CHM 113 Week 4 Ch. 8 & 9

Electromagnetic radiation

Energy that travels through space as electric and magnetic waves.

Speed of light (c)

3.00 × 10^8 m/s

Wavelength (λ)

Distance between corresponding points on adjacent waves.

Frequency (ν)

Number of waves passing a point per second; units Hz or s^-1.

Relationship between wavelength and frequency

Inversely proportional: wavelength ↑, frequency ↓.

Equation relating wavelength and frequency

c = λν

1 nm in meters

1 nm = 1 × 10^-9 m

Electromagnetic spectrum from lowest energy to highest

Radio → Microwave → Infrared → Visible → Ultraviolet → X-ray → Gamma

Electromagnetic spectrum from longest wavelength to shortest

Radio → Microwave → Infrared → Visible → Ultraviolet → X-ray → Gamma

Electromagnetic spectrum from lowest frequency to highest

Radio → Microwave → Infrared → Visible → Ultraviolet → X-ray → Gamma

Visible light from longest wavelength to shortest

Red → Orange → Yellow → Green → Blue → Violet

Which visible color has the highest energy?

Violet

Which visible color has the longest wavelength?

Red

Planck's constant (h)

6.626 × 10^-34 J·s

Photon energy equation using frequency

E = hν

Photon energy equation using wavelength

E = hc/λ

Energy vs frequency

Directly proportional: frequency ↑, energy ↑.

Energy vs wavelength

Inversely proportional: wavelength ↑, energy ↓.

Shortest wavelength means what energy?

Highest energy

Highest frequency means what energy?

Highest energy

Ground state

Electron is in its lowest possible energy level.

Excited state

Electron has absorbed energy and moved to a higher energy level.

Absorption

Electron moves from lower n to higher n.

Emission

Electron moves from higher n to lower n.

Large energy transition produces

High photon energy, high frequency, short wavelength.

Small energy transition produces

Low photon energy, low frequency, long wavelength.

Rydberg equation

1/λ = RH(1/n1^2 − 1/n2^2)

Rydberg constant (RH)

1.097 × 10^7 m^-1

Rydberg rule for n1 and n2

n1 is the smaller energy level and n2 is the larger energy level.

After calculating 1/λ in the Rydberg equation

Take the reciprocal to get λ.

Rydberg wavelength units before conversion

meters

Maximum electrons in one orbital

2

s subshell

1 orbital; 2 electrons maximum

p subshell

3 orbitals; 6 electrons maximum

d subshell

5 orbitals; 10 electrons maximum

f subshell

7 orbitals; 14 electrons maximum

Principal quantum number n

Energy level or shell.

Allowed values of n

Positive whole numbers: 1, 2, 3, ...

Angular momentum quantum number l

Orbital shape or subshell.

Allowed values of l

0 through n − 1

l = 0

s

l = 1

p

l = 2

d

l = 3

f

Magnetic quantum number ml

Orbital orientation.

Allowed values of ml

−l through +l

Spin quantum number ms

Electron spin.

Allowed values of ms

+1/2 or −1/2

How many quantum numbers describe an orbital?

3: n, l, ml

How many quantum numbers describe an electron?

4: n, l, ml, ms

As n increases, orbital size

Increase

As n increases, orbital energy

Increase

Aufbau principle

Fill lowest-energy orbitals first.

Hund's rule

Place one electron in each equal-energy orbital before pairing.

Pauli exclusion principle

Maximum 2 electrons per orbital and they must have opposite spins.

Paramagnetic

Has one or more unpaired electrons.

Diamagnetic

All electrons are paired.

Electron configuration filling order

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

Electron configuration of phosphorus

1s^2 2s^2 2p^6 3s^2 3p^3

Condensed electron configuration of phosphorus

[Ne]3s^2 3p^3

How to identify an element from an electron configuration

Add all superscripts; total electrons equals atomic number for a neutral atom.

Core electrons

Inner electrons not in the highest occupied principal energy level.

Valence electrons for main-group elements

Electrons in the highest occupied principal energy level.

Cation

An atom that loses electron(s) and becomes positive.

Anion

An atom that gains electron(s) and becomes negative.

Transition-metal cation electron removal rule

Remove electrons from the highest principal energy level first.

Neutral Fe electron configuration

[Ar]4s^2 3d^6

Fe3+ electron configuration

[Ar]3d^5

Atomic radius trend

Increases down a group and to the left across a period.

Ionization energy trend

Increases up a group and to the right across a period.

Effective nuclear charge trend

Increases left to right across a period.

Metallic character trend

Increases down a group and to the left.

Electron affinity general trend

Generally becomes more negative from left to right across a period.

Effective nuclear charge equation

Zeff = Z − S

Z in Zeff equation

Atomic number or number of protons.

S in Zeff equation

Shielding constant; approximately the number of core electrons.

Higher Zeff causes

Smaller atomic radius because electrons are pulled closer.

Approximate Zeff of chlorine

+7

Cation size compared with neutral atom

Cation is smaller.

Anion size compared with neutral atom

Anion is larger.

Why are cations smaller?

Electrons are removed, reducing electron-electron repulsion.

Why are anions larger?

Electrons are added, increasing electron-electron repulsion.

Isoelectronic species

Species with the same number of electrons and the same electron configuration.

Isoelectronic size rule

More protons means smaller radius.

Largest ion in an isoelectronic series

Ion with the fewest protons.

Smallest ion in an isoelectronic series

Ion with the most protons.

Order of S2-, Cl-, K+, Ca2+ from largest to smallest

S2- > Cl- > K+ > Ca2+

Ionization energy

Energy required to remove an electron from a gaseous atom or ion.

General ionization energy trend

Increases left to right and decreases top to bottom.

First ionization energy exception: Group 3A

Group 3A is lower than Group 2A.

First ionization energy exception: Group 6A

Group 6A is lower than Group 5A.

Successive ionization energies

IE1 < IE2 < IE3 < IE4 ...

A huge jump in successive ionization energy means

You have started removing a core electron.

If the huge ionization energy jump occurs after IE3

The atom has 3 valence electrons.

Electron affinity

Energy change when a neutral gaseous atom gains an electron.

More negative electron affinity means

Greater tendency to accept an electron.

Electron affinity exception: Group 2A

Filled s subshell makes electron gain less favorable.

Electron affinity exception: Group 5A

Half-filled p subshell makes electron gain less favorable.

Electron affinity exception: Group 8A

Filled p subshell makes electron gain unfavorable.

Elements in the same group generally share

Related valence-shell electron configurations and similar chemical properties.