s


Wave-Particle Duality

• Introduced by Einstein in 1905• Light exhibits particle-like

properties as well as wave-like ones.


Photon

• Particle of electromagnetic radiation that has zero mass and carries a quantum of energy.


Photoelectric effect

• Einstein explained• In order for an electron to be

knocked loose, it must be struck by a single photon with a high enough energy.

• This requires a high enough frequency.


Ground State

• The lowest energy state of an atom


Excited State

• A state in which an atom has a higher potential energy than its ground state.

• When an atom returns to its ground state, it gives of the extra energy as EM radiation.


Continuous spectrum

• Continuous range of frequencies (colors)

• Expected from classical theory


Line-emission spectrum

• Separated bands of light• Have different frequencies• Produced by passing light through

a thin slit.• Different elements have different

spectra.



Bohr model

• Used quantum theory to explain line spectra.

• Electrons only exist in specific energy states called orbitals.

• Since the change in energy from one state to another is fixed, only certain frequencies are emitted.


Bohr model

• Successful for the hydrogen atom• Needs tweaking for other atoms


De Broglie hypothesis

• 1923• Louis De Broglie’s dissertation• Planck’s quantum theory implied

that light, which had formerly been thought of as a wave, behaves as a particle.

• De Broglie hypothesized that the reverse is true.


Interference

• Occurs when waves overlap• Results in a reduction of energy in

some areas and an increase of energy in others


A two-slit light diffraction-interference pattern

A two-slit electron diffraction-interference pattern

Chemistry chapter 4 23Diffraction pattern from X-rays with l = 7.1 x 10-11 m

Chemistry chapter 4 24Diffraction pattern from 600 eV electrons with l = 5.0 x 10-11 m

Chemistry chapter 4 25Diffraction pattern from .0568 eV neutrons with l = 1.2 x 10-10 m

Chemistry chapter 4 26An electron microscope


An electron micrograph of DNA


Heisenberg

• Pointed out that it is impossible to know both the exact position and the exact momentum of an object at the same time.

• By measuring one, we change the other.


Measuring position

• If we measure the position of an object by hitting it with a photon of energy, the collision with the photon changes its momentum.


Measuring momentum

• If we measure an objects momentum by observing its collision with another object, we have altered its position.


Schrödinger

• Heisenberg treated the electron as a particle.

• Schrödinger treated it as a wave

• Formulated a difficult wave equation with solutions called wave functions


Quantum theory

• Describes mathematically the wave properties of electrons and other very small particles


Probability

• Wave functions only give the probability of finding an electron at a given place

• Electrons don’t travel in neat orbits• “God doesn’t play dice” - Einstein


Orbital

• A 3D region around the nucleus that indicates the probable location of an electron


Discuss

• What is quantum theory?• What radical new idea did de

Broglie introduce?• What is interference?


• Omaha zip codes

• 681 -

http://www.hunteltelephone.com/service_information/images/omahazipspg55.jpg


Four quantum numbers

• Specify the properties of atomic orbitals and the properties of electrons in orbitals

• Each electron in an atom has a unique set of quantum numbers


Principal quantum number, n• Indicates the main energy level

occupied by the electron• Start numbering with 1 at the level

closest to the nucleus.


Electrons in a given level

• The greatest number of electrons that can be in a given level is calculated by the formula 2n2.

• So, in the first level there can be 2 ∙ 12 = 2 electrons.


The angular momentum quantum number, l• Indicates the shape of the orbital• Allowed values: 0, 1, 2, … n – 1


s orbital

• When l = 0• spherical


p orbital

• When l = 1• Dumbbell shaped


d orbital

• When l = 2• More complex


f orbital• When l = 3• Too complex for this class


Sublevels

• Each orbital is designated by the principal quantum number and the orbital letter.

• Examples:• 1s• 2p• 4f


Magnetic quantum number, m• Indicates the orientation of an

orbital around the nucleus• Allowed values: -l to l


s orbital

• Each s sublevel only has one s orbital

• m = 0


p orbitals

• Each p sublevel has 3 different p orbitals

• m = -1, m = 0, m = 1


d orbitals• There are 5 different d orbitals in

each d sublevel• m = -2, m = -1, m = 0, m = 1, m =

2


Spin quantum number

• Has only two values, + ½ and – ½

• Indicate the two fundamental spin states of an electron

• Spin up or spin down


Discussion

• Study table 4-2 on page 104• Answer the section review

questions on that page.


Electron configuration

• The arrangement of electrons in an atom

• Each element has its own unique one.


Ground-state configuration

• Has the lowest energy• All systems in nature tend to be in

lowest energy state.• We can determine the ground-

state configuration with rules


Aufbau principle

• An electron occupies the lowest-energy orbital that can receive it.

• For the order, see page 105 or the chart on the wall.

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


Orbital filling diagram

• This diagram is on your objectives sheet.


Pauli exclusion principle

• No two electrons in the same atom have the same set of four quantum numbers.

• The principal, angular momentum and magnetic quantum numbers specify the orbital.

• The spin number specifies the electron.• Each orbital can hold two electrons.


Hund’s rule

• Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron.

• All single electrons must have the same spin.


Orbital notation

• An unoccupied orbital is represented by a line, ___, with the orbital’s name written under the line.

• An orbital containing an electron is written as ↑ .

• An orbital containing two electrons is written as ↑↓ .


Electron-configuration notation• Uses superscripts instead of

arrows.• The superscript 1 indicates 1

electron in the sublevel.• The superscript 2 indicates 2

electrons in the sublevel.


Hydrogen

• One electron.

H Configuration: 1s1

1s


Helium

• Two electrons.

He 1s2

1s


Lithium

• Three electrons.

Li 1s22s1

1s 2s


Carbon

• Six electrons.

C 1s22s22p2

1s 2s 2p


Oxygen

• Eight electrons.

O 1s22s22p4

1s 2s 2p


Discuss

• What is an atom’s electron configuration?

• What three principles guide the electron configuration of an atom?


Discuss

• The electron configuration of fluorine is 1s22s22p5. How many electrons does fluorine have? What is its atomic number?

• Write the electron configuration of sulfur, which has an atomic number of 16.


Highest occupied level

• The highest n level that contains an electron.


Inner-shell electrons

• Aren’t in the highest occupied level


octet

• The 8 electrons (or electron spaces) in the highest occupied level.

• The s and p orbitals.• If they are all occupied, then the

octet is full.


Noble Gases

• Group 18 elements• Have a full octet


Noble Gas notation

• Shorthand for electron-configuration.

• Start with the noble gas from the period above, then add on.

• Example: magnesium• 1s22s22p63s2

• [Ne]3s2


You try

• Write the electron configuration and noble gas notations for titanium

• 1s22s22p63s23p63d24s2

• [Ar] 3d24s2

• Notice that we write the 3d sublevel before the 4s, even though the 4s fills first


You try

• Write the electron configuration and noble gas notations for copper

• 1s22s22p63s23p63d104s1

• [Ar]3d104s1


Exceptions to the rules

• An electron will leave the 4s orbital to create a half-filled or a filled 3d sublevel.

• This configuration is more stable.• Chromium [Ar]3d54s1


Exceptions to the rules

• Sometimes an electron will leave the 5s sublevel to go to the 4d sublevel. This makes the atom more stable.

• There isn’t a pattern like the 4s to 3d switch.


Discuss

• Write the noble-gas notation for aluminum. • [Ne]3s23p1

• How many outer-shell electrons does an atom of aluminum have? • 3

• How many unpaired electrons does an atom of aluminum contain?• 1


Extra Credit• Write a paragraph explaining how

you can use the periodic table to determine the order in which orbitals are filled. Your explanation should include references to group and period numbers. Worth up to 10 extra credit points.

Date post:	30-Dec-2015
Category:	Documents
Upload:	brynn-landry
View:	30 times
Download:	0 times

Chapter 4

Documents