Begin Chapter 1 Atomic Structure Emission Line Spectra Historical

1

Begin Chapter 1

Electronic Structure of the Atom (A review)

Atomic Structure

• Historical Perspective (Newton, Bunsen, Bohr, Schrodinger)• Rutherford’s experiments• Bohr model –> Interpretation of hydrogen atom spectra• Wave - particle duality• Wave mechanics– Heisenberg’s uncertainty principle– Electron density and orbitals– The Schrödinger equation and its solutions– Electron spin, the Pauli principle, Hund’s rule– Aufbau principle– Effective nuclear charge, shielding and penetration– Structure of the periodic table

Continuous Spectrum

Line Spectrum

Line Spectrum forHydrogen (Visible Region)

GasDischargeTube

Emission Line Spectra HistoricalIsaac Newton – (circa 1700) passage of sunlight through a prism

produced a continuous visible spectrum

Continuous Spectrum

Robert Bunsen – (1860) (1) observed that the emission spectra, rather than being continuous,

Instead, were series of colored lines (line spectra)(2) each chemical element produced a unique and characteristic spectrum

Line SpectrumGasDischargeTube

Example of Line Spectrum

Emission and Absorption Spectra

• Atoms are excited either via electrical discharge (A) or with a light source (B)

(B)

(A)

Other Investigators – several sets of spectral lines for H atom: in UV, VIS, and IR

How do we explain these line spectral phenomena? Introduction to Neils Bohr

Historical

2

Bohr Model (1913)

• First model that could account for the spectra of atomic hydrogen

+

- me

mnEnergy = kinetic energy + potential energy

centrifugal force = coulombic attraction

-The single hydrogen electron is located only in certain allowed orbits circlingaround the nucleus

-These orbits are symbolized by n = 1, n = 2, n = 3, n = 4, etc.-These orbits have fixed or “quantized” energies (in Joules)

Bohr’s Model of the Hydrogen Atom

-The single hydrogen electron is located only in certain allowed orbits circlingaround the nucleus-These orbits are symbolized by n = 1, n = 2, n = 3, n = 4, etc.-These orbits have fixed or “quantized” energies (in Joules)

E = - RH/n2 where n = 1,2,3…∞ and Rydberg’s constant for H is RH = 2.179 x 10-18 J

Bohr explained the emission of light

(1) Absorption

(2) Emission

nexcited

n1

Absorption Emission

When energy is absorbed by an atom from a flame or electrical discharge, electronsmove from 1 quantum number to one or more higher energy levels. When electronsreturn to ground state, light is emitted and wavelength of emitted light correspondsto energy of separation between initial and final quantum levels.

Bohr introduced quantization

• Atomic spectra of hydrogen is not continuous but consists of discrete lines–> Bohr suggested that the electron can adopt only certain distances r (orbits)

where k is a constant (Bohr radius = 52.9 pm, also a0), and n is any integer = QUANTUM NUMBER of the orbit

• Each allowed orbit corresponds to a different energy level: i.e. Bohr introduced quantization –>(E, electron energy; n, quantum number; Z, nuclear charge; k’, constant)

En = - me4Z = - k’8ε0

2h2n2 n2

m =9.10939 x 10–31 kg 1 J = 5.034 x 1022 cm–1

e = 1.60218 x 10–19 C 1 eV = 1.602 x 10–19 Jh = 6.62608 x 10–34 J·s ε0 = 8.85419 x 10–12 F·m–1

Permittivity of free space

More mathematical backgroundFrom centrifugal force = coulombic attraction 4πε0rmv2 = Ze2

Quantizing of the angular momentum: mvr = nh2π

Energy = kinetic energy + potential energy E = ½ mv2 – 1 . Ze2

4πε0 r

With these three equations, the radius, the energy and the velocity of the electron of the H atom with quantum number n (and nuclear charge Z) can be calculated:

m =9.10939 x 10–31 kg 1 J = 5.034 x 1022 cm–1

e = 1.60218 x 10–19 C 1 eV = 1.602 x 10–19 Jh = 6.62608 x 10–34 J·s ε0 = 8.85419 x 10–12 F·m–1

From Equations

With these equations,

• Radius

• Energy.

• Velocity of the electron

of the H atom with quantum number n (and nuclear charge Z) can be calculated:

Energy levels

• Each orbit corresponds to a specific energy level

• The allowed energies are often displayed in a energy level diagram:

Note: The energy levels are negativenumbers and indicate the energy of anelectron in the corresponding orbit:

->This is the energy required to removethis electron from the orbit ( = Ionization energy).

->The principle quantum number ndetermines this energy value.

3

Emission Spectrum of Hydrogen

Transitions between energy levels

RH = 13.06 eV or 109678 cm–1

1st four Bohr radii

Where• k is a constant (Bohr radius = 52.9 pm, also a0), • n is any integer = QUANTUM NUMBER of the orbit• Z = nuclear charge• R = Rydberg Constant (13.605 eV)

E = - RH1n2

Conclusions• Ionization energy is proportional to Z2 (nuclear charge)

• Radius of hydrogen atom in ground state (n=1) is 52.9 pm (= Bohr radius)The radius is inversely proportional to Z

• The excited state radius is proportional to n2

• In the ground state the electron has a velocity, v = 2.187•108 cm/s

• Bohr’s model allows accurate prediction of the hydrogen atom spectrumBUT fails to(1) describe atoms or ions with more than ONE electron(2) explain the splitting of the spectral lines in a magnetic field (Zeeman effect)(3) not include the uncertainty principle, which says that we cannot determine

both the position and the velocity of the electron at the same time, which theBohr model requires us to do.

• New theory is required, which is not based on classical mechanics

Electron: Particle or Wave?

• Depending on the experimental conditions, electrons appear either as particles, or show properties usually associated with waves

– Electrons are diffracted by crystalline materials, just as observed with X-rays

– De Broglie relationship: moving particles could exhibit wavelike properties

h: Planck’s constantmev: momentum of electron

• The photoelectric effect revealed a linear relationshipbetween the kinetic energy of the photon and frequency:

E = ½ mev2 = h(v-v0) ∆E = hν

From Orbits to Orbitals: The Uncertainty Principle

• Werner Heisenberg (1927):

It is not possible to determine simultaneously the position and momentum of anelectron with good precision:

• But: the probability of finding an electron at a particular point can be calculated

• this probability distribution is called an ORBITAL rather than an orbit.

Representation of Orbitals

4

The Schrödinger Equation

• The probability distribution and energy levels for electrons atoms andmolecules can be calculated using the Schrödinger equation:

H Ψ = E ΨH: Hamilton Operator (Energy)

E: Energy of solution ΨΨ: wavefunction

• Each solution Ψ of the equation corresponds to a differentelectron probability distribution with a distinct energy E

• The probability of finding an electron at some point isproportional to ΨΨ* (Ψ* is the complex conjugate of Ψ)

Probability of Finding an Electron About the Nucleus

• Ψ2 is the probability of finding an electron at any point in the regionsurrounding the nucleus

• Therefore, electrons no longer define specific orbits around the nucleusof an atom, but rather atomic orbitals!

Quantum Numbers

Quantum Numbers are n l ml ms

Principal Quantum Number (n):• determines the energy of an electron and the size of orbital• n = 1, 2, 3, 4, … (whole numbers)• maximum number of electrons in an energy level (or shell) = 2n2

Angular Momentum Quantum Number (l):• determines shape of the orbital• allowed values of l are 0, 1, 2, … up to (n – 1)• a shell (n) contains subshells

Quantum Numbers

Magnetic Quantum Number (ml)• determines orientation of orbital in space• allowed whole number values of ml are from – l to + l

if l = 0, then ml = 0 (1 orbital in s subshell)if l = 1, then ml = –1, 0, and +1 (3 orbitals in p subshell)if l = 2, then ml = –2, –1, 0, +1, and +2 (5 orbitals in d subshell)

• number of allowed orbitals in a subshell is (2 l + 1)

Electron Spin Quantum Number (ms)• determines direction (orientation) of electron spin in an orbital• values of ms are – ½ (spin down) and +½ (spin up)• each orbital can hold a maximum of 2 electrons

Give the values of the quantum numbers and number of orbitals associated with the following subshells. Also how many electrons are in these subshells?

2p:5s:5d:

Which one of the following sets of quantum numbers (n, l, ml, ms) is valid?(3, 3, –1, –1/2)(3, 0, 1, –1/2)(0, 0, 0, +1/2)(2, 1, –1, +1/2)(2, –1, 0, –1/2)

Quantum Numbers – Examples Shapes of the Atomic OrbitalsAtomic Orbital –

Hydrogen 1s and 2s Orbitals

5

Shapes of the Atomic Orbitals

2p subshell contains 3 orbitals:

2px 2py 2pz

Shapes of the Atomic Orbitals

3d subshell contains 5 orbitals: 3dx2–y2, 3dz2, 3dxy, 3dxz, 3dyz

Polyelectronic Atom

Aufbau Principle –The electrons fill up the available energy levels(orbitals), starting with the lowest available level

Pauli Exclusion Principle – No two electrons can possess identical setsof the 4 quantum numbers

Order of Filling Orbitals by Electrons in Many-Electron Atoms (Memorize!)1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s 7p

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

Writing Electron Configurations for Neutral Atoms using the Periodic Table

What is the electron configuration for P?

What is the electron configuration for Ba using the noble-gas corenotation?

Hund’s Rule

• There is often more than one way of arranging electrons within a set of degenerate (=energetically equivalent) orbitals• –> they correspond often to different energies• Hund’s rule of maximum multiplicity:

“The ground state of an atom will be the one having the greatest multiplicity”

Electron Configuration for Atoms

• Some atoms do not have theelectron configuration you would expect based on the Aufbau principle:

• Half filled and filled subshells(here the five 3d orbitals)provide additional stabilization

[Ar] 4s13d10 is the lower energy configuration compared to 4s23d9

V = [Ar] 4s23d3

Cr = [Ar] 4s13d5

Mn = [Ar]4s23d5

Fe = [Ar]4s23d6

Ni = [Ar]4s23d8

Cu = [Ar]4s13d10

Zn = [Ar]4s23d10

6

Aufbau Principle(Filling Principle)

1. Based on the Pauli principle,the distribution of electronsamong the atom orbitals can be determined (= electron configuration)

2. The electrons fill up theavailable energy levels(orbitals), starting with thelowest available level

Ener

gy

Exceptions to the Aufbau Principle

Exceptions to the Aufbau Principle (i.e. electron(s) stolen from an s orbital topartially fill or completely fill a d orbital):

Cr, Cu, Nb, Mo, Ru, Rh, (Pd), Ag, Pt, and Au

For those in bold, only 1 electron stolen from ns to fill (n-1)dFor Pd only, 2 electrons stolen from 5s to fill 4d

Electron Configurations:Cr: [Ar] 4s1 3d5Cu: [Ar] 4s1 3d10Nb: [Kr] 5s1 4d4Mo: [Kr] 5s1 4d5Ru: [Kr] 5s1 4d7Rh: [Kr] 5s1 4d8Pd: [Kr] 5s0 4d10 (2 electrons stolen from 5s to fill 4d)Ag: [Kr] 5s1 4d10Pt: [Xe] 6s1 4f14 5d9Au: [Xe] 6s1 4f14 5d10

Writing Electron Configurations

•What is the electron configuration for Pb using the noble-gas core?

•What is the electron configuration for Au using the noble-gas core?

•What is the electron configuration for Cu using the noble-gas core?

•What is the electron configuration for W using the noble-gas core?

•What is the electron configuration for Fe using the noble-gas core?

Ion Electron Configurations

ANIONSO2–:

CATIONSCa2+:

Mn2+:

Electron Configuration for Ions

• Ease of removal of an electron (=Ionization energy) often does not mirror the filling order embodied in the Aufbau principle

Mn [Ar] 4s23d5

Mn2+ [Ar] 4s03d5 and not [Ar] 4s23d3

Different electron configurations correspond to different species, and have different properties (color, magnetic behavior, etc.)

Note: All transition metal atoms lose their ns electrons before their (n–1)delectrons!

Cl–

Pb2+

Fe2+

Mo6+

Ni2+

Ion Electron Configurations

7

Magnetic Properties of Atoms

Diamagnetic Substance: If atoms containing only spin-paired electrons are placed in a magnetic field, they are weakly repelled by the field = phenomena is diamagnetismNe: 1s2 2s2 2p6: ____ ____ ____ ____ ____

1s 2s 2p 2p 2p

Paramagnetic Substance: atoms containing one or more unpaired electrons are attracted by the magneticN: 1s2 2s2 2p3: ____ ____ ____ ____ ____

1s 2s 2p 2p 2p

Ferromagnetic Substance: Unpaired electrons are aligned with their neighbors even in the absence of a magnetic field.

(e.g., recording media chromium (IV) oxide, γ-iron(III)oxide)

Chapter 2: Periodic Table, Overview, and Trends in Atomic Properties

•Periodic Table history“Triads” –“Law of Octaves” –Mendeleev and Meyer –Flaws in Mendeleev’s Periodic TableMoseley –

•Effective nuclear charge•Ionization energies•Electron affinities•Covalent and ionic radii•Bond strength•Electronegativity•Orbital energies•Promotion energies•Common oxidation states•Relativistic effects

Organization of the Modern Periodic Table

• Elements placed in order of increasing atomic number (# of protons)

Electron Affinity• Energy change associated with addition of an electron to the lowest energy unoccupied orbital of a free atom

X(g) + e- X-(g) 1st electron affinity (EA)• Favorable process for most elements

• Influenced by Zeff and atom size (principle quantum number)

• Note: Positive sign per definition:F + e- F- EA = + 328 kJ/mol ∆H = – 328kJ/mol

• Trends in EAs parallel those of IPs• Exceptions:

– EA of F is lower than of Cl– EA of N is lower than of P– EA of O is lower than of S

–> smaller size of F (and N or O)causes greater electron-electronrepulsion!

Electron AffinityElectron-Gain Enthalpy ΔHeg = enthalpy change when a gas-phase

atom gains an electron as shown below:A(g) + e- (g) A- (g)

Electron gain may be exothermic or endothermic.

Although electron-gain enthalpy (ΔHeg) is the thermodyanmically appropriate termmuch of inorganic chemistry is discussed in terms of a closely related property (Electron Affinity, EA) of an element, which is the negative of itselectron gain enthaply.

If EA of an element A is higher than B, then electron gain is more exothermic for A than B

Note: electron gain enthalpy (kJ/mol) whereas electron-gain energy ( eV)

An element has a high EA if the additional electron can enter a shell where it experiences a strong effective nuclear charge (e.g., top right of periodic table)Therefore elements close to F can be expected to have the highest EA values

Electron Affinity (more)The 2nd electron-gain enthalpy (the enthalpy change for the attachment of a 2nd e-

to an initially neutral atom, is invariably endothermic since the e-repulsionoutweighs the nuclear attraction

8

Trends in Electron AffinitiesAdded e- - e- pair repulsion from 2p3 to 2p4

High Zeff for 2p electronsoutweighs interelectron

repulsion factor

Addition from 2p2 to 2p3,energy advantage

Ionization EnergyIonization Energy (I) = the ease with which an e- can be removed from an atom is

measured by its ionization energy, I, the minimum energy needed to removean e- from a gas-phase atom: A (g) A+ (g) + e- (g)

1st IE = ionization of the least tightly bound e- from the neutral atom2nd IE = ionization of the resulting cation, and so on

IE conveniently expressed in eV, where 1 eV is the energy acquired by an e- when itfalls through a potential difference of 1 V. Hence, 1 eV = 96.5 kJ/molExample….Ionization Energy of H atom is 13.6 eV, so to remove an e- from

a H atom is equivalent to moving the e- through 13.6 V.

In thermodynamic calculations, it is convenient to use the IONIZATION ENTHALPY,the ionization enthalpy is TR larger than the ionization energy

Because RT is only 2.5 kM/mol at room temperature, the difference betweenionization energy and enthalpy can often be ignored

Ionization energies (eV) versus Ionization enthalpies (kJ/mol)

Periodic Trends in Ionization Energies

Half-filled Filled shell

IPs do not uniformly increase fromleft to right–> changing orbital and spin pairingbreaks trendIPs do not always decrease goingdown a group–> transition series and actinidesupset this trend

Generally, Ionization Energies vary systematically through the periodic table, being smallest at lower left (near Cs) and greatest at upper right (near He). Variation follows pattern of effective nuclear charge in connection with Aufbau principle,But yet a few exceptions.

Covalent Radii

• Covalent radius of an atom A = half the distance between a diatomic molecule A–A

• They are approximately additive: For a non-polar molecule A–B the sum of the radii of A and B should be equal to the A–B bond length

• Note: For polar molecules this approximation does not work as well!

Electronegativity (EN)

• Pauling (1930):

Electronegativity = “the ability of an atom to attract electron density towards itself in a molecule”

– Not amenable to direct experimental measurement– But: very useful concept which allows to predict, whether a givencombination of elements is likely to result in a molecule with polar bonds

• Various quantifications of EN:

– Pauling: based on bond-strength– Alfred-Rochow: based on size and effective nuclear charge– Mulliken: based on IPs and EAs– Allen’s spectroscopic values: based on orbital energies

Periodic Trends in Electronegativity

9

Promotion Energy

Carbon Atom:1s22s22p2 –> has only 2 unpaired electrons, but forms 4 covalent bonds

Prior to bonding interaction, one of the 2selectrons is promoted to the empty 2porbital (promotion energy):

1s22s12p3 –> 4 unpaired electrons canform 4 bonds

Note: Promotion energies increase whengoing down the B or C group

Many atoms form more bonds than expected based on the number of unpaired electrons in their ground states

Relativistic EffectsImportant for heavy metal elements:

(1) Schrodinger’s equation fails to take into account the effects of relativity on the electrons.

(2) Einstein’s relativity theory: Objects moving close to the speed of light increase in mass

(3) Due to the high nuclear charge of heavy elements, electrons close to the nucleus (s orbitals!) have a large velocity. Ramifications?

–> mass of electron increases–> effective size of orbital decreases (relativistic orbital contraction)

s and p orbitals contractd and f orbitals expand (more strongly shielded)

(4) The contraction of the s orbitals (and somewhat also the p orbitals) leads to an expansion of the d and f orbitals due to increased shielding of the nuclear charge

Organization of the Modern Periodic Table – Terminology

Period –

Group (or Family) –

Main Group Elements –-Alkali Metals-Alkaline Earth Metals-Pnicogens-Chalcogens-Halogens-Noble Gases

Transition Metals –

Lanthanoids and Actinoids –

Existence of the Elements

Big Bang Theory –

Exothermic Nuclear Reactions –

Endothermic Nuclear Reactions –

Stability of the Elements and Their Isotopes

Only 81 stable elements!

Why?

Quantum (or Shell) Model of the Nucleus

Quantum (or Shell) Model of the Nucleus

Magic Number –For Electrons: 2, 10, 18, 36, 54, 86For Protons: 2, 8, 20, 28, 50, 82, (114)For Neutrons: 2, 8, 20, 28, 50, 82, 126

Examples:

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Classification of the Elements

Standard Temperature and Pressure (STP) – 0 C, 101 kPa

Standard Ambient Temperature and Pressure (SATP) – 25 C, 100 kPa

Criteria for Classification of Metals and Nonmetals1) Luster2) Density3) Hardness4) Malleability or Ductility5) Thermal Conductivity6) Three Dimensional Electrical Conductivity

Semimetals: B, Si, Ge, As, and Te

Weak metals: Be, Al, Zn, Ga, Sn, Sb, Bi, Po

Periodic Properties

1) Effective Nuclear Charge (Zeff)2) Atomic Radius3) Ionization Energy4) Electron Affinity

Atomic RadiusIncreases

ZeffIonization EnergyElectron Affinity

INCREASES

Orbital Energies

• Energy of orbitals in a multielectron atom is not identical with the corresponding orbitals of the hydrogen atom–> the electrons experience an effective nuclear charge (Zeff) that is different from 1.0.

• Each electron is attracted by the nucleus and repelled by other electrons

Zeff = Z −σ

Z+

attraction

attraction

repulsion shieldingnuclear charge

Effective nuclear charge

-

-

Shielding or Screening and Zeff

Inner-core electrons act as a shield for outer-core electrons farther out from the nucleus

Effective Nuclear Charge (Zeff)

-Increases going left to right across a period-Decreases going down a group

Shielding

The shielding constant isdependent on which orbitalsare occupied:

1s electrons shieldthe third 2s electron from the full effect of the nuclear charge

Radial distribution function for Li

Slater’s Rules

Rules for determining “σ = Slater’s screening factor”1) The electronic structure of the atom is written in groupings as follows:

(1s) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4d) (4f) (5s, 5p), etc.2) Electrons in higher groups (to the right in the list above) do not shield

those in lower groups (they contribute zero)3) For ns or np valence electrons:

a) Other electrons in the same (ns, np) group contribute 0.35b) Electrons in the n–1 group contribute 0.85c) Electrons in the n–2 group contribute 1.00

4) For nd or nf valence electrons:a) Other electrons in the same (nd) or (nf) group contribute 0.35b) Electrons in all the other groups to the left contribute 1.00

Calculate Zeff for on of the 2p electrons in the oxygen atom (1s22s22p4)σ = (2 x 0.85) + (5 x 0.35) = 3.45Zeff = Z – σ = 8 – 3.45 = 4.55

Zeff = Z – σ

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Periodic Properties: Atomic Radius

What is Atomic Size?1) Covalent radius (rcov) – ½ distance between the nuclei of 2 atoms of the same element2) van der Waals radius (rvdw) – ½ distance between the nuclei of 2 atoms of neigboring molecules3) metallic radius – ½ distance between the nuclei of 2 neighboring atoms in the solid metal

Factors:1) Zeff2) increasing size of orbitals

Covalent Radius Covalent RadiusF: 64 pm Li: 134 pmCl: 99 pm Be: 106 pmBr: 114 pm B: 88 pmI: 133 pm C: 77 pm

N: 70 pmO: 66 pm

Atomic Radius and Relativistic Effects

Variations in Atomic RadiusAluminum (rcov = 126 pm) and Gallium (rcov = 126 pm):

Relativisitic Effects in Heavier Elements (Period 6)When Z = 1, velocity of electrons are ~7.3 x 10–3

When Z = 80, velocity of electrons are ~0.58

Therefore, electrons are ~20% heavier when Z = 80

1) Relativistic Contraction of s orbitals

2) Relativistic Expansion of d and f orbitals

Overall, Atomic Radius Decreases

Cr (129pm) Mo (140 pm) W (141 pm)

Periodic Properties: Ionization Energy

Ionization Energy (IE) –the ease with which an e- can be removed from an atom is measured by its ionization energy, I, the minimum energy needed to remove an e- from a gas-phase atom (IE can be measured with great precision)

A (g) A+ (g) + e- (g)

First Ionization Energies (IE1):energy (IE1) + Li Li+ + e– IE1 = 520 kJ/molenergy (IE1) + Ne Ne+ + e– IE1 = 2080 kJ/mol

Factors:1) IE increases going left to right across a period because increase in Zeff.2) IE decreases going down a group because electrons in inner orbitals shield

electrons in outer orbitals and the successive outer orbitals themselves arelarger.

Electron Affinity• Energy change associated with addition of an electron to the lowest energy unoccupied orbital of a free atom

X(g) + e- X-(g) 1st electron affinity (EA)• Favorable process for most elements

• Influenced by Zeff and atom size (principle quantum number)

• Note: Positive sign per definition:F + e- F- EA = + 328 kJ/mol ∆H = – 328kJ/mol

• Trends in EAs parallel those of IPs• Exceptions:

– EA of F is lower than of Cl– EA of N is lower than of P– EA of O is lower than of S

–> smaller size of F (and N or O)causes greater electron-electronrepulsion!

Ionic Radii

Ion sizes follow similar trends

BUT: Changes in charge have a significant impact on size

Radius (Å)Covalent +2 ionic +4 ionic

Ge 1.22 0.87 0.67Sn 1.40 1.12 0.83

Bond Strength

• The bond energy is the energy needed to break a chemical bond

• Strong bond: > 800 kJmol–1• Average bond: 500 kJmol–1• Weak bond: < 200 kJmol–1

• Bond strength often depends upon the size of the elements thatare bonded together

–> bond strength often decreases going down a group

Example:HF 568 kJ/mol, HCl 432 kJ/mol, HBr 366 kJ/mol, HI 298 kJ/mol

12

Begin Chapter 9

Periodic Trends

Alkali Metals

• Melting/Boiling point trends• Ionization energy, Hydration enthalpy,

electrode potential• Chemical Reactivity trend with water

2M(s) + 2H20(l) -> 2MOH(aq) + H2(g)

Alkali

69029Cs70039Rb75964K89298Na1330180LiBP (C)MP (C)

Trends: (1) Decrease in mp and bp down the group(2) Why trend? Weakening metallic bond as the atomic

radius of the metal increases

Alkalis (Reactivity,)

2 M(s) + 2 H2O (l) 2 MOH (aq) + H2 (g)

Examples of increasing reactivity down the group(Li) bubbles quietly to produce the hydroxide and H2 gas(Na) melts,skating around on the water surface as a silvery globule, and

the H2 gas produced usually burns(Heavier Alkalis) rxn is extremely violent: explosions often occur when small

chunks of Rb and Cs are dropped into water (explosions fromignition of H2 gas)

IE and Hydration Enthalpies of all Alkalis are low and parallel one another

Alkalis (Other Trends)IE and Hydration Enthalpies of all Alkalis are low and parallel one another

-3.03-276 (Cs+)376Cs

-2.92-301 (Rb+)403Rb

-2.94-322 (K+)425K

-2.71-406 (Na+)502Na

-3.04-519 (Li+)526Li

Electrode Potential

(V)

Hydration Enthalpy (kJ/mol)

Ionization Energy (kJ/mol)

Halogens

• Reverse trend in melting/boiling points than Alkalis

Element M.P. (C) B.P. (C)

F -219 -188Cl -101 -34Br -7 60I 114 185

Why?

13

Halogens: Explanation of m.p./b.p. Trend

Why? Intermolecular forces between neighboring diatomic molecules.

Dispersion Forces increase in strength with number of electrons.

Dispersion (London) Forces

(a) (b) (c)δ+δ- δ- δ- δ+δ+

δ+ δ+ δ-δ-

(d)

Average electron density for an atom

Instantaneous electron density produces a temporary dipole

Instantaneous attraction between neighboring molecules

Reverse of polarity in the next instant

Summary Dispersion Forces

• Strength of dispersion forces relates to number of electrons

• Number of electrons determines how easily the electron density can be polarized

• The greater the polarization, the stronger the dispersion forces

• The stronger the intermolecular forces, the higher the melting and boiling points

Halogens – Physical PropertiesElement Bond Electron Electrode

Energy Affinity Potential(kJ/mol) (kJ/mol) (acidic, V)

F 155 -328 +3.05

Cl 240 -349 +1.36

Br 190 -331 +1.09

I 149 -301 +0.54

Fluorine breaks trend

Chemical Reactivity is ReverseOrder of Alkali Metals

Great OxidizingAgent

Halogens – Chemical Reactivity

• Why? Reactivity related to Bond Energy and ElectronAffinity

• Chemical Reactivity is Reverse Order of Alkali Metals(e.g. reaction with hydrogen)

X2(g) + H2(g) -> 2HX(g)

F2(g) + H2(g) -> 2HF(g) Explosive (rocket fuel)

Cl2(g) + H2(g) -> 2HCl(g) Violent, but need catalyst (light)

Br2(g) + H2(g) -> 2HBr(g) Slow

I2(g) + H2(g) <-> 2HI(g) Equilibrium Mixture

Group V Elements

Change from nonmetallic to metallicbehavior (see large jump in m.p/b.p.)

1564271Bi

1387631Sb

615 (sub)As

28144P (P4)

-196-210N

Boiling Point (C)

Melting Point (C)

Element

N2: Only weak dispersion forces, Hence, very low mp/bpP4: Compared to N2, higher mp/bp – more electrons and cluster of 4

atomsAs,Sb,Bi Grey solids with electrical conductivity increasing down the series.As Sublimes to As4 (like P4)Sb, Bi long liquid range (characteristic for metals)

14

Group V – Chemical Reactivity

• No consistent Trend. Why? Members don’t share common bond type

N2 (Gas, Diatomic)P4 (Solid, White Phosphorus Allotrope)As (Solid, Layer Structure of Network

Covalent Bonding)Sb,Bi More interaction between layers in

Solid State, giving predominantly Metallic Bonding Type

Group V – Oxoanion Formation

• Oxoanions easyNO3

-, PO43-, AsO4

3-

• Oxoanions difficult (prefer cations like metals)Sb, Bi

Row 2 Trends

-249-219-229-2104100(sub)

21801287180

NeF2O2N2CBBeLi

Melting Point Trends

Li, Be Shiny with high electrical conductivityLi only 1 outer electron, larger radius than Be: Li has weak metallic bonding

resulting in low mp and high chemical reactivityBe 2 outer electrons and very much smaller radius, stronger metallic bond

resulting in higher mp.B NOT metallic bonding but classified as semimetallic

Row 3 Bonding Trends

• Na, Mg, Al, Si, P, S, Cl, Ar

MetallicBonding

NetworkCovalentBonding

SingleCovalentBonding

P4 (white waxy solid bound together by single covalent bonds)S8 (yellow solid, atoms held together by single covalent bonds)Cl2 (gas, diatomic, simple covalent bonds)Ar (gas, monatomic)

Row 3 Melting Points

-189-10111944142066064998m.p.(oC)

ArCl2S8P4SiAlMgNa

Weakness of dispersion forces -> low melting points

Metallic Bonding NetworkCovalentBonding

Increasing m.p. Increasing m.p.

SingleCovalentBonding

Rows 2,3 – Bonding Trends, F

Covalent(gas)

Covalent(gas)

Covalent(gas)

Covalent(gas)

NetworkCovalent(solid)

Ionic(solid)

Ionic(solid)

Bonding Type (phase)

ClF5SF6PF5SiF4AlF3MgF2NaFCompound

n/aCovalent(gas)

Covalent(gas)

Covalent(gas)

Covalent(gas)

NetworkCovalent(solid)

Ionic(solid)

Bonding Type (phase)

n/aOF2NF3CF4BF3BeF2LiFCompound

Very low m.p., weak intermolecular forces (dispersion and dipole-dipole)

Ionic compoundsm.p. high, must break ionic bonds in crystal lattice

NetworkCovalent, m.p. high, covalent bonds broken

No C

lF7

(Geom

etry Issue, 7 C

laround 1F)

15

Rows 2,3 – Bonding Trends, O

Covalent(liquid)

Covalent(solid)

Covalent(solid)

Network Covalent(solid)

Ionic(solid)

Ionic(solid)

Ionic(solid)

Bonding Type (phase)

Cl2O7(SO3)3P4O10SiO2Al2O3MgONa2OCompound

Covalent(gas)

n/aCovalent(gas)

Covalent(gas)

Network Covalent (solid)

Ionic(solid)

Ionic(solid)

Bonding Type (phase)

F2On/aN2O5CO2B2O3BeOLi2OCompound

Very low m.p., weak intermolecular forces (dispersion and dipole-dipole)

Ionic compoundsm.p. high, must break ionic bonds in crystal lattice

NetworkCovalent, m.p. high, covalent bonds broken

Electronegativity Higher

Rows 2,3 – Bonding Trends, O

>+270(Decomposes explosively)

-371-2700-856ΔGfo

Cl2O7(SO3)3P4O10SiO2

+42n/a+115-386-1194ΔGfo

F2O---N2O5CO2B2O3

• Stability of covalent oxides decreases left to right -> see free energies of formation table• N, F, Cl are strong oxidizing agents, themselves being easily reduced• More positive ΔG, the less stable the compound

(ΔGfo)

Rows 2,3 – Bonding Trends, H

Covalent(gas)

Covalent(gas)

Covalent(gas)

Covalent(gas)

NetworkCovalent(solid)

Ionic(solid)

Ionic(solid)

Bonding Type (phase)

HClH2SPH3SiH4(AlH3)xMgH2NaH(Stable in dry air, but reaction with H2O

Compound

Covalent(liquid)

Covalent(liquid)

Covalent(gas)

Covalent(gas)

Covalent(gas)

NetworkCovalent(solid)

Ionic(solid)

Bonding Type (phase)

HFH2ONH3CH4B2H6

Diborane(Spont. Burns in air)

(BeH2)xLiHCompound

Free Energy of Formation (Hydrides)

-95-34+13+57ΔGfo

HClH2SPH3

(Phosphine,Ignites from trace impurities

SiH4

(Silane, spontaneously flammable)

Compound

-275-237-16-51+87ΔGfo

HFH2ONH3CH4B2H6Compound

Need ignition source

Flammability of nonmetal hydrides correlates with the free energy of formation

Acid-Base TrendsBasic Amphoteric Acidic

AcidicAcidicAcidicAcidicAmphotericBasicBasic

Cl2O7(SO3)3P4O10SiO2Al2O3MgONa2O

2Na2O (s) + H2O (l) NaOH(Base)

SO3 (s) + H2O (l) H2SO4(Acid)

(Al, like Zn, Sn) => both metal&nonmetal properties

• Oxides react with acids (like metals)Al2O3(s) + 6H+ (aq) 2Al3+ (aq) + 3H2O (l)

• Oxides react with bases (like nonmetals)Al2O3(s) + 2OH-(aq) + H2O (l) [Al(OH)4]- (aq)

Acid-Base (Group V)

BasicBi2O3

AmphotericSb2O3

AcidicAs2O3

AcidicP4O10

AcidicN2O5

Non-Metal Oxides

16

Acid-Base

StronglyAcidic

Very weaklyAcidic

NeutralNeutral

HClH2SPH3SIH4

WeaklyAcidic

---BasicNeutralNeutral

HFH2ONH3CH4B2H6

Non-Metal Hydrides