1 Atomic Structure and Periodicity AP Chemistry Chapter 7

Atomic Structure Atomic Structure and Periodicityand Periodicity

AP Chemistry

Chapter 7

• Electromagnetic radiation – a form of energy that exhibits wavelike behavior as it travels through space.

• Types include visible light, X rays, ultraviolet light, infrared light, microwaves, and radio waves.

• Electromagnetic spectrum – All the forms of electromagnetic radiation together

7.1 EMR7.1 EMR7.1 EMR7.1 EMR

• Waves have a wavelength – Waves have a wavelength – distance between corresponding distance between corresponding points on adjacent wavespoints on adjacent waves

• Use the Greek letter “lambda”, Use the Greek letter “lambda”, , for wavelength, and units are , for wavelength, and units are length units (m, cm, nm)length units (m, cm, nm)

Electromagnetic Electromagnetic RadiationRadiation

• Waves have a frequency – Waves have a frequency – number of waves that pass a number of waves that pass a given point in a specific timegiven point in a specific time

• Use the Greek letter “nu”, Use the Greek letter “nu”, , for , for frequency, and units are “cycles frequency, and units are “cycles per sec” or Hertz (Hz) per sec” or Hertz (Hz)

• All radiation travels at the same All radiation travels at the same speed of light. speed of light.

• c = 3.00 x 10c = 3.00 x 1088 m/s m/s

• • = c= c• This means that This means that must be in must be in

meters and meters and must be in Hertz (1/s) must be in Hertz (1/s) so that units cancel.so that units cancel.

Electromagnetic Electromagnetic SpectrumSpectrum

Long wavelength Long wavelength small frequency small frequency

Short wavelength Short wavelength high frequency high frequency

increasing increasing frequencyfrequency

increasing increasing wavelengthwavelength

ElectroElectromagneticmagnetic SpectrumSpectrum

Radio waves

Low frequency, Long wavelength1 m - 1km

Microwaves

Infra-Red 0.01 mm

Visible light 400-700 nm

Ultra-Violet 100 nm

X-Rays 1 nm

Gamma Rays

High frequency,

0.01 nmShort wavelength

• Scientific belief around the 1900’s was that there was NO relationship between matter and light

• Light given off by objects that were heated to high temperatures could not be explained.

Problems with Wave Theory of Light

Black Body RadiationBlack Body Radiation

Max Planck

• Stated that objects radiated energy in small packets of energy called quanta

quantum- a specific amount of energy that can be gained or lost by an atom

7.2 Nature of Matter7.2 Nature of Matter

• Energy and frequency are directly related

E=h• E is energy (J)

• h is Planck’s constant

• h = 6.626 x 10-34 J s

Particle Behavior of Light

Thomson (1839)

• First to observe the photoelectric effect

• photoelectric effect - the emission of electrons from a metal surface when exposed to light of a specific energy.

Photoelectric EffectPhotoelectric Effect

1905 Albert Einstein

• stated that EMR could be viewed as a stream of particles “photons”

• photon- a quantum of light

• energy of these photons could be calculated by Planck’s equation

• stated that the photons strike the electrons therefore ejecting them from the metal

Photoelectric EffectPhotoelectric Effect

• The success of Einstein’s work in explaining the photoelectric effect was largely responsible for the acceptance of the particle behavior of light

• Ephoton = h• E = mc2

Dual Wave-Particle Behavior Of Light

• Using Einstein’s and Planck’s equations, de Broglie derived:

• The momentum, mv, is a particle property, whereas is a wave property.

• In one equation de Broglie summarized the concepts of waves and particles as they apply to low mass, high speed objects.

Can matter act as a wave?

• Compare the wavelength for an electron (mass = 9.11 x 10-31 kg) traveling at a speed of 1.0 x 107 m/s with that for a ball (mass = 0.10 kg) traveling at 30 m/s.

Sample ProblemSample Problem

• Energy is a form of matter.

• All matter exhibits both particle and wave properties.

• Large pieces of matter (i.e. baseball) exhibits mostly particle properties.

• Tiny pieces of matter (i.e. photons) exhibits mostly wave properties.

• Pieces of matter somewhere in the middle (i.e. electrons) clearly show both types of properties!

Dual Wave-Particle Behavior Of MatterDual Wave-Particle Behavior Of Matter

• Observed that light was given off when they heated different chemicals in their designed burner

• They passed the light through a prism and saw separate lines instead of a continuous spectrum.

Kirchoff and Robert Bunsen (1854)

Emission spectra- the colors produced by an object when burned or heated.

Absorption spectra- the colors that are not shown, rather absorbed in the spectrum

http://chemistry.beloit.edu/BlueLight/moviepages/ab_em_el.htm

Absorption and Emission SpectraAbsorption and

Emission Spectra

• Only four lines are emitted:– Red, green, blue, violet

• Only certain energies are allowed.

7.3 Hydrogen Spectrum

To understand emission spectrum, we must understand these two terms:

Ground state: the lowest energy state for the electron

Excited state: state where electron has higher energy than ground state.

Why do elements produce these lines?

Atoms are heated, which adds energy. The electron become excited (thus unstable). They want to return to their normal, or ground state. To do so, they give off energy in the form of EMR.

Why do elements produce these lines?

• Balmer: developed a numerical relationship between the wavelength of the lines in the spectrum and the amount of energy

• Lyman: discovered lines produced in the UV range.

• Paschen: discovered lines produced in the IR range

Scientists associated with the H spectrum Scientists associated with the H spectrum

1913 Neils Bohr

-worked with Rutherford to study the H spectrum.

-Bohr’s model is sometimes referred to as the “Planetary model” based upon his postulates.

7.4 The Bohr Model7.4 The Bohr Model

Bohr Model of the Atom

Postulates of Bohr’s model:

1. The single electron of hydrogen can circle the nucleus in fixed paths called orbits or stationary states.

2. The electron can jump to higher orbits when energy is added.

3. The angular momentum of the electron is quantized.

-The electron’s energy can be calculated in the different orbits.

Bohr Model of the AtomHow does this relate to the

Hydrogen spectrum?

• Bohr calculated the energy that the electron would lose as it fell from higher orbits to lower orbits.

• Bohr’s calculations agreed exactly with Lyman, Balmer and Paschen’s observations.

• Calculate the energy required to excite the hydrogen electron from n=1 to n=2.

• Calculate the energy required to completely remove the electron from a hydrogen atom in its ground state.

ninitial = 1 to nfinal = ∞

• Bohr’s model worked very well for the Hydrogen atom.

• Through Bohr’s work, as well as the other scientists mentioned, a very good understanding of the electron within the atom was now in place.

Downfalls to Bohr’s Model

1. Bohr’s model of the atom worked very well for the hydrogen atom and the He+, but failed when applied to multielectron atoms.

2. Bohr’s model could not explain why the electron could not exist between orbits.

Downfalls to Bohr’s Model of the Atom Downfalls to Bohr’s Model of the Atom

• We need a new approach to the atom!

• Big Three: de Broglie, Heisenberg & Schrodinger

• Developed wave mechanics AKA quantum mechanics (7.5)

Now What?Now What?

With respect to atomic particles, we cannot determine exactly

1. the position

2. direction of motion

3. speed

simultaneously.

Heisenberg’s Uncertainty Principle

• Schrödinger proposed an equation that contains both wave and particle terms.

• Solving the equation leads to wave functions .

• The wave function gives the probability distribution of an electron.

• We call wavefunctions orbitalsorbitals.

Schrodinger’s Wave Equation

7.6 Quantum Numbers

• When we solve the Schrödinger equation for the hydrogen atom, we find many wave functions (orbitals) that satisfy it.

• Each orbital is characterized by a

series of numbers called quantum numbers that describe various properties of the orbital.

• Related to the size and energy of the orbital – think energy level

• n has integer values: 1,2,3…

• As n becomes larger, the atom becomes larger and the electron is further from the nucleus.

• A larger n value also corresponds to higher energy because the electron is less tightly bound to the nucleus.

Principal Quantum Number, n

• Related to the shape of the atomic orbitals

• This quantum number depends on the value of n. The values of l begin at 0 and increase to (n - 1).

• Because we use numbers to describe the first quantum number, we usually use letters for l (s for l =0, p for l = 1, d for l =2 and f for l = 3).

• Usually we refer to the s, p, d and f-orbitals.

Angular Momentum Quantum Number, lAngular Momentum Quantum Number, l

• Provides the 3D orientation of the orbital in space

• Value depends on l. The magnetic quantum number has integer values between -l and +l.

Magnetic Quantum number, ml

Quantum NumbersQuantum Numbers

• Each orbital has a unique probability distribution.

• Nodes = areas of zero probability

• To simplify, we think of orbitals in terms of their overall shapes, which becomes larger as n increases.

7.7 Orbital Shapes & Energies

p orbitalsp orbitals

d orbitalsd orbitals

f orbitalsf orbitals

• For hydrogen, energy is determined by value of n

• All orbitals with the same value of n have the same energy – they are degenerate.

Energies of orbitals in Hydrogen

• Developed by Samuel Goudsmit & George Uhlenbeck (University of Leyden in the Netherlands)

• 4th quantum number necessary to account for the details of emission spectra of atoms

• Electron has a magnetic moment with two possible orientations when placed in an external magnetic field.

• Magnetic spin quantum number ms can only have two possible values +½ and -½

7.8 Electron Spin & Pauli Exclusion

Principle

• Wolfgang Pauli developed Pauli exclusion principle

• In a given atom, no two electrons can have the same set of four quantum numbers

• An orbital can hold only 2 electrons, and they must have opposite spins

Principle

Effective Nuclear ChargeEffective Nuclear Charge

•Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge.

•The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons.

•As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases.

•As the distance from the nucleus increases, S increases and Zeff decreases.

7.9 Polyelectronic Atoms

• Dobereiner – triads (groups of 3 elements share similar properties)

• Newlands – octaves (certain properties repeat for every eighth element)

• Meyer & Mendeleev – present form of periodic table

• Mendeleev – considered father of periodic table because he predicted the existence and properties of still unknown elements and left space for them in his periodic table

• Fundamental difference – modern periodic table organized by atomic number not mass

7.10 History of the Periodic Table

• As protons are added one by one to the nucleus to build up the elements, electrons are similarly added

• Electron configurations tells us in which orbitals the electrons for an element are located.

7.11 Aufbau Principle & the Periodic Table

Three rules:

– electrons fill orbitals starting with lowest n and moving upwards;

– no two electrons can fill one orbital with the same spin (Pauli);

– for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule).

Periods 1 - 3Periods 1 - 3

• After Ca the d orbitals begin to fill.

• After the 3d orbitals are full the 4p orbitals being to fill.

• From Ce onwards the 4f orbitals begin to fill.

• Note: La: [Xe]6s25d14f0

• Elements Ce - Lu have the 4f orbitals filled and are called lanthanides.

• Elements Th - Lr have the 5f orbitals filled and are called actinides.

• Most actinides are not found in nature.

Period 4 and Beyond

• The periodic table can be used as a guide for electron configurations.

• The period number is the value of n.

• Groups 1 and 2 have the s-orbital filled.

• Groups 13 - 18 have the p-orbital filled.

• Groups 3 - 12 have the d-orbital filled.

• The lanthanides and actinides have the f-orbital filled.

• Note that the 3d orbital fills after the 4s orbital. Similarly, the 4f orbital fills after the 6s orbital.

Electron Configurations and the Periodic Table

• There is a shorthand way of writing electron configurations

• Write the core electrons corresponding to the filled Noble gas in square brackets.

• Write the valence electrons explicitly.

• Example, P: 1s22s22p63s23p3

• but Ne is 1s22s22p6

• Therefore, P: [Ne]3s23p3.

Noble Gas NotationNoble Gas Notation

Determine the expected electron configurations for each of the following:

• Ba

• Ni2+

• Eu

• Ti+

Practice ProblemPractice Problem

• Many properties of atoms depend on electron configurations and how strongly valence electrons are attracted to the nucleus.

• Coulomb’s Law – strength of the interaction between 2 electrical charges depends on the size of the charges and the distance between them.

• Zeff = Z – S where Z is # protons in nucleus and S is number of core electrons

• Explains differences in sublevel energies but also describes periodic trends.

Effective Nuclear Charge - revisitedEffective Nuclear Charge - revisited

• The effective nuclear charge increases as we move across any row (period) of the periodic table (Z gets larger while S stays the same)

• The effective nuclear charge also increases as we go down a column (group) of the periodic table, but the effect is far less than going across a row.

Effective nuclear charge

Atomic RadiusAtomic RadiusSimple diatomic molecule•The distance between the two nuclei is called the bond distance.•If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom.

• Atomic size varies consistently through the periodic table.

• As we move down a group, the atoms become larger.

• As we move across a period, atoms become smaller.

• There are two factors at work:– principal quantum number, n (down a group)

– the effective nuclear charge, Zeff (across a period)

Atomic RadiusAtomic Radius

• Ionization energy – minimum amount of energy required to remove an electron from the ground state of an isolated gas atom or ion.

• Na(g) Na+(g) + e- First ionization energy

• Na+(g) Na2+(g) + e- Second ionization energy

• The greater ionization energy, the more difficult it is to remove the electron.

Ionization EnergyIonization Energy

• Ionization energy increases for each additional electron removed from an atom.

• There is a sharp increase in ionization energy when a core (non-valence) electron is removed.

Ionization EnergyIonization Energy

• Same factors influence ionization energy – effective nuclear charge & distance of electron from nucleus.

• Increasing effective charge or decreasing distance from nucleus increases attraction between electron & nucleus – more difficult to remove an electron so ionization energy increases. (Both happen when move across row)

• As we move down group, the atomic radius increases (due to larger n) while effective nuclear charge only increases slightly. Attraction between nucleus & electron decreases, so ionization energy decreases.

Ionization Energy Trend

1 Atomic Structure and Periodicity AP Chemistry Chapter 7

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