1 electronic configurations and the periodic table chapter 7 robert w. bunsen 1811-1899. established...

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Electronic Configurations and Electronic Configurations and the Periodic Tablethe Periodic Table

Chapter 7Chapter 7

Robert W. Bunsen 1811-1899. Established the science of spectroscopy.

Niels H. Bohr 1885-1962.* Related Periodic Table to atomic atomic structure and spectral behavior.

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The Wave Nature of LightThe Wave Nature of Light

•Modern atomic theory arose out of studies of the interaction of radiation with matter.

•Electromagnetic radiation moves through a vacuum with a speed of 2.99792458 108 m/s.•(We will usually round off to 3 x 108 m/s)

•Electromagnetic waves have characteristic wavelengths ( and frequencies () .

•Example: visible radiation has wavelengths between • 400 nm (violet) and 750 nm (red).

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Some examples of EM radiation:

visible lightultraviolet radiation (UV)infrared radiation (IR)radio wavesmicrowavesx-rays?sound waves?

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The frequency, , of a wave is the number of cycles which pass a point in one second.

The speed of a wave, c, is given by its frequency multiplied by its wavelength:c = where c = 3 x 108 m.s-1 (the “speed of light”)

The wavelength, λ, is the distance between crests in a wave

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Short wavelength High frequency as λ ν

Long wavelengthSmall frequency as λ ν

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Quantized Energy and PhotonsQuantized Energy and Photons

Some problems affecting science in 1900:

Blackbody radiationBlackbody radiationPhotoelectric effectPhotoelectric effectLine spectraLine spectraSpiraling of electrons into nucleusSpiraling of electrons into nucleus

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Quantized Energy and PhotonsQuantized Energy and PhotonsPlanckPlanck: energy can only be absorbed or released from atoms in certain amounts called quantaquanta.

This energy is proportional to the frequency ,, of the radiation. The proportionality constant, h, is called Planck’s constant and is equal to: 6.626 10-34 J.s

E = h

Matter can absorb or emit energy only in

quantum units which are multiples of h :

ΔE = nh where n=integer

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The Photoelectric EffectThe Photoelectric Effect•The photoelectric effect provides evidence for the particle nature of light -- “quantization”.•If light shines on the surface of a metal, there is a point at which electrons are ejected from the metal.

•The electrons will only be ejected once the threshold frequency, υ, is reached.

•Below the threshold frequency, no electrons are ejected.

•This is consistent with the Planck statement that light energy depends only on frequency.

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The Photoelectric EffectThe Photoelectric Effect

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Energy ofejectedelectrons

thresholdfrequency

o frequency,

The Photoelectric EffectThe Photoelectric Effect

•Einstein assumed that light traveled in energy packets called photons.

•The energy of one photon, E = h.

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1. Calculate the wavelength and energy of a photon with frequency of 2.3 x 1014 Hz

2. Calculate the frequency and energy of a photon with a wavelength of 420 nm

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Lothar Meyer 1830-1895. Codiscoverer of Periodic Table.

Meyer noticed a periodic trend in atomic sizes andsuggested there was an inner structure to the atom,although it was too early to guess what the natureof this structure was.

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Bohr’s Model of the Hydrogen Atom Bohr’s Model of the Hydrogen Atom

Bohr’s ModelBohr’s Model•Rutherford assumed the electrons orbited the nucleus analogous to planets around the sun.•However, a charged particle moving in a circular path should lose energy.•This means that the atom should be unstable according to Rutherford’s theory.•Bohr noted the line spectra of certain elements and assumed the electrons were confined to specific energy states. These were called orbits.

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Na

H

Bohr’s Model of the Hydrogen AtomBohr’s Model of the Hydrogen AtomLineLine Spectra SpectraColors from excited gases arise because electrons move between energy states in the atom. These arecalled line spectra.

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gas discharge tube

prism

Line Spectra

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Absorption Emission

E1

E2

E3

E4

Bohr’s Model of the Hydrogen AtomBohr’s Model of the Hydrogen Atom

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E1

E2

E3

Absorption

Emission

N

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Bohr’s Model of the Hydrogen Atom Bohr’s Model of the Hydrogen Atom The first orbit in the Bohr model has energy E1 and is assigned a quantum number, n=1. The next orbit hasenergy E2 and quantum number, n=2.

hEEE if BUT.......

Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta (h).

The amount of energy absorbed or emitted on movement between states is given by

The Bohr model doesn’t work!

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The Wave Behavior of MatterThe Wave Behavior of MatterKnowing that light has a particle nature, it seems reasonable to ask if matter has a wave nature.Using Einstein’s and Planck’s equations, de Broglie supposed:

mvh

The momentum, mv, is a particle property, where as is a wave property.

In one equation de Broglie summarized the concepts of waves and particles, with noticeable effects if the objects are small.

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Quantum Mechanics and Atomic OrbitalsQuantum Mechanics and Atomic Orbitals

•Solving the equation leads to the wave function Ψ These are later referred to as orbitals.•Wave functions (orbitals, Ψ ) are mathematical quantities.•The square of the wave function, Ψ 2 gives the probability of finding the electron in some region of space.•Orbitals have shapes and spatial orientations.• The old Newtonian, classical theory is deterministic.•The new quantum wave theory is probabilistic.

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

tizyxV

zyxm

hh

),,(2 2

2

2

2

2

22 (impressyour friends)

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Quantum Mechanics and Atomic Quantum Mechanics and Atomic OrbitalsOrbitals

Ψ2 = probability of finding electronin box

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Orbitals and Quantum NumbersOrbitals and Quantum NumbersIf we solve the Schrödinger equation, we get wave

functions (orbitals), energies and quantum numbers. Schrödinger’s equation requires 3 quantum numbers:

Principal Quantum Number, n. This is the same as

Bohr’s n. As n becomes larger, the atom becomes

larger and the electron is further from the nucleus.

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Quantum Mechanics and Atomic Quantum Mechanics and Atomic OrbitalsOrbitalsOrbitals and Quantum NumbersOrbitals and Quantum Numbers

Azimuthal Quantum Number, l. This quantum number

depends on the value of n. l=0,1,2,....n-1

We usually use letters for l (s, p, d and f for l = 0, 1, 2, and 3).

l has to do with orbital shape.

Magnetic Quantum Number, ml. This quantum number depends on l. The magnetic quantum number has integral values between -l and +l. Magnetic quantum numbers give the 3D orientation of each orbital.

l = 0 1 2 3 4 s p d f g

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n l ml

1 0 (s) 0 1s 2 0 (s) 0 2s 1 (p) 1,0,-1 2p 3 0 (s) 0 3s 1 (p) 1,0,-1 3p 2 (d) 2,1,0,-1,-2 3d 4 0 (s) 0 4s 1 (p) 1,0,-1 4p 2 (d) 2,1,0,-1,-2 4d 3 (f) 3,2,1,0,-1,-2,-3 4f

ShellK

L

M

N

SubShell

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Representation of OrbitalsRepresentation of Orbitals

As n increases, the s-orbitals get larger.

All s-orbitals are spherical in shape.

The The ss Orbitals Orbitals

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The The pp Orbitals Orbitals

There are three p-orbitals px, py, and pz in each l=1 subshell.

These correspond to allowed values of ml

of -1, 0, and +1.)

The orbitals are dumbbell shaped.

As n increases, the p-orbitals get larger.

All p-orbitals have a node at the nucleus.

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The The pp Orbitals Orbitals

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The The dd and and ff Orbitals Orbitals

There are 5 d-orbitals in l=2 subshells. There are 7 f-orbitals in l=3 subshells

Four of the d-orbitals have four lobes each.One d-orbital has two lobes and a collar.

(We will not be concerned with the f-orbitals)

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Representation of OrbitalsRepresentation of OrbitalsThe The dd Orbitals Orbitals

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Orbitals and Quantum NumbersOrbitals and Quantum Numbers

Orbitals can be ranked in terms of increasing energy to yield an Aufbau diagram.Note that the Aufbau diagram on next slide is for a single electron system.As n increases, note that the spacing between energy levels becomes smaller.

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Quantum Mechanics and Atomic OrbitalsQuantum Mechanics and Atomic Orbitals

Orbitals and Quantum Orbitals and Quantum NumbersNumbersEach square is an orbital.There are n2 orbitals in each shell.

Note that orbitals with same value of n have the same energy

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Orbitals in Orbitals in Many ElectronMany Electron Atoms (everything above Atoms (everything above hydrogen)hydrogen) Energies of OrbitalsEnergies of Orbitals

Note: 2p above 2s 3p above 3s 3d above 3pBUT: 4s below 3d

Rule:for given value of nenergy of orbital increases with increasing value of lExample: for n=3l = 0 1 2

s p d< <Increasing energy

All due to screening(next chapter).

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Orbitals in Orbitals in Many ElectronMany Electron Atoms Atoms

1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f 5g6s 6p

Mnemonic Device to remember Orbital Sequence

1s 2s 2p3s 3p4s 3d4p5s 4d5p6s

When filled with electrons, it’s

1s22s22p63s23p64s23d104p65s2...

(superscripts are numbers of electrons in orbital series)

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Orbitals in Orbitals in Many ElectronMany Electron Atoms Atoms

Electron Spin and the Pauli Exclusion PrincipleElectron Spin and the Pauli Exclusion Principle

Studies in the 1920s demonstrated that the electron hasa property called spin.

This introduced a new (and fourth) quantum number,the electron spin (or just spin) quantum number, ms

There are two values for ms: +1/2 and -1/2

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Orbitals in Orbitals in Many ElectronMany Electron Atoms AtomsElectron Spin and the Pauli Exclusion PrincipleElectron Spin and the Pauli Exclusion Principle

ms=+1/2 ms =-1/2

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Orbitals in Orbitals in Many ElectronMany Electron Atoms AtomsElectron Spin and the Pauli Exclusion PrincipleElectron Spin and the Pauli Exclusion Principle

Pauli’s Exclusion Principle:Pauli’s Exclusion Principle: no two electrons can have

the same set of 4 quantum numbers.

Therefore, two electrons in the same orbital must have

opposite (+1/2 and -1/2) spins.

Can use arrows ( and ) to represent +1/2 and -1/2spins, respectively.For example, in the 1s orbital, n=1, l=0 and ml=0.One electron can be placed (+1/2) and the other at (-1/2). But, that’s it! No more electrons in the1s orbital. Why?

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Electron ConfigurationsElectron Configurations

Periods 1, 2, and 3Periods 1, 2, and 3

Electron configurations tell us in which orbitals the electrons for an element are located.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 (orbitals with same energy), electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule).

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1s

2s

2p Hund’s Rule

3s

3p Hund’s Rule

4s

3d

(Hund’s Rule)

1s2 2s2 2p6 3s2 3p6 4s2 3d10

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Electron ConfigurationsElectron Configurations

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Electron Configurations and the Periodic Electron Configurations and the Periodic TableTable

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Electron Configurations and the Periodic Electron Configurations and the Periodic TableTable

There is a shorthand way of writing electron configurationsWrite the core electrons corresponding to the preceding filled Noble gas in square brackets.Write the valence electrons explicitly.Example, P: 1s22s22p63s23p3

but Ne is : 1s22s22p6

Therefore, P: [Ne]3s23p3.

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Development of the Periodic TableDevelopment of the Periodic Table

•How do we organize elements in a meaningful way that will allow us to make predictions about undiscovered elements?•Arrange elements to reflect the trends in chemical and physical properties.•First attempt (Mendeleev and Meyer) arranged the elements in order of increasing atomic weight.•Modern periodic table: arrange elements in order of increasing atomic number.•Elements in same column (group) have same or similar properties.•Properties are periodic.

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+e

ee

e-

this electron

is partially shielded from this nucleus

by all the othersurrounding electrons in atom

Screening occurs when you have more than one electron to consider.

Effective Nuclear Charge and ScreeningEffective Nuclear Charge and Screening

This electron experiences an effective nuclear charge, Zeff

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Effective nuclear charge, Zeff, and screeningOuter (valence) electrons experience an effective (not full) nuclear charge because of screening (or shielding).

Effective nuclear charge

Zeff = Z – no. of core electrons

(Z = actual nuclear charge)

e.g., for Na (1s22s22p63s1) Zeff = 11-10 = +1 Mg (1s22s22p63s2) Zeff = 12-10 = +2

Al (1s22s22p63s23p1) Zeff = 13-10 = +3

Note that Zeff increases across a row

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The Periodic Table is used to organize the 114 elements in a meaningful way

The Periodic TableThe Periodic Table

As a consequence of this organization, there are periodic properties associated with the periodic table.

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Development of the Periodic TableDevelopment of the Periodic Table

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Don’t forget to review historical questionsregarding the Periodic Table that are found

in the Problem Bank:

Dalton, Mendeleev, Rutherford, Lavoisier, Bunsen,Berzelius, Seaborg, Curie, Davy, Bohr, Moseley, Soddy, the Ancients, all discussed in Chapters 2 and 7.

Review groups (families) of elements in Chapter 2:alkali metals, alkaline earths, rare earths (lanthanides),actinides, transuraniums, halogens, noble (inert)gases.

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In Au, internuclear distance = 2.88 Å

Array of Au atoms:

Therefore, atomic radius = 2.88 = 1.44 Å 2

Sizes of Atoms

Covalent (atomic) radii

(1 Å = 10-10 m)

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Sizes of AtomsSizes of Atoms•As a consequence of the ordering in the periodic table, properties of elements vary periodically.•Atomic size varies consistently through the periodic table.•As we move down a group, the atoms become larger.

Each row (period)

is a new shell.

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•As we move across a period, atoms become smaller.Sizes of AtomsSizes of Atoms

•As we move across the periodic table, the number of core electrons remains constant. But, the nuclear charge increases. •Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease.

•principal quantum number, n, and•the effective nuclear charge, Zeff.

There are two factors at work:

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Sizes of AtomsSizes of Atoms

In general, atomic size increases this way in the table.

inc

incinc

Atomic Size

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AtomicRadii

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Problem: Estimate the As-I bond length

Radii:As = 1.19 A I = 1.33 A

1.33 1.19

I As

Bond Length=1.33 + 1.19= 2.52 A

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Ionization EnergyIonization Energy

• The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom:

Na(g) Na+(g) + e-.•The second ionization energy, I2, is the energy required to remove an electron from a gaseous ion:

Na+(g) Na2+(g) + e-.

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

There is a sharp increase in ionization energy when a core electron is removed.

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Ionization Energy

element at end of row

element at beginning of row

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Ionization EnergyIonization EnergyPeriodic Trends in Ionization EnergyPeriodic Trends in Ionization Energy•Ionization energy decreases down a group.

This means that the outermost electron is more readily removed as we go down a group. Why?

Group IALi (520 kJ/mol)Na (496)K (419)Rb (403)Cs (376)

As the atom gets bigger, it becomes easier to remove an electron from the most spatially extended orbital.

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•Ionization energy generally increases across a • period.•As we move across a period, Zeff increases. Therefore, it becomes more difficult to remove an electron.

Ionization EnergyIonization Energy

Period 2: Li Be B C N O F NeIE: 520 899 801 1086 1402 1314 1681 2081

(all in kJ/mol)

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IonizationEnergy

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Ionization EnergyIonization EnergyIn general, ionization energy increases this way in the table.

inc

incinc

Ionization Energy

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Electron AffinitiesElectron Affinities•Electron affinity is the opposite of ionization energy.•Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion:

Cl(g) + e- Cl-(g)•Electron affinities are usually exothermic (as in the above example)

inc

inc

inc

Electron Affinity

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Electron AffinitiesElectron AffinitiesThe added electron in Cl is placed in the 3p orbital to form the stable 3p6 electron configuration.

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Ionic RadiiIonic RadiiJust as atom size is periodic, ion size is also periodic.

Cations (+ ions)•To form cations, outermost or valence electrons are removed. •The effective nuclear charge has increased.•Therefore, the cation is smaller than the parent.

Anions (- ions) To form anions, electrons are added to the

outermost orbital. •The nuclear charge has remained the same, but the number of screening electrons is increased.•Therefore, anions are larger than their parents.

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Ionic RadiiIonic Radii

All the members of an isoelectronic series have the same number of electrons.

For ions of the same charge, ion size increases down a group.

As nuclear charge increases in an isoelectronic series the ions become smaller:

O2- F- Na+ Mg2+ Al3+ (all have 10 e’s)

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Ionic RadiiIonic RadiiIsoelectronic Species

O2-

Z=8#e-=10

F-

Z=9#e-=10

Ne

Z=10#e-=10

Na+

Z=11#e-=10

Mg2+

Z=12#e-=10

Attraction to nucleus increases

Size Decreases

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Ionic and Atomic Radii

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Ionic RadiiIonic Radii

Which of the following species is the largest?

B B+ Al Al+

Which of the following species is the smallest?

P P- S S-

Which of the following species is the largest?

N- P+ P- P

Which of the following species is the smallest?

Rb+ K+ Cl- Ar

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Metals, Nonmetals, and MetalloidsMetals, Nonmetals, and Metalloids

MetalsMetals•Metallic character refers to the properties of metals (shiny or lustrous, malleable and ductile)• Metals react with acids (H+)

•Metallic character increases down a group.

•Metallic character decreases across a period.

inc

•Metals have low ionization energies.

dec

incinc

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•When metals are oxidized they tend to form characteristics cations.•All group 1A metals form M+ ions.•All group 2A metals form M2+ ions.•Most transition metals have variable charges.

MetalsMetals

Metal oxides are basic: CaO(s) + H2O Ca(OH)2 Na2O(s) + H2O 2NaOH

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NonmetalsNonmetalsWhen nonmetals react with metals, nonmetals tend to gain electrons metals tend to lose electrons:

metal + nonmetal salt e.g. 2Al(s) + 3Br2(l) 2AlBr3(s)

Mg(s) + P ? K(s) + O2 (g) ? Ca(s) + Cl2 ?

Nonmetal oxides tend to be acidic:

CO2 + H2O H2CO3

NO2 + H2O HNO3

SO3 + H2O H2SO4

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SemimetalsSemimetals

Semimetal oxides tend to be amphoteric (they can react either as weak acids or weak bases):

B2O3 + 3H2O 2B(OH)3

As an acid:B2O3 + 6NaOH 2Na3BO3 + 3H2O strong base

As a base:B2O3 + 4HCl 2Cl2B(OH) + H2O strong acid

1 electronic configurations and the periodic table chapter 7 robert w. bunsen 1811-1899. established...

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