ATOMIC STRUCTURE AND PERIODICITY

By: Ms. Buroker

ELECTROMAGNETIC RADIATION

Energy that moves through space in a wavelike manner.

Electromagnetic radiation is self-propagating (i.e. it doesn’t require a medium to travel through. It can travel through the vacuum of space.)

ELECTROMAGNETIC RADIATION

ELECTROMAGNETIC SPECTRUM

EXAMPLE PROBLEMS

1.) Which color in the visible spectrum has the highest frequency? Which has the lowest frequency?

2.) Is the frequency of the radiation used in a microwave oven higher or lower than that from a FM radio station broadcasting at 91.7MHz?

3.) Is the wavelength of x-rays longer or shorter than that of ultraviolet light?

EXAMPLE PROBLEM

The brilliant red colors seen in fireworks are due to the emission of light with wavelengths around 650nm when strontium salts such as Sr(NO3)2 and SrCO3 are heated. Calculate the frequency of red light of wavelength

6.50 x 102nm.

Traditional Theories:Intensity of radiation

should increasecontinuously with

decreasing wavelength

A QUANDARY

Are matter and energy different?

Matter was thought to consist of particles, whereas energy in the form of light (electromagnetic radiation) was described as a wave. Particles were things that had mass and whose position in space could be specified. Waves were described as massless and delocalized; that is, their position in space could not be specified, It also was assumed that there was no intermingling of the matter and light.

MAX PLANK (1858- 1947)

Lead to the discovery of a quantum … which is like a packet of energy released when there is a transfer.

Vibrations in atoms are quantized (i.e. Only certain vibrations with certain frequencies are allowed)

THE PHOTOELECTRIC EFFECT

Assumption:Energy is carried on the amplitude of a wave (as in classical waves).

Prediction: Since the amplitude of an EM wave correlates with the brightness of the light, light of high enough intensity irradiating a metal for a long enough period of time should be able to eject electrons from the surface of a metal.Reality: It was the frequency of the light that determined whether or not electrons were ejected regardless of the time of irradiation. The brightness (amplitude) only determined how many were ejected per unit time once ejection started occurring.

EXAMPLE PROBLEM

The blue color in fireworks is often achieved by heating copper(I) chloride (CuCl) to about 1200○C. Then the compound emits blue light having a wavelength of 450nm. What is the increment of energy (the quantum) that is emitted at 4.50 x 102nm by CuCl?

EINSTEIN AND THE PHOTON

Einstein (1905) incorporated Planck’s equation with the idea that light had (mass-less) particle properties (photon). These photons are “packets” of energy where E depends on the frequency of the photon (E = hν).

EINSTEIN AND THE PHOTON He proposed that electromagnetic

radiation is itself is quantized and can be viewed as a stream of particles called photons.

Ephoton = h = n hc

l

EXAMPLE PROBLEM

A 60W, monochromatic laser beam gives off photons of wavelength 650nm. How long does it take for 2.5moles of these photons to be given off?

Answer:E = hc/λ = (6.626x10-34Js)(3.0x108m/s)/6.50x10-7m= 3.06x10-19J/photon

(3.06x10-19J/photon)(6.022x1023photons/mol)(2.5mol) = 4.60x105J

4.60x105J / (60J/s) = 7673.4s /(3600s/1hr) = 2.13h

EXAMPLE PROBLEMCompare the energy of a mole of photons of orange light (625nm)with the energy of a mole of photons of microwave radiation having a frequency of 2.45GHz (1GHz = 109s-1). Which has the greater energy? By what factor is one greater than the other?

Answer:E = (6.626x10-34Js)(3.00x108m/s)/(6.25x10-7m) = 3.18x10-19J3.18x10-19J(6.022x1023) = 1.92x105J/mol

E = (6.626x10-34Js)(2.45x109Hz) = 1.623x10-24J1.623x10-24(6.022x1023) = 9.78x10-1J/mol

1.9x105 / 9.7x10-1 = 1.96x105

Atomic Line Spectra

Atomic Emission Spectra

ATOMIC EMISSION SPECTRA

BOHR MODEL OF THE ATOM

First connection between line spectra and quantum ideas ofPlanck and Einstein.

Bohr Model: Electrons orbit the nucleus of the atom like planetsgoing around the sun.

Only certain stable orbits are allowed (quantized) to keep theelectron (a charged, accelerating particle) from crashing into thenucleus.

n is an integer equal to or greater than 1 andRhc = 2.179x10-18J/atom or 1312kJ/mol

Principle quantum number

Note: Potential energy = 0 at infinity

ELECTRON TRANSITION DIAGRAM

1) Electron in ground state2) Electron jumps to “excited state” from absorption of outside energy (Energy absorbed = positive)3) Electron transitions back down giving off photon that is equal in energy to the transition down (energy emitted = negative)

Lyman (UV),Balmer(Visible), andPaschen (IR),series of thehydrogenatom.

EXAMPLE PROBLEM

Calculate the energy of the n=3 state of the H atom in a) joules peratom and b) kilojoules per mole.Rhc = 2.179x10-18J/atom or 1312kJ/mol

Answer:E = -Rhc/n2 = -2.179x10-18J/atom / 32 = -2.421x10-19JE = -1312kJ/mol / 32 = -145.8kJ/mol

QUANTUM NUMBERS

Each electron in an atom has a unique set of 4 quantum numbers which describe it.

Principal quantum number Angular momentum quantum number Magnetic quantum number Spin quantum number

Pauli Exclusion Principle

No two electrons in an atom can have the same four quantum numbers.

Wolfgang Pauli

Principal Quantum Number

Generally symbolized by n, it denotes the shell (energy level) in which the electron is located.

Number of electrons that can fit in a shell:

2n2

Angular Momentum Quantum Number

The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.

Magnetic Quantum Number

The magnetic quantum number, generally symbolized by ml, denotes the orientation of the electron’s orbital with respect to the three axes in space.

Assigning the Numbers

The three quantum numbers (n, l, and ml) are integers. The principal quantum number (n) cannot be zero 1, 2, 3, etc. The angular momentum quantum number (l ) can be any integer between 0 and n - 1. For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (ml) can be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1, +2.

Spin Quantum Number

Spin quantum number (ms) denotes the behavior (direction of spin) of an electron within a magnetic field.

Possibilities for electron spin: 1

2

1

2

An orbital is a region within an atom where thereis a probability of finding an electron. This is a probability

diagram for the s orbital in the first energy level…

Orbital shapes are defined as the surface that contains 90% of the total electron probability.

Schrodinger Wave Equation

22

2 2

8dh EV

m dx

Equation for probability of a single electron being found along a single axis (x-axis)Erwin Schrodinger

Heisenberg Uncertainty Principle

You can find out where the electron is, but not where it is going.

OR…

You can find out where the electron is going, but not where it is!

“One cannot simultaneously determine both the position and momentum of an electron.”

Werner Heisenberg

Orbitals of the same shape (s, for instance) grow larger as n increases

Nodes are regions of low probability within an orbital.

The s orbital has a spherical shape centered around the origin of the three axes in space.

s orbital shape

P orbital shape

Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of:

…and a “dumbell with a donut”!

“double dumbells”

d shaped orbitals

Shape of f orbitals

Orbital filling table

Irregular confirmations of Cr and Cu

Chromium steals a 4s electron to half fill its 3d sublevel

Copper steals a 4s electron to FILL its 3d sublevel

Half of the distance between nuclei in covalently bonded diatomic molecule "covalent atomic radii"

Periodic Trends in Atomic Radius

Radius decreases across a period Increased effective nuclear charge due to decreased shielding

Radius increases down a group Addition of principal quantum levels

Determination of Atomic Radius

Table of Atomic Radii

Tends to increase across a period

Electrons in the same quantum level do not shield as effectively as electrons in inner levels

Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove

Tends to decrease down a group

Outer electrons are farther from the nucleus

IONIZATION ENERGYIonization Energy: the energy required to remove an electron from an atom

Mg + 738 kJ Mg+ + e-

Mg+ + 1451 kJ Mg2+ + e-

Mg2+ + 7733 kJ Mg3+ + e-

Ionization of Magnesium

Table of 1st Ionization Energies

ELECTRON AFFINITY

Electron Affinity is the energy change associated with the addition of an electron. So … think the opposite of ionization energy.

1.) Electron affinity tends to increase across a period.

2.) Affinity tends to decrease as you go down in a group.Electrons farther from the nucleus experience less nuclear attraction

Some irregularities due to repulsive forces in the relatively small p orbitals

Table of Electron Affinities

ELECTRONEGATIVITY

Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons.

Trend ….1.) Electronegativity tends to increase as you go across a period.

2.) Electronegativity tends to decrease as you go down a group or remain the same.

FACTORS DETERMINING ELECTRONEGATIVITY

When you subtract the electronegativity values of two atoms bound together … you use the value to determine what kind of bond you have.

Non-polar covalent= 0-0.3

Polar Covalent Bonds= 0.3- 1.7

Ionic Bonds= 1.7- 3.3

Periodic Table of Electronegativities

Cations Positively charged ions

  Smaller than the corresponding atomAnions

  Negatively charged ions   Larger than the corresponding atom

Ionic Radii

Table of Ion Sizes

Summary of Periodic Trends