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Chapter 7Chapter 7

Atomic StructureAtomic Structure

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Our goal:• Understand why some substances behave as they do.• For example: Why are K and Na reactive metals? Why do

H and Cl combine to make HCl? Why are some compounds molecular rather than ionic?

Atom interact through their outer parts, their electrons.

The arrangement of electrons in atoms are referred to as their electronic structure.

Electron structure relates to:• Number of electrons an atom possess.

• Where they are located.• What energies they possess.

Electronic StructureElectronic Structure

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Study of light emitted or absorbed by substances has lead to the understanding of the electronic structure of atoms.

Light made up of electromagnetic (E.M) radiation

Characteristics of light: • All waves have a characteristic wavelength, λ, and amplitude, A.

• The frequency, ν, of a wave is the number of cycles which pass a point in one second. Measured in hertz , 1 hertz = 1 cycle/second

• The speed of a wave, v, is given by its frequency multiplied by its wavelength: λ α (1/ ν) λ ν = constant (c)

c = λ ν c : speed of light = 3 x 108 m/s

The Wave Nature of LightThe Wave Nature of Light

Identifying λ and ν

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

• Electromagnetic (E.M.) radiation moves through a vacuum with a speed of 2.99792458 × 108 m/s.

• There are many kind of E.M. radiation with different wavelengths and frequencies shown in the following figure.

• Visible radiation is the only part our eye can detect. It has wavelengths between 400 nm (violet) and 750 nm (red).

Electromagnetic RadiationElectromagnetic Radiation

The Electromagnetic Spectrum

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Example(1):

What is the wavelength of light with a frequency 5.89 x 105 Hz.

Example (2):

What is the frequency of blue light with a wavelength of 484 nm?

λ =cν = 3 x 108 m/s

5.89 x 105 s-1 = 509 m (Radio wave)

ν =cλ = 3 x 108 m/s

484 x 10-9 m = 6.2 x 1014 s-1 or Hz

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In 1990 Matter and energy were seen as different from each other in fundamental ways

Matter: � consist of particles� Particles have a mass � Its position in space can be specified.

Energy: � could come in waves, with any frequency.� Massless and delocalized.� Their position in space could not be specified.

It was assumed that there was no intermingling of matter and light

The Nature of Matter The Nature of Matter

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At the beginning of 20 century, certain experimental results suggested that this picture was incorrectThe first important advance came from Max Planck, he found that the cooling of hot objects couldn’t be explained by viewing energy as a wave.Plank found that the results could not be explain in term of the physics of his day (matter absorb or emit any quantity of energy) .Plank account for these observation by postulating that:The energy can be gained or lost only in whole-number multiple of the quantity hhνν

∆ ∆ E = nhE = nhνν where n is an integer (1, 2, where n is an integer (1, 2, 3…).3…).

h : is Planck’s constant = 6.626 x 10h : is Planck’s constant = 6.626 x 10 -34-34 J s J sIt seemed clear that energy is It seemed clear that energy is quantizedquantized and can occur in discrete and can occur in discrete unit of sizeunit of size hhνν , these packets of energy (h , these packets of energy (hνν) are called ) are called quantum.quantum.A system can transfer energy only in whole quanta.

Thus energy seems to have particulate properties

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The next development came when Einstein proposed that E.M. radiation is itself quantized.

He suggested that E.M. radiation can be viewed as a stream of “particles” called photons

Each photon has energy Ephoton = hν = hc/λ

Combine this with E = mc2 (Einstein equation)

m = E /c2 m =

you get the apparent mass of a photon

m = h / (λc)

Does a photon really have a mass? The answer appears to be yes.

However, it is clear that photons do not have mass in the classical sense. A photon has mass only in relativistic sense – it has no rest mass.

hc/λc2

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We can summarize the important conclusions from the work of Plank and Einstien as follows: Energy is quantized, it can occur only in discrete unit called quanta. E.M. radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well.

This phenomenon is referred to as the dual nature of light

Is the opposite is true? That is, does matter exhibit wave properties.

de Brolie supplied the answer to this question.

m = h /(λc) for a particle with velocity v

m = h /(λv) λ = h/(mv)

This equation, called de Brolie equation, allow us to calculate the wave lenghth of particle

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Example:The laser light of a CD is 7.80 x 102 m. calculate

A) What is the frequency of this light?

B) What is the energy of a photon of this light?

C)What is the apparent mass of a photon of this light?

A) ν = c/λ ν = 3 x 108(m/s)/ 7.80x102 m = 3.85 x 105s-1

B) Ephoton = hν

Ephoton = 6.626 x 10-34 J s x 3.85 x 105s-1

Ephoton = 2.55 x 10-28 J

C) m = h / (λc)

m = 6.626 x 10-34 J s /7.80x102 m x 3x108(m/s)

m = 2.83x10-45 Js2/m2 = 2.83x10-45(kg m2/s2) s2/m2

m = 2.83x10-45kg

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Example:

What is the wavelength of an electron with a mass of

9.11 x 10-31 kg traveling at 1.0 x 107 m/s?

m = h / (λc)

λ = h /mc

λ = 6.626 x 10-34 J s

9.11 x 10-31 kg x 1.0 x 107 m/s = 7.2 x 10-11 m

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The atomic spectrum of hydrogenThe atomic spectrum of hydrogen

Another important experiment was the study of the emission light by excited H-atoms. When hydrogen gas receives high-energy spark, H2 molecules absorb energy, some of H-H bonds are broken. The resulting H-atoms are excited; that is they contain excess energy which they release by emitting light of various wavelength to produce what is called the emission spectrum of H-atoms.

To understand the significance of H-emission spectrum, we must describe the continuous spectrum that results when white light is passed through a prism,

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continuous spectrum� Contain all the wavelength of visible (white) light.� All the colors are possible.� Like the rainbow.

When H-emission spectrum in visible region is passed through prism, only a few lines can be seen, each correspond to a discrete wavelength. The H-emission spectrum is called line spectrum.

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Continuous spectrum

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Hydrogen spectrumHydrogen spectrum� Emission spectrum because these Emission spectrum because these

are the colors it gives off or emitsare the colors it gives off or emits� Called a line spectrum.Called a line spectrum.� There are just a few discrete lines There are just a few discrete lines

showingshowing

410 nm

434 nm

486 nm

656 nm

•Spectrum

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What is the significance of line spectrum of hydrogen?� It indicate only certain energies are allowed for the

hydrogen atom.� Energy of electron H-is quantized � Only certain energies are possible.� Use ∆E = hν = hc / λ

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BohrBohr Model ModelHe developed the quantum model of the hydrogen atom.� He proposed that the atom was like a solar system, the

electron in H-atom move around the nucleus only in certain allowed circular orbit

� The electrons were attracted to the nucleus because of opposite charges.

� Didn’t fall in to the nucleus because it was moving around

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� He didn’t know why but only certain energies were allowed.

� He called these allowed energies: energy levels.� Putting energy into the atom moved the electron away

from the nucleus from ground state to excited state.� When it returns to ground state it gives off light of a

certain energy� The energy levels for H-atom are shown in the following

figure.

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The Bohr Ring AtomThe Bohr Ring Atom

n = 3n = 4

n = 2n = 1

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The Bohr ModelThe Bohr Model

� for each energy level the energy is:

E = -2.178 x 10-18 J (Z2 / n2 )

n: is the energy level� Z: is the nuclear charge, which is +1 for hydrogen.� n = 1 is called the ground state� when the electron is removed from the atom, n = ∞� When the electron moves from one energy level to

another.

� ∆E = Efinal - Einitial

� ∆E = -2.178 x 10-18 J Z2 (1/ nf2 - 1/ ni

2)

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Example:

Calculate the energy need to move an electron from its ground state to the third energy level.

∆E = -2.178 x 10-18 J Z2 (1/ nf2 - 1/ ni

2)

∆E = -2.178 x 10-18 J (+1)2 (1/9 – 1/1)

∆E = +1.936 x 10-18 J (+ mean energy absorbed)

Example:

Calculate the energy released when an electron moves from n= 4 to n=2 in a hydrogen atom.

∆E = -2.178 x 10-18 J (+1)2 (1/4 – 1/16)

∆E = - 5.2125 x 10-19 J

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Example:

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

∆E = -2.178 x 10-18 J (+1)2 (1/∞ – 1/1)

∆E = 2.178 x 10-18 J

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Bohr model:� Only works for hydrogen atoms and other monoelectronic

species.� electrons don’t move in circles� the quantization of energy is right, but not because they

are circling like planets.

The negative sign of the energy level :� increase the energy of the electron when you make it

further to the nucleus.� the maximum energy an electron can have is zero, at an

infinite distance (n = ∞ ).

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The Quantum Mechanical Model of the atomThe Quantum Mechanical Model of the atom

A totally new approach was needed.

Three physicists were at the forefront of this effort: Heisenberg, de Broglie, and Schrödinger. The approach they developed known as wave mechanics or quantum mechanics

� De Broglie said matter could be like a wave.� Schrödinger proposed an equation that contains both

wave and particle terms.� Much math, but what is important are the solutions.� Solving the equation leads to wave functions.

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• The wave function is a F(x, y, z) Actually F(r,θ,φ)

• Solutions to the equation are called orbitals (not Bohr orbits).

• Each solution is tied to a certain energy level.

• The wave function gives the shape of the electronic orbital.

• The square of the wave function, gives the probability of finding the electron, that is, gives the electron density for the atom.

• There is a limit to what we can know from Schrödinger equation.

• We can’t know how the electron is moving or how it gets from one energy level to another.

Electron Density Distribution

•Probability of finding an electron in a

hydrogen atom in its ground state.

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Quantum NumbersQuantum Numbers

There are many solutions to Schrödinger’s equation

Each solution can be described with quantum numbers that describe some aspect of the solution.

� Principal quantum number (n):

has an integral value: 1, 2, 3, ……, it is related to the size and energy of an orbital.

As (n) increase: orbital become larger, electron spends more time farther

from the nucleus higher energy, because the electron is less tightly bound

to nucleus, energy is less negative.

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� Angular momentum quantum number (ℓ): has integer values from 0 to n-1 for each value of n It is related to the shape of the orbital (as shown in the

following figures the value of (ℓ) for a particular orbital is commonly

assigned a letter: ℓ = 0 is called s , ℓ = 1 is called p

ℓ =2 is called d , ℓ =3 is called ƒ , ℓ =4 is called g

� Magnetic quantum number (m ℓ):

– integer values between - ℓ and + ℓ including zero.

– The value of mℓ is related to the orientation of the orbital in space relative to the other orbitals in the atom.

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s-orbitalsAll s-orbitals are spherical.

1s 2s 3s

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P orbitals

There are three p-orbitals px, py, and pz.

The three p-orbitals lie along the x-, y- and z- axes of a Cartesian system.

The letters correspond to allowed values of mℓ of -1, 0,an +1.

Electron-distribution of a 2p orbital.

d-orbitals

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f-orbitals

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� Electron spin quantum number (m s):

the electron has a magnetic moment with two possible when the atom placed in an external magnetic field– Can have 2 values , either +1/2 or -1/2

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For our purpose, the main significance of electron spin is connected with the postulate of Pauli: in a given atom no two electrons can have the same set of four quantum numbers (n , ℓ, m ℓ, and ms ), this is called Pauli exclusion principle. Since electrons in the same orbital have the same value of n , ℓ, m ℓ , they must have different values of ms .

Then, since only two value of ms are allowed, an orbital can hold only two electrons, and they must have opposite spin.

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Quantum number for the first four level of orbitals in H-atom:

n ℓ Orbital designation

m ℓ No. of orbitals

1 0 1s 0 1

2 0 2s 0 1

1 2p -1 ,0 ,+1 3

3 0 3s 0 1

1 3p -1 ,0 ,+1 3

2 3d -2- ,1 ,0 ,+1,+2 5

4 0 4s 0 1

1 4p -1 ,0 ,+1 3

2 4d -2- ,1 ,0 ,+1,+2 5

3 4f -3-,2- ,1 ,0 ,+1 ,+2 ,+3

7

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Example: For n = 4, what are the possible values of ℓ.

ℓ = 0→ n -1 , so ℓ = 0 → 4-1

ℓ = 0, 1, 2, 3

s, p, d, f For ℓ = 2. What are the possible values of mℓ

mℓ = - ℓ → +ℓ mℓ = -2 → +2

mℓ = -2, -1, 0, +1, +2

How many possible values for ℓ and mℓ are there when

n = 3

ℓ = 3-1 = 2 ℓ = 0, 1, 2

for ℓ = 0 mℓ = 0 , , for ℓ = 1 mℓ = -1, 0, +1

for ℓ = 2 mℓ = -2, -1, 0, +1, +2

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Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Orbitals and Their Energies

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The Periodic Table The Periodic Table

� Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s)

� Didn’t know much about atom.� Put atoms in columns by similar properties.� Predicted properties of missing elements.

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Aufbau Principle and the periodic tableAufbau Principle and the periodic tableOur main assumption is that the atoms have the same type of orbitals as have been described from the hydrogen atom.

As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these H-like orbitals. This is called aufbau principle

H has one electron, occupy the 1s orbital

The configuration for H can be represent as:

H: 1s1

Helium has two electron

1s

Quantum no. for the electron is: n=1, ℓ = 0, mℓ =0, ms =+1/2

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Helium has two electrons

He: 1s2

Lithium has three electrons

Li: 1s2 2s1 2p

Be: 1s2 2s2 2p

B: 1s2 2s2 2p1

Quantum no. for the first electron is: n=1, ℓ = 0, mℓ =0, ms =+1/2

Quantum no. for the second electron is: n=1, ℓ = 0, mℓ =0, ms = -1/2

1s

1s 2s 2p

1s 2s 2p

1s 2s 2p

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C: 1s2 2s2 2p2

Two electrons occupy 2p orbital, since there are three 2p orbitals with the same energy, the mutually repulsive electrons will occupy separate 2p orbitals

Hund’s rule: the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons.

N: 1s2 2s2 2p3

O: 1s2 2s2 2p4

1s 2s 2p

1s 2s 2p

1s 2s 2p

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F: 1s2 2s2 2p5

Ne: 1s2 2s2 2p6

With neon, the orbital with n =1 and n = 2 are now completely filled.

Na: 1s2 2s2 2p63s1 can be abbreviate as Na : [Ne] 3s1

Write the symbol of the noble gas before the element

Then the rest of the electrons.

Mg: [Ne] 3s2

Al: [Ne] 3s2 3p1

1s 2s 2p

1s 2s 2p

Ne

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At this point it is useful to introduce the following concepts: � Valence electrons- the electrons in the outermost principle

quantum level of an atom (not d).� Core electrons- the inner electrons� Hund’s Rule- The lowest energy configuration for an atom

is the one have the maximum number of unpaired electrons in the orbital.

Example:

element valence electrons core electrons

O 6 2

N 5 2

Ne 8 2

Mg 2 10

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K: 1s2 2s2 2p63s13p64s1 or [Ar] 4s1 (valence electrons = 1)

Ca: 1s2 2s2 2p63s13p64s2 or [Ar] 4s2 (valence electrons = 1)

Sc: [Ar] 4s23d1 Ti: [Ar] 4s23d2 V: [Ar] 4s23d3

Valence electrons: 3 4 5

The expected configuration for chromium is:

Cr: [Ar] 4s23d4 however, the observed configuration is:

Cr: [Ar] 4s13d5 both 4s and 3d half-filled

Also the expected configuration for Cu is: Cu : [Ar] 4s23d9

The observed configuration is: Cu : [Ar] 4s13d10

4s is half-filled, 3d is filled

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Fill from the bottom up following the Fill from the bottom up following the arrowsarrows

1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

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The periodic table can be used as a guide for electron configurations.• the groups label (1A-8A) called the main-groups or representative elements (no. of the group = valence electrons• the groups label (1B-8B) called the Transion elements• the (n+1)s orbital is always fill before nd orbitals.• The period number is the value of n.• Groups 1A and 2A (1 & 2) have the s-orbital filled.• Groups 3A - 8A (13 - 18) have the p-orbital filled.

• Groups 3B - 2B (3 - 12) have the d-orbital filled.• The lanthanides and actinides have the f-orbital filled.

Electron Configurations and the Periodic TableElectron Configurations and the Periodic Table

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� Elements in the same column have the same electron configuration.

� Put in columns because of similar properties.� Similar properties because of electron configuration.� Noble gases have filled energy levels.� Transition metals are filling the d orbitals

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Periodic TrendsPeriodic Trends in atomic propertiesin atomic properties� Ionization energy (I.E.):

Ionization energy the energy required to remove an electron form a gaseous atom or ion in its ground state.

X(g) → X+(g) + e

We will consider the energy required to remove several electrons from Al in the ground state.

Al(g) → Al+(g) + e I1 = 580 kJ/mol

Al+(g) → Al+2

(g) + e I2 = 1815 kJ/mol

Al+2(g) → Al+3

(g) + e I3 = 2740 kJ/mol

Al+3(g) → Al+4

(g) + e I4 = 11600 kJ/mol

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Several points can be illustrated from these results:Highest energy electron (the one bound least tightly) that is removed first.The first ionization energy I1 is the energy required to remove the first electron (highest-energy electron) The value of I1 is considerably smaller than the value of I2 (second ionization energy). The primary factor is simply charge, electron is removed from +1 ion (Al+) .The increase in positive charge bind the electron more firmly, and the ionization energy increases. The same trend shows up in I3 and I4, where the electron is removed from Al+2 and Al+3 ions respectively.The increase in I.E. from I1 to I2 occur also because the first electron is removed from 3p orbital that is higher in energy than 3s orbital from which the second electron is removed.The largest jump in I.E. by far occur in going from the I3 and I4 because Al+3 has the configuration (1s2 2s2 2p6), the core electrons are bound much more tightly than valence electrons.

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Symbol I1 I2 I3HHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

1181014840 3569 4619 4577 5301 6045

6276

In the following table, the ionization energies for all the period 3 are given. Note the large jump in energy in each case in going

from removal of valence electrons to removal of core electrons

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Firs

t Ion

izat

ion

en

erg

y

Atomic number

H

He

Li

Be

B

C

N

The values of the first I.E. for the elements are shown in the following figure:

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Note that: As you go down a group first I.E. decreases because of electron being removed are, on average, farther from the nucleus. As n increases, the size of the orbital increases, and the electron is easier to remove. As you go across a period from left to right, first I.E. increases because - Same shielding (same principle quantum level) .

- Increasing nuclear charge (electrons are strongly bound)

There are some trends in I.E. in going across period. For example, trends occur from Be to B and from N to O. it can be explain in term of electron repulsion. Half-filled and filled orbitals are harder to remove electrons from

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The ionization energies for the representative elements are summarized in the following figure

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� Atomic Size or radius:The size of the orbital cannot be specified exactly (The electron cloud

doesn’t have a definite edge), neither can the size of an atom.

We can make some arbitrary choice to obtain values for atomic radii.

These values can be obtain by measuring the distance between atoms in chemical compounds.

For example, in Br2 molecule, the distance between the two nuclei is 228 pm. The Br atomic radius is assumed to be half this distance, or 114 pm, as shown in the following figure.

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Atomic SizeAtomic Size

� Atomic Radius = half the distance between Atomic Radius = half the distance between two nuclei of a diatomic moleculetwo nuclei of a diatomic molecule

}Radius

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� Influenced by two factorsInfluenced by two factors� ShieldingShielding� More shielding is further awayMore shielding is further away� Charge on nucleusCharge on nucleus� More charge pulls electrons in More charge pulls electrons in

closercloser

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As we go down a As we go down a group each atom group each atom has another has another energy level.energy level.

So the atoms get So the atoms get biggerbigger

HLi

Na

K

Rb

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� As you go across a period the radius As you go across a period the radius gets smaller.gets smaller.

� Same energy levelSame energy level� More nuclear chargeMore nuclear charge� Outermost electrons are closerOutermost electrons are closer

Na Mg Al Si P S Cl Ar

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OverallOverall

Atomic Number

Ato

mic

Ra

diu

s (n

m)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

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Electron Affinity:

The energy change associated with adding an electron to a gaseous atom.

X(g) + e X-(g)

High electron affinity gives you energy- � exothermic� More negative

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In general, electron affinity becomes more exothermic as you go from left to right across a row.

Increase (more - ) from left to right (greater nuclear charge).

� Decrease as we go down a group (More shielding)

65© 2009, Prentice-Hall, Inc.

Trends in Electron AffinityTrends in Electron Affinity

There are There are again, again, however, two however, two discontinuitiediscontinuities in this trend.s in this trend.

66© 2009, Prentice-Hall, Inc.

Trends in Electron AffinityTrends in Electron Affinity

� The first occurs The first occurs between Groups IA between Groups IA and IIA.and IIA.– The added electron The added electron

must go in a must go in a pp--orbital, not an orbital, not an ss--orbital.orbital.

– The electron is The electron is farther from nucleus farther from nucleus and feels repulsion and feels repulsion from the from the ss-electrons.-electrons.

67© 2009, Prentice-Hall, Inc.

Trends in Electron AffinityTrends in Electron Affinity

� The second occurs The second occurs between Groups IVA between Groups IVA and VA.and VA.– Group VA has no Group VA has no

empty orbitals.empty orbitals.– The extra electron The extra electron

must go into an must go into an already occupied already occupied orbital, creating orbital, creating repulsion.repulsion.

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Ionic SizeIonic Size� Cations form by losing electronsCations form by losing electrons� Cations are smaller than the atom Cations are smaller than the atom

they come fromthey come from� Metals form cationsMetals form cations� Cations of representative elements Cations of representative elements

have noble gas configuration.have noble gas configuration.

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Ionic sizeIonic size� Anions form by gaining electronsAnions form by gaining electrons� Anions are bigger than the atom they Anions are bigger than the atom they

come fromcome from� Nonmetals form anionsNonmetals form anions� Anions of representative elements Anions of representative elements

have noble gas configuration.have noble gas configuration.

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Configuration of IonsConfiguration of Ions� Ions always have noble gas Ions always have noble gas

configurationconfiguration

� Na is 1sNa is 1s222s2s222p2p663s3s11

� Forms a 1+ ion - 1sForms a 1+ ion - 1s222s2s222p2p66 � Same configuration as neonSame configuration as neon� Metals form ions with the Metals form ions with the

configuration of the noble gas before configuration of the noble gas before them - they lose electronsthem - they lose electrons

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Configuration of IonsConfiguration of Ions� Non-metals form ions by gaining Non-metals form ions by gaining

electrons to achieve noble gas electrons to achieve noble gas configuration.configuration.

� They end up with the configuration They end up with the configuration of the noble gas after them.of the noble gas after them.

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Group trendsGroup trends� Adding energy levelAdding energy level� Ions get bigger as Ions get bigger as

you go downyou go downLi+1

Na+1

K+1

Rb+1

Cs+1

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Periodic TrendsPeriodic Trends� Across the period nuclear charge Across the period nuclear charge

increases so they get smaller.increases so they get smaller.� Energy level changes between Energy level changes between

anions and cationsanions and cations

Li+1

Be+2

B+3

C+4

N-3O-2 F-1

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Size of Isoelectronic ionsSize of Isoelectronic ions� Iso - sameIso - same� Iso electronic ions have the same # Iso electronic ions have the same #

of electronsof electrons

� AlAl+3+3 Mg Mg+2 +2 NaNa+1 +1 Ne FNe F-1 -1 OO-2 -2 and Nand N-3-3 � all have 10 electronsall have 10 electrons

� all have the configuration 1sall have the configuration 1s222s2s222p2p66

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Size of Isoelectronic ionsSize of Isoelectronic ions� Positive ions have more protons so Positive ions have more protons so

they are smallerthey are smaller

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

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ElectronegativityElectronegativity

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ElectronegativityElectronegativity� The tendency for an atom to attract The tendency for an atom to attract

electrons to itself when it is electrons to itself when it is chemically combined with another chemically combined with another element.element.

� How “greedy”How “greedy”� Big electronegativity means it pulls Big electronegativity means it pulls

the electron toward itself.the electron toward itself.� Atoms with large negative electron Atoms with large negative electron

affinity have larger electronegativity.affinity have larger electronegativity.

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Group TrendGroup Trend� The further down a group more The further down a group more

shieldingshielding� Less attracted (ZLess attracted (Zeffeff))

� Low electronegativity.Low electronegativity.

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Periodic TrendPeriodic Trend� Metals are at the left endMetals are at the left end� Low ionization energy- low effective Low ionization energy- low effective

nuclear chargenuclear charge� Low electronegativityLow electronegativity� At the right end are the nonmetalsAt the right end are the nonmetals� More negative electron affinityMore negative electron affinity� High electronegativityHigh electronegativity� Except noble gasesExcept noble gases

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Ionization energy, electronegativity

Electron affinity INCREASE

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Atomic size increases,

Ionic size increases

82

Parts of the Periodic TableParts of the Periodic Table

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The information it hidesThe information it hides� Know the special groupsKnow the special groups� It is the number and type of valence It is the number and type of valence

electrons that determine an atom’s electrons that determine an atom’s chemistry.chemistry.

� You can get the electron configuration You can get the electron configuration from it.from it.

� Metals lose electrons have the lowest IEMetals lose electrons have the lowest IE� Non metals- gain electrons most Non metals- gain electrons most

negative electron affinitiesnegative electron affinities

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The Alkali MetalsThe Alkali Metals� Doesn’t include hydrogen- it behaves Doesn’t include hydrogen- it behaves

as a non-metalas a non-metal� decrease in IEdecrease in IE� increase in radiusincrease in radius� Decrease in densityDecrease in density� decrease in melting pointdecrease in melting point� Behave as reducing agentsBehave as reducing agents

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Reducing abilityReducing ability� Lower IE< better reducing agentsLower IE< better reducing agents� Cs>Rb>K>Na>LiCs>Rb>K>Na>Li� works for solids, but not in aqueous works for solids, but not in aqueous

solutions.solutions.� In solution Li>K>NaIn solution Li>K>Na� Why?Why?� It’s the water -there is an energy It’s the water -there is an energy

change associated with dissolvingchange associated with dissolving

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Hydration EnergyHydration Energy� LiLi++(g) (g) → Li→ Li++(aq)(aq) is exothermic is exothermic

� for Lifor Li++ -510 kJ/mol -510 kJ/mol

� for Nafor Na+ + -402 kJ/mol-402 kJ/mol

� for Kfor K++ -314 kJ/mol -314 kJ/mol� Li is so big because of it has a high Li is so big because of it has a high

charge density, a lot of charge on a charge density, a lot of charge on a small atom.small atom.

� Li loses its electron more easily Li loses its electron more easily because of this in aqueous solutionsbecause of this in aqueous solutions

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The reaction with waterThe reaction with water� Na and K react explosively with waterNa and K react explosively with water� Li doesn’t.Li doesn’t.� Even though the reaction of Li has a Even though the reaction of Li has a

more negative more negative ∆∆H than that of Na and H than that of Na and KK

� Na and K meltNa and K melt� ∆∆H does not tell you speed of reactionH does not tell you speed of reaction� More in Chapter 12.More in Chapter 12.