Electrons in Atoms Chapter 3 Standard 2.3 Cambridge Standards 2.3 Electrons: energy levels, atomic orbitals, ionization energy, electron affinity 2.3

Electrons in AtomsElectrons in Atoms

Chapter 3Standard 2.3

Cambridge StandardsCambridge Standards

• 2.3 Electrons: energy 2.3 Electrons: energy levels, atomic orbitals, levels, atomic orbitals, ionization energy, electron ionization energy, electron affinityaffinity

Leaning Outcomes• Describe the number and relative energies of the s,p,and d

orbitals for the principal quantum numbers 1,2 and 3 and also the 4s and 4p orbitals

• Describe the shapes of s and p orbitals• State the electronic configuration of atoms and ions given

the proton number and charge, using the convention 1s22s22p6, etc

• Explain and use the term ionisation energy, and the factors influencing the ionisation energies of elements

• Use isonisation energy data to:– Explain the trends across a period and down a group of the

periodic table– Deduce the electronic configurations of elements

• Interpret successive ionisation energy data of an element in terms of the position of that element within the periodic table

Electron OrbitalsElectron Orbitals

Cartoon courtesy of lab-initio.com

• The arrangement of electrons within an atom determines how that atom interacts with other atoms

• Electrons are organized into different energy levels

What is an atomic orbital?

Orbitals and orbits•When a planet moves around the sun, you can plot a definite path for it which is called an orbit. A simple view of the atom looks similar and you may have pictured the electrons as orbiting around the nucleus. The truth is different, and electrons in fact inhabit regions of space known as orbitals.•Orbits and orbitals sound similar, but they have quite different meanings. It is essential that you understand the difference between them.

• Electrons are arranged in orbitals or energy levels or principal quantum shells

• Symbol n• Lowest energy level, n=1,

is closest to nucleus• Arrangement of electrons

in atoms is called it electronic structure or electronic configuration

p+

no

e-

Quantum Theory

Energy levels, n

n = 1

n = 2

n = 3

n = 4

Electron Energy Level (Shell)Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. “n” is also known as the Principle Quantum Principle Quantum numbernumberNumber of electrons that can fit in a shell: 2n2

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

An orbital is a region within an energy level where there is a probability of finding an electron.

Electron OrbitalsElectron Orbitals

The impossibility of drawing orbits for electrons

• To plot a path for something you need to know exactly where the object is and be able to work out exactly where it's going to be an instant later. You can't do this for electrons.

• The Heisenberg Uncertainty Principle says - loosely - that you can't know with certainty both where an electron is and where it's going next. (What it actually says is that it is impossible to define with absolute precision, at the same time, both the position and the momentum of an electron.)

• That makes it impossible to plot an orbit for an electron around a nucleus. Is this a big problem? No. If something is

impossible, you have to accept it and find a way around it.

Hydrogen's electron

the 1s orbital

• Suppose I asked everyone to draw the hydrogen atom and place it’s one electron in position. Then somehow I superimposed everyone’s drawings on top of each other. I doubt any of the drawn electrons would line up. So who put the electron in the correct place?

• Everyone. We would end up with a sort of 3D map of the place that the electron is likely to be found.

• In the hydrogen case, the electron can be found anywhere within a spherical space surrounding the nucleus. The diagram shows a cross-section through this spherical space.

• Most of the time,the electron will be found within a fairly easily defined region of space quite close to the nucleus. Such a region of space is called an orbital. You can think of an orbital as being the region of space in which the electron lives.

• If you wanted to be absolutely 100% sure of where the electron is, you would have to draw an orbital the size of the Universe!

Each orbital has a name

• The orbital occupied by the hydrogen electron is called a 1s orbital. The "1" represents the fact that the orbital is in the energy level closest to the nucleus.

• The "s" tells you about the shape of the orbital.

• s orbitals are spherically symmetric around the nucleus.

2s orbital• This is similar to a 1s orbital except

that the region where there is the greatest chance of finding the electron is further from the nucleus - this is an orbital at the second energy level.

• "Electron density" is another way of talking about how likely you are to find an electron at a particular place

• 3s, 4s (etc) orbitals get progressively further from the nucleus.

• The nearer the nucleus the electrons get, the lower their energy.

p orbitals

• Not all electrons inhabit s orbitals (in fact, very few electrons live in s orbitals). At the first energy level, the only orbital available to electrons is the 1s orbital, but at the second level, as well as a 2s orbital, there are also orbitals called 2p orbitals.

• At any one energy level it is possible to have three absolutely equivalent p orbitals pointing mutually at right angles to each other. These are arbitrarily given the symbols px, py and pz.

• This is simply for convenience - what you might think of as the x, y or z direction changes constantly as the atom tumbles in space.

p Orbitalshttp://youtu.be/VfBcfYR1VQo

http://youtu.be/VfBcfYR1VQo

http://youtu.be/VfBcfYR1VQo

• The p orbitals at the second energy level are called 2px, 2py and 2pz. There are similar orbitals at subsequent levels - 3px, 3py, 3pz, 4px, 4py, 4pz and so on.

• All levels except for the first level have p orbitals. At the higher levels the lobes get more elongated, with the most likely place to find the electron more distant from the nucleus.

• Start at the third energy level • 5 different shapes• Each can hold 2 electrons

http://youtu.be/K-jNgq16jEY

http://youtu.be/K-jNgq16jEY

d and f orbitals• Here is where it gets really weird• At the third level, there is a set of

five d orbitals (with complicated shapes and names) as well as the 3s and 3p orbitals (3px, 3py, 3pz). At the third level there are a total of nine orbitals altogether.

• At the fourth level, as well the 4s and 4p and 4d orbitals there are an additional seven f orbitals - 16 orbitals in all. s, p, d and f orbitals are then available at all higher energy levels as well.

• Start at the fourth energy level• Have seven different shapes• 2 electrons per shape

1

2

3

4

5

6

7

6

7

1A

2A

3B 4B 5B 6B 7B 8B 8B 8B 1B 2B

3A 4A 5A 6A 7A

8Agroup # = # valence (outside) e-

d p

f

sRow

=# shells

Summary

s

p

d

f

# of shapes

Max electrons

Starts at energy level

1 2 1

3 6 2

5 10 3

7 14 4

Fitting electrons into orbitals• You can think of an atom as a very bizarre hotel (like an

inverted pyramid- remember the Wayside School books?) - with the nucleus living on the ground floor, and then various rooms (orbitals) on the higher floors occupied by the electrons.

• On the first floor there is only 1 room (the 1s orbital)• On the second floor there are 4 rooms (the 2s, 2px, 2pyand

2pz orbitals)

• On the third floor there are 9 rooms (one 3s orbital, three 3p orbitals and five 3d orbitals); and so on.

• But the rooms aren't very big . . . Each orbital can only hold 2 electrons.

Energy sublevels are contained within the principal energy levels.

"Electrons-in-boxes"

• Orbitals can be represented as boxes with the electrons in them shown as arrows.

• An up-arrow and a down-arrow are used to show that the electrons spin is different directions

• A 1s orbital holding 2 electrons would be drawn as

• A 1s orbital holding 2 electrons would be drawn as shown on the right, but it can be written even more quickly as 1s2. This is read as "one-s-two" not as "one-s-squared".

• Don’t confuse the two numbers in this notation

In an orbital (box) diagram a box represents each orbital within subshells, and arrows represent electrons. The arrows’ directions represent electron spins; opposing spins are paired.

Electron Configurations

N:

Hund's rule• This filling of orbitals singly

where possible is known as Hund's rule. It only applies where the orbitals have exactly the same energies (as with p orbitals, for example), and helps to minimize the repulsions between electrons and so makes the atom more stable.

• The ground state electron configuration for nitrogen would be 1s22s22p3.

• A drawback to this method of showing the electron configuration is that it does not tell us how the three 2p electrons are distributed among the three 2p orbitals.

• We can show this by using an orbital diagram in which boxes are used to indicate orbitals within subshells and arrows to represent electrons in these orbitals. The direction of the arrows represent the directions of the electron spins. The orbital diagram for nitrogen is

Hund’s Rule(of maximum multiplicity)

“For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

NOT:

Elements of the Second Period• In the first-period elements, hydrogen and helium,

electrons occupy the orbital of the first main energy level.

• According to the Aufbau principle, after the 1s orbital is filled, the next electron occupies the s sublevel in the second main energy level.

Orbital Diagrams

• Aufbau Rule has us fill low energy orbitals first

• Hunds Rule tell us that if the energies are equivalent (like in the p orbital) give each an electron to minimize electron repulsion

Electron Configurations and Orbital Diagrams

Atom Configuration Diagram

Hydrogen 1s1 1s

Helium 1s2 1s

Beryllium 1s22s2 1s 2s

Carbon 1s22s22p2 1s 2s 2p

Oxygen 1s22s22p4 1s 2s 2p

↑↓

↑↓

↑

↑↓

↑↓ ↑↓ ↑ ↑

↑↓ ↑↓ ↑↓ ↑ ↑

Orbital Notation for Three Noble Gases

Orbital Notation for Argon and Potassium

Orbitals being filled for elements in various parts of the periodic table.

• d electrons are almost always described as, for example, d5 or d8- and not written as separate orbitals.

• Remember that there are five d orbitals, and that the electrons will inhabit them singly as far as possible.

• Up to 5 electrons will occupy orbitals on their own. After that they will have to pair up.

d5 means

d8 means

Electron Configurations• The electron configuration of an atom is a shorthand

method of writing the location of electrons by sublevel.

• The sublevel is written followed by a superscript with the number of electrons in the sublevel.

– If the 2p sublevel contains 2 electrons, it is written 2p2

• An electron configuration describes the distribution of electrons among the various orbitals in the atom.

• Electron configuration is represented in two ways.

Electron Configurations

The spdf notation uses numbers to designate a principal shell and letters (s, p, d, f) to identify a subshell; a superscript indicates the number of electrons in a designated subshell.

the Aufbau Principle

• Aufbau is a German word meaning building up or construction. We imagine that as you go from one atom to the next in the Periodic Table, you can work out the electronic structure of the next atom by fitting an extra electron into the next available orbital.

• Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones. Where there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible.

• Aufbau comes from the German word "Aufbauen" which means "to build". In essence when writing electron configurations we are building up electron orbitals as we proceed from atom to atom. As we write the electron configuration for an atom, we will fill the orbitals in order of increasing atomic number.

Aufbau Principle• Electrons occupy the orbitals of lowest energy first• An orbital can hold at most 2 electrons.• The number of electrons in a sublevel is indicated by

adding a superscript to the sublevel designation

Aufbau PrincipleAn electron occupies

the lowest-energy orbital that can receive

it (Always start with n = 1 and work your way up)

Hydrogen = 1 electron = 1s1

Helium = 2 electrons = 1s2

Lithium = 3 electrons = 1s2 2s1

Nitrogen = 7 electrons = 1s2 2s2 2p3

Order of Subshell Energies

• Follow the arrows from the top: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

• Subshells that are far from the nucleus may exhibit exceptions to the filling order.

Electron ConfigurationsElectron Configurations

1s2

1s

2s

2p 2p 2p

The ground state electron configuration of helium

Electron ConfigurationsElectron ConfigurationsThe Aufbau principle states that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals.

1s22s1

1s

2s

2p 2p 2p

The ground state electron configuration of Li

The 1s orbital can only accommodate 2 electrons (Pauli exclusion principle)

The third electron must go in the next available orbital with the lowest possible energy.

Li has a total of 3 electrons


The Aufbau principle states that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals.

1s

2s

2p 2p 2p 1s22s2

The ground state electron configuration of Be

Be has a total of 4 electrons

Electron ConfigurationsElectron ConfigurationsThe Aufbau principle states that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals.

1s

2s

2p 2p 2p

The ground state electron configuration of B

1s22s22p1

B has a total of 5 electrons


According to Hund’s rule, the most stable arrangement of electrons is the one in which the number of electrons with the same spin is maximized.

1s22s22p2

1s

2s

2p 2p 2p

The ground state electron configuration of C

The 2p orbitals are of equal energy.

Put 1 electron in each before pairing (Hund’s rule).

C has a total of 6 electrons



1s22s22p3

1s

2s

2p 2p 2p

The ground state electron configuration of N

The 2p orbitals are of equal energy.

Put 1 electron in each before pairing (Hund’s rule).

N has a total of 7 electrons



1s22s22p4

1s

2s

2p 2p 2p

The ground state electron configuration of O

O has a total of 8 electrons

Once all the 2p orbitals are singly occupied, additional electrons will have to pair with those already in the orbitals.



1s22s22p5

1s

2s

2p 2p 2p

The ground state electron configuration of F

F has a total of 9 electrons



1s22s22p6

1s

2s

2p 2p 2p

The ground state electron configuration of Ne

Ne has a total of 10 electrons

Electron configurations


General rules for writing electron configurations:

1) Electrons will reside in the available orbitals of the lowest possible energy.

2) Each orbital can accommodate a maximum of two electrons.

3) Electrons will not pair in degenerate orbitals if an empty orbital is available.

4) Orbitals will fill in the order indicated in the figure.

• Notice in what follows that all the 3-level orbitals are written together - with the 4s electrons written at the end of the electronic structure

Sc 1s22s22p63s23p63d14s2

Ti 1s22s22p63s23p63d24s2

V 1s22s22p63s23p63d34s2

Cr 1s22s22p63s23p63d54s1

Whoops! Chromium breaks the sequence. In chromium, the electrons in the 3d and 4s orbitals rearrange so that there is one electron in each orbital. It would be convenient if the sequence was tidy - but it's not!

Survey of the Periodic Table

Rule: Sublevels are most stable when they are either half or completely

filled. Electrons will shift to different energy levels to

accommodate this stability whenever possible.

another awkward one!

Mn 1s22s22p63s23p63d54s2

Fe 1s22s22p63s23p63d64s2

Co 1s22s22p63s23p63d74s2

Ni 1s22s22p63s23p63d84s2

Cu 1s22s22p63s23p63d104s1

Zn 1s22s22p63s23p63d104s2

And at zinc the process of filling the d orbitals is complete

There are several notable exceptions to the order of electron filling for some of the transition metals.

Chromium (Z = 24) is [Ar]4s13d5 and not [Ar]4s23d4 as expected.Copper (Z = 29) is [Ar]4s13d10 and not [Ar]4s23d9 as expected.

The reason for these anomalies is the slightly greater stability of d subshells that are either half-filled (d5) or completely filled (d10).

4s 3d 3d 3d 3d 3d[Ar]Cr

Greater stability with half-filled 3d subshell

Electron Configurations and the Periodic TableElectron Configurations and the Periodic Table

There are several notable exceptions to the order of electron filling for some of the transition metals.

Chromium (Z = 24) is [Ar]4s13d5 and not [Ar]4s23d4 as expected.Copper (Z = 29) is [Ar]4s13d10 and not [Ar]4s23d9 as expected.

The reason for these anomalies is the slightly greater stability of d subshells that are either half-filled (d5) or completely filled (d10).


4s 3d 3d 3d 3d 3d[Ar]Cu

Greater stability with filled 3d subshell

Chromium and Copper“expect the unexpected!”

• Chromium, Cr- atomic number 24 expected: 1s22s22p63s23p64s23d4

actual: 1s22s22p63s23p64s13d5

• Copper, Cu- atomic number 29 expected: 1s22s22p63s23p64s23d9

actual: 1s22s22p63s23p64s13d10

Some AnomaliesSome irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

Exceptions to the Aufbau PrincipleHalf-filled d subshell plus half-filled s subshell has slightly lower in energy than s2 d4.

Filled d subshell plus half-filled s subshell has slightly lower in energy than s2 d9.

More exceptions occur farther down the periodic table. They aren’t always predictable, because energy levels get closer together.

Order of Orbitals

Energy levels

• Notice that the s orbital always has a slightly lower energy than the p orbitals at the same energy level, so the s orbital always fills with electrons before the corresponding p orbitals

• The real oddity is the position of the 3d orbitals. They are at a slightly higher level than the 4s - and so it is the 4s orbital which you fill first, followed by all the 3d orbitals and then the 4p orbitals.

Electron Electron Filling Filling OrderOrder

Using the Periodic Table to Write Electron Configurations

The electron configuration of Si ends with 3s2 3p2

The electron configuration of Rh ends with 5s2 4d7

• Electrons ordinarily occupy orbitals of the lowest energy available.

• No two electrons in the same atom may have all four quantum numbers alike.

• Pauli exclusion principle: one atomic orbital can accommodate no more than two electrons, and these electrons must have opposing spins.

• Of a group of orbitals of identical energy, electrons enter empty orbitals whenever possible (Hund’s rule).

• Electrons in half-filled orbitals have parallel spins (same direction).

Rules for Electron Configurations

• The Aufbau principle describes a hypothetical “building-up” of an atom from the one that precedes it in atomic number.

(Z = 1) H 1s1

(Z = 2) He 1s2

(Z = 3) Li 1s2 2s1

• Noble-gas-core abbreviation: we can replace the portion that corresponds to the electron configuration of a noble gas with a bracketed chemical symbol. It’s easier to write …

(Z = 3) Li [He]2s1

(Z = 22) Ti [Ar]4s2 3d2

The Aufbau Principle

To get He, add one electron to H.

To get Li, add one electron to He.


The electron configurations of all elements except hydrogen and helium can be represented using a noble gas core.

The electron configuration of potassium (Z = 19) is 1s22s22p63s23p64s1.

Because 1s22s22p63s23p6 is the electron configuration of argon, we can simplify potassium’s to [Ar]4s1.

1s22s22p63s23p64s1

The ground state electron configuration of K:

[Ar] [Ar]4s1

1s22s22p63s23p64s1

Short cut for writing electron configurations

Noble Gas Configuration

Noble gas notation uses noble gas symbols in brackets to shorten inner electron configurations of other elements.

• The valence shell is the outermost occupied principal shell. The valence shell contains the valence electrons.

• For main group elements, the number of valence shell electrons is the same as the periodic table group number (2A elements: two valence electrons, etc.)The period number is the same as the principal quantum number n of the electrons in the valence shell.

• Electrons in inner shells are called core electrons.

Example: As [Ar] 4s2 3d104p3

Valence Electrons and Core Electrons

Five valence electrons, for which n = 4

28 core electrons

• To obtain the electron configuration of an anion by the aufbau process, we simply add the additional electrons to the valence shell of the neutral nonmetal atom.

• The number added usually completes the shell.• A nonmetal monatomic ion usually attains the

electron configuration of a noble gas atom.O2– : [Ne]Br– : [Kr]

Electron Configurations of Ions

• A metal atom loses electrons to form a cation.• Electrons are removed from the configuration of

the atom.• The first electrons lost are those of the highest

principal quantum number.• If there are two subshells with the same highest

principal quantum number, electrons are lost from the subshell with the higher l.


Atom IonF 1s2 2s22p5 F– 1s2 2s22p6 [Ne]S [Ne] 3s2 3p4 S2– [Ne] 3s2 3p6 [Ar]

Fe [Ar] 4s2 3d6 Fe2+ [Ar] 4s2 3d6 [Ar] 3d6

Ti [Ar] 4s2 3d2 Ti4+ [Ar] 4s2 3d2 [Ar]Sr [Kr] 5s2 Sr2+ [Kr] 5s2 [Kr]

What would be the configuration of Fe3+? Of Sn2+?

Valence electrons are lost first.


White Board

Write the electron configuration of the Co3+ ion in a noble-gas-core abbreviated spdf notation.

White Board

Write electron configurations for sulfur, using both the spdf notation and an orbital diagram.

1s22s22p63s23p4 (2.8.6)

[Ne]3s2 3p4

ResonanceResonanceResonance is invoked when more than one valid Lewis structure can be written for a particular molecule.

The actual structure is an average of the resonance structures.

H

H

H

H

H

H

H

H

H

H

H

H

Benzene, C6H6

The bond lengths in the ring are identical, and between those of single and double bonds.

Resonance Bond Length and Bond EnergyResonance Bond Length and Bond Energy

Resonance bonds are shorter and stronger than single bonds.

Resonance bonds are longer and weaker than double bonds.

H

H

H

H

H

H

H

H

H

H

H

H

Resonance in Ozone, OResonance in Ozone, O33

Neither structure is correct.

O O O

O O O

Oxygen bond lengths are identical, and intermediate to single and double bonds

Resonance in a carbonate ion:

Resonance in an acetate ion:

Resonance in Polyatomic IonsResonance in Polyatomic Ions

Ionization Energy• Ionization energies (IE) have to do with the making of

ions. Ions are atoms which have gained or lost electrons. • The ionization energy is the amount of energy it takes to

detach one electron from a neutral atom.

Some elements actually have several ionization energies. When this is the case, we refer to them as the:

• "first ionization energy" or IE1,

• "second ionization energy" or IE2

• and so on.

• Ionization energy is the energy required to remove an electron from a ground-state gaseous atom.

• IE is usually expressed in kJ per mole of atoms.Mg(g) Mg+(g) + e– ΔH = I1

Mg+(g) Mg2+(g) + e– ΔH = I2

Mg2+(g) Mg3+(g) + e– ΔH = I3

Ionization Energy

Ionization Energy• ionization energy, IE1 - energy required to

remove an electron from a neutral atom:

• Depends on attraction between nucleus and outer electron

• ion- atom or group of atoms that has a positive or negative charge

• ionization- process that results in the formation of an ion

A + ionization energy A+ + e-

IONISATION ENERGY

• The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

• It is the energy needed to carry out this change per mole of X

Things to notice about the equation

• The state symbols (g) are essential. When you are talking about ionisation energies, everything must be present in the gas stategas state.

• Ionization energies are measured in kJ mol1 (kilojoules per molekilojoules per mole).

• They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).

Ionization EnergyIonization EnergyIonization EnergyIonization Energy

IE = energy required to remove an electron IE = energy required to remove an electron from an atom in the gas phase.from an atom in the gas phase.

Mg Mg (g) (g) + 738 kJ + 738 kJ Mg Mg++ (g) (g) + e-+ e-

Mg (g) + Mg (g) + 735735 kJ kJ MgMg++ (g) + e- (g) + e-MgMg+ + (g) + (g) + 14511451 kJ kJ MgMg2+2+ (g) + e- (g) + e-

MgMg2+2+ (g) (g) + + 77337733 kJ kJ MgMg3+3+ (g) (g) + e-+ e-

Energy cost is very high to dip into a shell of lower n. Energy cost is very high to dip into a shell of lower n.

Ionization EnergyIonization EnergyIonization EnergyIonization Energy

Ionization Energy

Ionization energies in kJ/mol1 2 3 4 5 6 7 8

H 1312

He 2372 5250

Li 520 7297 11810

Be 899 1757 14845 21000

B 800 2426 3659 25020 32821

C 1086 2352 4619 6221 37820 47260

N 1402 2855 4576 7473 9442 53250 64340

O 1314 3388 5296 7467 10987 13320 71320 84070

F 1680 3375 6045 8408 11020 15160 17860 92010

Ne 2080 3963 6130 9361 12180 15240

Na 496 4563 6913 9541 13350 16600 20113 25666

Mg 737 1450 7731 10545 13627 17995 21700 25662

• All elements have a first ionization energy - even atoms which don't form positive ions in test tubes.

• The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.

• Remember that helium is a noble gas.

• First ionisation energy shows periodicityperiodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.

• These variations in first ionization energy can all be explained in terms of the structures of the atoms involved.

Ionization Energy Trends• Ionization energy decreases down a group

due to shielding effect • shielding effectshielding effect- - inner core electrons shield

outermost (valence) electrons from nucleus positive charge, reducing the attractive forces

• ionization energy increases across a period increases across a period due to increased nuclear charge and lack of shielding effect

Trends in Ionization Energy

• The energy required to remove the first electron from the outer shell is called the first ionization energy.

• The second ionization energy is the energy required to remove the second electron.– Always greater than first IE.

• The third IE is the energy required to remove a third electron.– Greater than 1st or 2nd IE.

Ionization Energy - Group trends

• As you go down a group, the ionisation energy decreases because...–The electron is further away from

the attraction of the nucleus, and and thus easier to remove the outermost one.

–There is more shielding.

Ionization Energy - Period trends

Going across a period, ionisation energy generally increases from left to right because…..

the atomic radius decreases, that is, the atom is smaller. The outer electrons are closer to the nucleus and more strongly attracted to the centre. Therefore, it becomes more difficult to remove the outermost electron.

Factors affecting the size of ionisation energy

• Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.

• The charge on the nucleus• The distance of the electrons from the nucleus• The number of electrons between the outer electrons

and the nucleus• Whether the electron is on its own or in an orbital or

paired with another electron

The charge on the nucleus

• The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.

+1 +2 +3

WHAT AFFECTS IONISATION ENERGY?WHAT AFFECTS IONISATION ENERGY?

The value of the 1st Ionisation Energy depends on the electronic structure

Hydrogen Helium Lithium

The value for helium is higher than that for hydrogen because there are now two protons in the nucleus. The nuclear charge is greater so the pull on the outer electrons is larger. More energy will be needed to pull an electron out of the atom.

519 kJ mol-1

1310 kJ mol-1 2370 kJ mol-1

The distance of the electron from the nucleus

• Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away

• Atomic radius

The number of electrons between the outer

electrons and the nucleus• Consider a sodium atom, with the electronic structure

2,8,1. • If the outer electron looks in towards the nucleus, it

doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the center.

• This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.

Shielding• The shielding effect is

the decrease in attraction between an electron and the nucleus in any atom.

Whether the electron is on its own in an orbital or paired with another electron

• Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.

Explaining the pattern in the first

few elements• Hydrogen has an electronic structure of 1s1. It

is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted.

• There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1).

• Helium has a structure 1s2. • The electron is being removed from the same

orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.

• Lithium is 1s22s1. • Its outer electron is in the second energy level, much

more distant from the nucleus. • You might argue that that would be offset by the

additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.

• You can think of the electron as feeling a net 1+ pull from the center (3 protons offset by the two 1s2 electrons).

• If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.

The patterns in periods 2 and 3• The first thing to realize is that the patterns in

the two periods are identical - the difference being that the ionisation energies in period 3 are all lower than those in period 2.

• In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons.

• The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies.

• In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level and are screened by the 1s22s22p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge.

Why the drop between Be-B?• You might expect the boron value to be more than the beryllium

value because of the extra proton. Offsetting that is the fact that boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is further from the nucleus. This has two effects.

1. The increased distance results in a reduced attraction and so a reduced ionisation energy.

2. The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.

Be 1s22s2 1st I.E. = 900 kJmol-1

B 1s22s22px1 1st I.E. = 799 kJmol-1

Why the drop between N and O?• Once again, you might expect the ionisation energy

of the group 6 element to be higher than that of group 5 because of the extra proton. What is offsetting it this time?

• N 1s22s22px12py

12pz1 1st I.E. = 1400 kJmol-1

• O 1s22s22px22py

12pz1 1st I.E. = 1310 kJmol-1

• The screening is identical (from the 1s2 and the 2s2 electrons), and the electron is being removed from an identical orbital.

• The difference is that in the oxygen case the electron being removed is one of the 2px

2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be.

• The drop in ionisation energy at sulfur is accounted for in the same way.

N 1s22s22px12py

12pz1 1st I.E. = 1400 kJmol-1

O 1s22s22px22py

12pz1 1st I.E. = 1310 kJmol-1

Selected Ionization Energies

Compare I2 to I1 for a 2A element, then for the corresponding 1A element.

Why is I2 for each 1A element so much greater than I1?

Why don’t we see the same trend for each 2A element? I2 > I1 … but only about twice as great …

Selected Ionization EnergiesGeneral trend in I1: An increase from left to right, but …

…I1 drops, moving from 2A to 3A.

The electron being removed is now a p electron (higher energy, easier to remove than an s).

I1 drops again between 5A and 6A.

Repulsion of the paired electron in 6A makes that electron easier to remove.

First Ionization EnergiesChange in trend occurs at 2A-3A and at 5A-6A for each period …

… but the change becomes smaller at higher energy levels.

How does ionization energy change down a group?

• The first ionization energy decreases as you move down a group.

• Why?– The size of the atom increases.– Electron is further from the nucleus.

Trends in ionisation energy

down a group

• As you go down a group in the Periodic Table ionisation energies generally fall

• Why is the sodium value less than that of lithium?

• There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening.

• Lithium's outer electron is in the second level, and only has the 1s2 electrons to screen it. The 2s1 electron feels the pull of 3 protons screened by 2 electrons - a net pull from the center of 1+.

• The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner electrons. The 3s1 electron also feels a net pull of 1+ from the center of the atom. In other words, the effect of the extra protons is compensated for by the effect of the extra screening electrons. The only factor left is the extra distance between the outer electron and the nucleus in sodium's case. That lowers the ionisation energy.

Trends in ionisation energy in a transition series

• Apart from zinc at the end, the other ionisation energies are all much the same.

• All of these elements have an electronic structure [Ar]3dn4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital.

• I1 < I2 < I3

– Removing an electron from a positive ion is more difficult than removing it from a neutral atom.

• A large jump in I occurs after valence electrons are completely removed (why?).

• I1 decreases from top to bottom on the periodic table.– n increases; valence electron is farther from nucleus.

• I1 generally increases from left to right, with exceptions.– Greater effective nuclear charge from left to right

holds electrons more tightly.

Ionization Energy Trends

Successive Ionisation Energies

• A measure of the energy required to remove each electron in turn.

• Mg(g) Mg+(g) + e- 1st I.E. =+738 kJ.mol-1

• Mg+(g) Mg2+(g) + e- 2nd I.E.= + 1451kJ.mol-1

• Mg2+(g) Mg3+(g) + e- 3rd I.E.= + 7733kJ.mol-1

• Mg3+(g) Mg4+(g) + e- 4th I.E.= + 10541kJ.mol-1

Electron Affinity

• What does the word ‘affinity’ mean?• Electron affinity reflects the ability of an atom to

accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom.

• Electron affinity is, essentially the opposite of the ionization energy: Instead of removing an electron from the element we add an electron to the element to create an anion.

Ionization energies measure the tendency of a neutral atom to resist the loss of electrons. It takes a considerable amount of energy, for example, to remove an electron from a neutral fluorine atom to form a positively charged ion. Why?

F(g) + 1681.0 kJ/mol F+(g) + e-

ΔE = 1681.0 kJ/mol

Ionization energies are always concerned with the formation of positive ions.

Electron affinities are the negative ion equivalent.

Electron affinityThe electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ionA fluorine atom in the gas phase, for example, gives off energy when it gains an electron to form a fluoride ion.

An isolated electron is brought from far away to undergo attraction to the nucleus which lowers its energy → Excess energy released.

9+2e-

7e-

e-

F(g) + e- → F-(g) + 328 kJ/mol ΔE=-328 kJ/mol

• Where ionization energy is always endothermic, electron affinity is usually exothermic (releases energy), but not always.

• A negative electron affinity means that energy is released when an electron is added to the atom; if energy is required, the electron affinity is positive.

• The more negative the E.A , the greater is the tendency of the atom to attract an electron.

• Halogens have the most negative electron affinities.

Note that the noble gases, alkali metals and alkali earth metals have E.A. close to zero - indicating that these groups

of elements do not particularly like to become anions. However, the nonmetals and especially the halogens are

highly negative and thus readily become anions.

Electron affinity (EA) is the energy change that occurs when an electron is added to a gaseous atom:

M(g) + e– M–(g) ΔH = EA1

• A negative electron affinity means that the process is exothermic.

• Nonmetals generally have more affinity for electrons than metals do. (Nonmetals like to form anions!)

• Electron affinity generally is more negative or less positive on the right and toward the top of the periodic table.

Electron Affinity

Selected Electron AffinitiesThe halogens have a greater affinity for electrons than do the alkali metals, as expected.

Electron Affinity• Electron affinity generally increases across periods.

• Electron affinity generally decreases down groups.– The larger an atom’s electron cloud is, the farther away its outer electrons are

from its nucleus.

Affinity for electron increases across a period Affinity for electron increases across a period (EA (EA becomes more negative)becomes more negative)..

Affinity decreases down a group Affinity decreases down a group (EA becomes (EA becomes lessless negative)negative). . Why???Why???

Atom EAAtom EA

FF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ

Atom EAAtom EA

FF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ

Trends in Size, IE, and EATrends in Size, IE, and EA• IE, and EA are the opposite of atomic radius

LargerLarger

LargerLarger

SizeAtomic Radii

Ionization energy

LargerLarger

Electron Affinity

videos• https://youtu.be/ywqg9PorTAw?list=PL166048DD75B05

C0D Periodic Table Trends: Ionization Energy

• https://youtu.be/uVWquWFjnCw First and Second Ionization Energy

• https://youtu.be/j0xQmCPXTvY Electron Affinity• https://youtu.be/rcKilE9CdaA crash course electrons• https://youtu.be/0tP6bV89log periodicity

https://youtu.be/ywqg9PorTAw?list=PL166048DD75B05C0D

https://youtu.be/ywqg9PorTAw?list=PL166048DD75B05C0D

https://youtu.be/uVWquWFjnCw

https://youtu.be/j0xQmCPXTvY

https://youtu.be/rcKilE9CdaA

https://youtu.be/0tP6bV89log

Certain physical and chemical properties recur at regular intervals, and/or vary in regular fashion, when the elements are arranged according to increasing atomic number.

Melting point, boiling point, hardness, density, physical state, and chemical reactivity are periodic properties.

We will examine several periodic properties that are readily explained using electron configurations.

Periodic Properties

Half the distance between the nuclei of two atoms is the atomic radius.Covalent radius: half the distance between the nuclei of two identical atoms joined in a molecule.Metallic radius: half the distance between the nuclei of adjacent atoms in a solid metal.

Periodic Properties: Atomic Radius

• Atomic radius increases from top to bottom within a group.

• The value of n increases, moving down the periodic table.

• The value of n relates to the distance of an electron from the nucleus.


• Atomic radius decreases from left to right within a period.• Why? The effective nuclear charge increases from left to

right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller.


Mg has a greater effective nuclear charge than Na, and is smaller than Na.

Atomic Radii of the Elements

Example 8.5With reference only to a periodic table, arrange each set of elements in order of increasing atomic radius:(a) Mg, S, Si (b) As, N, P (c) As, Sb, Se

The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion.

Ionic Radii

• Cations are smaller than the atoms from which they are formed; the value of n usually decreases. Also, there is less electron–electron repulsion.

Ionic Radii

• Anions are larger than the atoms from which they are formed.

• Effective nuclear charge is unchanged, but additional electron(s) increase electron–electron repulsion.

• Isoelectronic species have the same electron configuration; size decreases with effective nuclear charge.

Ionic Radii

SomeAtomic

andIonicRadii

** white board**

Identify in each case one element that has the following properties:1.The element with forms the largest cation 2.An element that floats on water and reacts with it3.An element in the s-block whose nitrate gives a brown gas on thermal decomposition

1.1. K or K+K or K+2.2. Na allow K Na allow K

or Lior Li3.3. Mg or Ca or Mg or Ca or

LiLi

Example 8.6Refer to a periodic table but not to Figure 8.14, and arrange the following species in the expected order of increasing radius:Ca2+, Fe3+, K+, S2–, Se2–

The ELECTRON: Wave – Particle DualityThe ELECTRON: Wave – Particle Duality

Graphic: www.lab-initio.com

The Dilemma of the Atom

• Electrons outside the nucleus are attracted to the protons in the nucleus

• Charged particles moving in curved paths lose energy

• What keeps the atom from collapsing?

• Electrons outside the nucleus are attracted to the protons in the nucleus

• Charged particles moving in curved paths lose energy

• What keeps the atom from collapsing?

Wave-Particle DualityJJ Thomson won the Nobel prize for describing the electron as a particle.

His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.

The electron is a particle! The electron is

an energy wave!

The Wave-like Electron

Louis deBroglie

The electron propagates through space as an energy wave. To

understand the atom, one must understand the behavior of

electromagnetic waves.

The electron propagates through space as an energy wave. To

understand the atom, one must understand the behavior of

electromagnetic waves.

c = c = c = speed of light, a constant (3.00 x 108 m/s)

= frequency, in units of hertz (hz, sec-1) = wavelength, in meters

Electromagnetic radiation propagates through space as a wave moving at the speed of light.

http://upload.wikimedia.org/wikipedia/commons/6/64/Wavelength.png

E = hE = h

EE = Energy, in units of Joules (kg·m= Energy, in units of Joules (kg·m22/s/s22))

hh = Planck’s constant (6.626 x 10-34 J·s)= Planck’s constant (6.626 x 10-34 J·s)

= frequency, in units of hertz (hz, sec= frequency, in units of hertz (hz, sec-1-1))

The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.

Long Wavelength

=Low Frequency

=Low ENERGY

Short Wavelength

=High

Frequency=

High ENERGY

Wavelength Table

Answering the Dilemma of the Atom

• Treat electrons as waves• As the electron moves toward the nucleus,

the wavelength shortens• Shorter wavelength = higher energy• Higher energy = greater distance from the

nucleus

• Treat electrons as waves• As the electron moves toward the nucleus,

the wavelength shortens• Shorter wavelength = higher energy• Higher energy = greater distance from the

nucleus

The Electromagnetic Spectrum

This produces bandsThis produces bandsof light with definite of light with definite wavelengths.wavelengths.

Electron transitionsinvolve jumps of definite amounts ofenergy.

…produces a “bright line” spectrum

Spectroscopic analysis of the hydrogen spectrum…

Flame Tests

strontium sodium lithium potassium copper

Many elements give off characteristic light which can Many elements give off characteristic light which can be used to help identify them.be used to help identify them.

Documents

Electrons in Atoms Chapter 3 Standard 2.3 Cambridge Standards 2.3 Electrons: energy levels, atomic orbitals, ionization energy, electron affinity 2.3