1 modern atomic theory (a.k.a. the electron chapter!) chapter 4

1Modern Atomic Theory

(a.k.a. the electron chapter!)

Chapter 4Chapter 4

2

Why are electrons important?• They are the “exterior” of the atom.

• Although the number of protons mainly determines the number of electrons, it is the electrons and how they are arranged that determine EVERY property of a substance!

• All bonding of elements into compounds is determined by the configuration (arrangement) of the atom’s electrons.

• CHEMISTRY IS ALL ABOUT ELECTRONS!

3

ELECTROMAGNETIC ELECTROMAGNETIC RADIATION – CH4.1RADIATION – CH4.1ELECTROMAGNETIC ELECTROMAGNETIC RADIATION – CH4.1RADIATION – CH4.1

4

• Why do we talk about electromagnetic radiation in this unit??????

• All electromagnetic radiation comes from disturbances in the electrons of atoms.

• By studying E.M. radiation, we can figure out how electrons are arranged in atoms.

5Electromagnetic radiation.

6

Electromagnetic Electromagnetic RadiationRadiation


• Most subatomic particles behave as Most subatomic particles behave as PARTICLES and obey the physics of PARTICLES and obey the physics of waves.waves.

7

wavelength Visible light

wavelength

Ultaviolet radiation

Amplitude

Node



8

• Waves have a frequency,Waves have a frequency, f f, units are “cycles , units are “cycles per sec” or Hertz. Hzper sec” or Hertz. Hz

denotes wavelength, units are denotes wavelength, units are meters, mmeters, m

• All electromagnetic radiation: All electromagnetic radiation: • • f f = c = c

where c = velocity of light = 3.00 x 10where c = velocity of light = 3.00 x 1088 m/sec m/sec• Now, rearrange this equation for equations for Now, rearrange this equation for equations for

f and C f and C



9Do these problems showing work using G.U.E.S.S.

• Remember n or nano = 10-9, Speed of light = 3X108 m/s1) Find the wavelength of a radio wave that has a frequency

of 93.1megahertz – the frequency of WXRT, Chicago’s finest rock.

2) Find the frequency of a 633nm laser beam.3) If a light beam has a wavelength of 500nm and a frquency

of 6 X 1014 Hz, find its speed.The following are NOT E.M. waves, you can’t use C=

3X108m/s4) If an ocean wave passes me at 0.1 waves every second

and there is 7 m between the waves, find its speed5) If a sound wave has a speed of 340m/s and a frequency of

200Hz, find its wavelength.6) If a sound wave has a speed of 340m/s and a wavelength

of 0.1cm, find its frequency7) If a by shaking a rope 3 times per second a wave is made

with a wavelength of 3m find its speed

10

Electromagnetic Electromagnetic SpectrumSpectrum

Electromagnetic Electromagnetic SpectrumSpectrum

Long wavelength --> small frequencyLong wavelength --> small frequency

Short wavelength --> high frequencyShort wavelength --> high frequency

increasing increasing frequencyfrequency

increasing increasing wavelengthwavelength

11

ElectroElectromagneticmagnetic SpectrumSpectrum

ElectroElectromagneticmagnetic SpectrumSpectrum

In increasing energy, RIn increasing energy, ROOYY GG BBIIVV

12

Quantum Theory – CH 4.2• By 1900, this wave theory of light

generally accepted. But some unexplained things remained..

• Radiant energy – why does color of an object change as the temperature changes?

• Why do certain elements emit only a particular color when heated or its gas has electricity passed through it?

13

Plank’s Theory• Explained radiant energy..

• “There is a fundamental restriction on the amount of energy that an object emits or absorbs”

• These amounts are called “quanta” (singular is quantum)

• E=hf

• energy = Plank’s Constant X frequency

14Photoelectric Effect• Einstein used Plank’s equation to

explain the photoelectric effect.• Electrons are released from metal when

exposed to light (used in solar cells etc.)• However, even very intense infra red

light or red light did not work whereas even very faint UV light works.

• Each “photon” particle of light has energy given by E=hf

• Compton : Light acts as both a wave and a particle

15

Atomic Line Emission Atomic Line Emission Spectra and Niels BohrSpectra and Niels BohrAtomic Line Emission Atomic Line Emission

Spectra and Niels BohrSpectra and Niels Bohr

Bohr’s greatest contribution Bohr’s greatest contribution to science was in building a to science was in building a simple model of the atom. It simple model of the atom. It was based on an was based on an understanding of theunderstanding of the LINE LINE EMISSION SPECTRAEMISSION SPECTRA of of excited atoms.excited atoms.

Niels BohrNiels Bohr

(1885-1962)(1885-1962)

16Line emission spectrum• When an atom is left alone, the

electrons all have certain energy levels. This “rest” state is called the GROUND STATE.

• When electrons receive energy from the environment (heat, electrical or kinetic), the electrons move to higher energy levels – EXCITED STATE.

• It is when these excited electrons fall back to the ground state that this energy is released as light.

• The bigger the fall, the higher the frequency.

17Making line emission spectra

18

Spectrum of White Spectrum of White LightLight

19Line Emission Spectra of Excited Line Emission Spectra of Excited AtomsAtoms

Line Emission Spectra of Excited Line Emission Spectra of Excited AtomsAtoms

• Excited atoms emit light of only certain frequencies. The frequencies are unique to each element.

• This suggests that the electrons are arranged in specific energy levels – these levels are known

20

Spectrum of Spectrum of Excited Hydrogen GasExcited Hydrogen Gas

21

Line Spectra of Other Line Spectra of Other ElementsElements

22

The Electric The Electric PicklePickle

• Excited atoms can emit light.

• Here the solution in a pickle is excited electrically. The Na+ ions in the pickle juice give off light characteristic of that element.

23

Light Spectrum Lab! Slit that Slit that allows light allows light insideinside

Line up the slit so Line up the slit so that it is parallel with that it is parallel with the spectrum tube the spectrum tube (light bulb)(light bulb)

ScaleScale

24

Light Spectrum Lab!

• Run electricity through various gases, creating light

• Look at the light using a spectroscope to separate the light into its component colors

• Using colored pencils, draw the line spectra (all of the lines) and determine the wavelength of the three brightest lines

• Once you line up the slit with the light, then look to the scale on the right. You should see the colored lines under the scale.

Slit that Slit that allows light allows light insideinside

EyepieceEyepiece

ScaleScale

25Light Spectrum Lab!

26

Atomic SpectraAtomic SpectraAtomic SpectraAtomic Spectra

+Electronorbit

One view of atomic structure in early 20th One view of atomic structure in early 20th century was that an electron (e-) traveled century was that an electron (e-) traveled about the nucleus in an orbit.about the nucleus in an orbit.

27

Atomic Spectra and Atomic Spectra and BohrBohr

Atomic Spectra and Atomic Spectra and BohrBohr

Bohr said classical view is wrong. Bohr said classical view is wrong.

Need a new theory — now called Need a new theory — now called QUANTUMQUANTUM or or WAVE MECHANICSWAVE MECHANICS..

e- can only exist in certain discrete e- can only exist in certain discrete orbitsorbits

e- is restricted to e- is restricted to QUANTIZEDQUANTIZED energy energy state (quanta = bundles of energy)state (quanta = bundles of energy)

28

Schrodinger applied idea of e- Schrodinger applied idea of e- behaving as a wave to the behaving as a wave to the problem of electrons in atoms.problem of electrons in atoms.

He developed the He developed the WAVE WAVE EQUATIONEQUATION

Solution gives set of math Solution gives set of math expressions called expressions called WAVE WAVE FUNCTIONS, FUNCTIONS,

Each describes an allowed energy Each describes an allowed energy state of an e-state of an e-

E. SchrodingerE. Schrodinger1887-19611887-1961

Quantum or Wave Quantum or Wave MechanicsMechanics

Quantum or Wave Quantum or Wave MechanicsMechanics

29Heisenberg Heisenberg Uncertainty PrincipleUncertainty Principle

Problem of defining nature Problem of defining nature of electrons in atoms of electrons in atoms solved by W. Heisenberg.solved by W. Heisenberg.

Cannot simultaneously Cannot simultaneously define the position and define the position and momentum (= m•v) of an momentum (= m•v) of an electron.electron.

We define e- energy exactly We define e- energy exactly but accept limitation that but accept limitation that we do not know exact we do not know exact position.position.

Problem of defining nature Problem of defining nature of electrons in atoms of electrons in atoms solved by W. Heisenberg.solved by W. Heisenberg.

Cannot simultaneously Cannot simultaneously define the position and define the position and momentum (= m•v) of an momentum (= m•v) of an electron.electron.

We define e- energy exactly We define e- energy exactly but accept limitation that but accept limitation that we do not know exact we do not know exact position.position.

W. HeisenbergW. Heisenberg1901-19761901-1976

30

Arrangement of Arrangement of Electrons in AtomsElectrons in Atoms

Arrangement of Arrangement of Electrons in AtomsElectrons in Atoms

Electrons in atoms are arranged asElectrons in atoms are arranged as

LEVELSLEVELS (n) (n)

SUBLEVELSSUBLEVELS (l) (l)

ORBITALSORBITALS (m (mll))

31

QUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERS

The The shape, size, and energyshape, size, and energy of each orbital is a function of each orbital is a function of 3 quantum numbers which describe the location of of 3 quantum numbers which describe the location of an electron within an atom or ionan electron within an atom or ion

n n (principal)(principal) ---> energy level---> energy level

ll (orbital) (orbital) ---> shape of orbital---> shape of orbital

mmll (magnetic)(magnetic) ---> designates a particular ---> designates a particular suborbitalsuborbital

The fourth quantum number is not derived from the The fourth quantum number is not derived from the wave functionwave function

ss (spin)(spin) ---> spin of the electron ---> spin of the electron (clockwise or counterclockwise: ½ or – ½)(clockwise or counterclockwise: ½ or – ½)

32

QUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERS

So… if two electrons are in the same place at So… if two electrons are in the same place at the same time, they must be repelling, so at the same time, they must be repelling, so at least the spin quantum number is different!least the spin quantum number is different!

The The Pauli Exclusion PrinciplePauli Exclusion Principle says that no two says that no two electrons within an atom (or ion) can have the electrons within an atom (or ion) can have the same four quantum numbers.same four quantum numbers.

If two electrons are in the same energy level, If two electrons are in the same energy level, the same sublevel, and the same orbital, they the same sublevel, and the same orbital, they must repel.must repel.

Think of the 4 quantum numbers as the address Think of the 4 quantum numbers as the address of an electron… Country > State > City > of an electron… Country > State > City > StreetStreet

33

Energy LevelsEnergy LevelsEnergy LevelsEnergy Levels

• Each energy level has a number Each energy level has a number called thecalled the PRINCIPAL PRINCIPAL QUANTUM NUMBER, nQUANTUM NUMBER, n

• Currently n can be 1 thru 7, Currently n can be 1 thru 7, because there are 7 periods on because there are 7 periods on the periodic tablethe periodic table

34

Energy LevelsEnergy LevelsEnergy LevelsEnergy Levels

n = 1n = 1

n = 2n = 2

n = 3n = 3

n = 4n = 4

35Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

36

Types of Orbitals

• The most probable area to find The most probable area to find these electrons takes on a shapethese electrons takes on a shape

• So far, we have 4 shapes. They So far, we have 4 shapes. They are named s, p, d, and f. are named s, p, d, and f.

• No more than 2 e- assigned to an No more than 2 e- assigned to an orbital – one spins clockwise, one orbital – one spins clockwise, one spins counterclockwisespins counterclockwise

37

Types of Orbitals Types of Orbitals ((ll))

s orbitals orbital p orbitalp orbital d orbitald orbital

38

p Orbitalsp Orbitalsp Orbitalsp Orbitals

this is a this is a p sublevelp sublevel with with 3 orbitals3 orbitals

These are called x, y, and zThese are called x, y, and z

this is a this is a p sublevelp sublevel with with 3 orbitals3 orbitals

These are called x, y, and zThese are called x, y, and z planar node

Typical p orbital

planar node

Typical p orbital

There is a There is a PLANAR PLANAR NODENODE thru the thru the nucleus, which is nucleus, which is an area of zero an area of zero probability of probability of finding an electronfinding an electron

3p3pyy orbital orbital

39

p Orbitalsp Orbitalsp Orbitalsp Orbitals

• The three p orbitals lie 90The three p orbitals lie 90oo apart in space apart in space

40

d Orbitalsd Orbitalsd Orbitalsd Orbitals

• d sublevel has 5 d sublevel has 5 orbitalsorbitals

typical d orbital

planar node

planar node

41

The shapes and labels of the five 3d orbitals.

42

f Orbitalsf Orbitalsf Orbitalsf Orbitals

For l = 3, For l = 3, ---> f sublevel with 7 orbitals---> f sublevel with 7 orbitals

43

Diagonal Rule• Must be able to write it for the test!

This will be question #1 ! Without it, you will not get correct answers !

• The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy

• Aufbau Principle states that electrons fill from the lowest possible energy to the highest energy

44Diagonal Rule – showing sublevels

ss

s 3p 3ds 3p 3d

s 2ps 2p

s 4p 4d 4fs 4p 4d 4f

s 5p 5d 5f 5g?s 5p 5d 5f 5g?

s 6p 6d 6f 6g? 6h?s 6p 6d 6f 6g? 6h?

s 7p 7d 7f 7g? 7h? 7i?s 7p 7d 7f 7g? 7h? 7i?

11

22

33

44

55

66

77

Steps:Steps:

1.1. Write the energy levels top to bottom.Write the energy levels top to bottom.

2.2. Write the orbitals in s, p, d, f order. Write Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy the same number of orbitals as the energy level.level.

3.3. Draw diagonal lines from the top right to the Draw diagonal lines from the top right to the bottom left.bottom left.

4.4. To get the correct order, To get the correct order,

follow the arrows!follow the arrows!

By this point, we are past By this point, we are past the current periodic table the current periodic table so we can stop.so we can stop.

45Rules of electron orbital filling• Aufbau PrincipleAufbau Principle states that electrons fill

from the lowest possible energy to the highest energy.

• Hund’s RuleHund’s Rule In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs. All single electrons must spin the same way

• Pauli Exclusion PrinciplePauli Exclusion Principle says that no two says that no two electrons within an atom (or ion) can have electrons within an atom (or ion) can have the same four quantum numbers.the same four quantum numbers.

• The quantum #’s refer to the principle The quantum #’s refer to the principle energy level (n=1->6), sublevel (spdf), orbital energy level (n=1->6), sublevel (spdf), orbital and spin (up or down)and spin (up or down)

46

Why are d and f orbitals always in lower energy levels?

• d and f orbitals require LARGE amounts of energy

• It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy

This is the reason for the diagonal rule! BE SURE TO FOLLOW THE ARROWS IN ORDER!

47

s orbitalss orbitals d orbitalsd orbitals

Number ofNumber oforbitalsorbitals

Number of Number of electronselectrons

p orbitalsp orbitals f orbitalsf orbitals

How many electrons can be in a sublevel?How many electrons can be in a sublevel?

Remember: A maximum of two electrons can Remember: A maximum of two electrons can be placed in an orbital.be placed in an orbital.

11 3 5 73 5 7

22 6 10 146 10 14

48

Electron ConfigurationsA list of all the electrons in an atom (or ion)

• Must go in order (Aufbau principle)

• 2 electrons per orbital, maximum

• We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons.

• The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule

1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p66 4s 4s22 3d 3d1010 4p 4p66 5s 5s22 4d 4d1010 5p 5p66 6s 6s22 4f 4f1414…… etc.etc.

49

Electron Configurations

2p4

Energy LevelEnergy Level

SublevelSublevel

Number of Number of electrons in electrons in the sublevelthe sublevel

1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p66 4s 4s22 3d 3d1010 4p 4p66 5s 5s22 4d 4d1010 5p 5p66 6s6s22 4f 4f1414…… etc.etc.

50Let’s Try It!

• Write the electron configuration for the following elements:

H

Li

N

Ne

K

Zn

Pb

51

An excited lithium atom emitting a photon of red light to drop to a

lower energy state.

52An excited H atom returns to a

lower energy level.

53Orbitals and the Orbitals and the Periodic TablePeriodic Table

• Orbitals grouped in s, p, d, and f orbitals Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental)(sharp, proximal, diffuse, and fundamental)

s orbitalss orbitalsp orbitalsp orbitals

d orbitalsd orbitals

f orbitalsf orbitals

54

Shorthand Notation

• A way of abbreviating long electron configurations

• Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

55

Shorthand Notation

• Step 1: It’s the Showcase Showdown!Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].

• Step 2: Find where to resume by finding the next energy level.

• Step 3: Resume the configuration until it’s finished.

56

Shorthand Notation• Chlorine

– Longhand is 1s2 2s2 2p6 3s2 3p5

You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6

The next energy level after Neon is 3

So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17

[Ne] 3s2 3p5

57

Practice Shorthand Notation

• Write the shorthand notation for each of the following atoms:

Cl K Ca I Bi Hg AgAl

58

Valence ElectronsValence ElectronsValence ElectronsValence ElectronsElectrons are divided between core and Electrons are divided between core and

valence electronsvalence electronsB 1sB 1s22 2s 2s22 2p 2p11

Core = [He]Core = [He] , , valence = 2svalence = 2s22 2p 2p11

Br [Ar] 3dBr [Ar] 3d1010 4s 4s22 4p 4p55

Core = [Ar] 3dCore = [Ar] 3d1010 , , valence = 4svalence = 4s22 4p 4p55

59

Rules of the GameRules of the GameRules of the GameRules of the GameNo. of valence electrons of a main group No. of valence electrons of a main group

atom = Group numberatom = Group number (for A groups) (for A groups)

Atoms like to either empty or fill their outermost Atoms like to either empty or fill their outermost level. Since the outer level contains two s level. Since the outer level contains two s electrons and six p electrons (d & f are always in electrons and six p electrons (d & f are always in lower levels), the optimum number of electrons lower levels), the optimum number of electrons is eight. This is called the is eight. This is called the octet rule.octet rule.

60

Keep an Eye On Those Ions!

• Electrons are lost or gained like they always are with ions… negative ions have gained electrons, positive ions have lost electrons

• The electrons that are lost or gained should be added/removed from the highest energy level (not the highest orbital in energy!)

61


• Tin

Atom: [Kr] 5s2 4d10 5p2

Sn+4 ion: [Kr] 4d10

Sn+2 ion: [Kr] 5s2 4d10

Note that the electrons came out of the highest energy level, not the highest energy orbital!

62


• Bromine

Atom: [Ar] 4s2 3d10 4p5

Br- ion: [Ar] 4s2 3d10 4p6

Note that the electrons went into the highest energy level, not the highest energy orbital!

63Try Some Ions!

• Write the longhand notation for these:

F-

Li+

Mg+2

• Write the shorthand notation for these:

Br-

Ba+2

Al+3

64

Exceptions to the Aufbau Principle

• Remember d and f orbitals require LARGE amounts of energy

• If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel)

• There are many exceptions, but the most common ones are

d4 and d9

For the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom.

(HONORS only)

65Exceptions to the Aufbau Principle

d4 is one electron short of being HALF full

In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4.

For example: Cr would be [Ar] 4s2 3d4, but since this ends exactly with a d4 it is an exception to the rule. Thus, Cr should be [Ar] 4s1 3d5.

Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d.

(HONORS only)

66Exceptions to the Aufbau

PrincipleOK, so this helps the d, but what about the

poor s orbital that loses an electron?

Remember, half full is good… and when an s loses 1, it too becomes half full!

So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital.

(HONORS only)

67Exceptions to the Aufbau Principle

d9 is one electron short of being full

Just like d4, one of the closest s electrons will go into the d, this time making it d10 instead of d9.

For example: Au would be [Xe] 6s2 4f14 5d9, but since this ends exactly with a d9 it is an exception to the rule. Thus, Au should be [Xe] 6s1 4f14 5d10.

Procedure: Same as before! Find the closest s orbital. Steal one electron from it, and add it to the d.

(HONORS only)

68

Try These!

• Write the shorthand notation for:

Cu

W

Au

(HONORS only)

69

Orbital Diagrams

• Graphical representation of an electron configuration

• One arrow represents one electron

• Shows spin and which orbital within a sublevel

• Same rules as before (Aufbau principle, d4 and d9 exceptions, two electrons in each orbital, etc. etc.)

70Orbital Diagrams

• One additional rule: Hund’s Rule

– In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs

– All single electrons must spin the same way

• I nickname this rule the “Monopoly Rule”

• In Monopoly, you have to build houses EVENLY. You can not put 2 houses on a property until all the properties has at least 1 house.

71

LithiumLithiumLithiumLithium

Group 1AGroup 1A

Atomic number = 3Atomic number = 3

1s1s222s2s11 ---> 3 total electrons ---> 3 total electrons

1s

2s

3s3p

2p

72

CarbonCarbonCarbonCarbon

Group 4AGroup 4A

Atomic number = 6Atomic number = 6

1s1s2 2 2s2s2 2 2p2p22 ---> --->

6 total electrons6 total electrons

Here we see for the first time Here we see for the first time

HUND’S RULEHUND’S RULE. When . When placing electrons in a set of placing electrons in a set of orbitals having the same orbitals having the same energy, we place them singly energy, we place them singly as long as possible.as long as possible.1s

2s

3s3p

2p

73Lanthanide Element Lanthanide Element

ConfigurationsConfigurations

4f orbitals used for Ce - Lu and 5f for Th - Lr

4f orbitals used for Ce - Lu and 5f for Th - Lr

74

Draw these orbital diagrams!

• Oxygen (O)

• Chromium (Cr)

• Mercury (Hg)

75

Ion ConfigurationsIon ConfigurationsIon ConfigurationsIon Configurations

To form anions from elements, add 1 or more To form anions from elements, add 1 or more e- from the highest sublevel.e- from the highest sublevel.

P [Ne] 3sP [Ne] 3s22 3p 3p33 + 3e- ---> P + 3e- ---> P3-3- [Ne] 3s [Ne] 3s22 3p 3p66 or [Ar] or [Ar]

1s

2s

3s3p

2p

1s

2s

3s3p

2p

76General Periodic General Periodic TrendsTrends

• Atomic and ionic sizeAtomic and ionic size

• Ionization energyIonization energy

• ElectronegativityElectronegativity

Higher effective nuclear chargeElectrons held more tightly

Larger orbitals.Electrons held lesstightly.

77

Atomic Atomic SizeSize

Atomic Atomic SizeSize

• Size goes UPSize goes UP on going down a group. on going down a group. • Because electrons are added further Because electrons are added further

from the nucleus, there is less from the nucleus, there is less attraction. This is due to additional attraction. This is due to additional energy levels and the energy levels and the shielding effectshielding effect. . Each additional energy level “shields” Each additional energy level “shields” the electrons from being pulled in the electrons from being pulled in toward the nucleus.toward the nucleus.

• Size goes DOWNSize goes DOWN on going across a on going across a period.period.

• Size goes UPSize goes UP on going down a group. on going down a group. • Because electrons are added further Because electrons are added further

from the nucleus, there is less from the nucleus, there is less attraction. This is due to additional attraction. This is due to additional energy levels and the energy levels and the shielding effectshielding effect. . Each additional energy level “shields” Each additional energy level “shields” the electrons from being pulled in the electrons from being pulled in toward the nucleus.toward the nucleus.

• Size goes DOWNSize goes DOWN on going across a on going across a period.period.

78

Atomic SizeAtomic SizeAtomic SizeAtomic Size

Size Size decreasesdecreases across a period across a period owing to increase in the positive owing to increase in the positive charge from the protons. Each added charge from the protons. Each added electron feels a greater and greater + electron feels a greater and greater + charge because the protons are pulling charge because the protons are pulling in the same direction, where the in the same direction, where the electrons are scattered.electrons are scattered.

LargeLarge SmallSmall

79

Which is Bigger?Which is Bigger?

• Na or K ?Na or K ?

• Na or Mg ?Na or Mg ?

• Al or I ?Al or I ?

80

Ion SizesIon SizesIon SizesIon Sizes

Li,152 pm3e and 3p

Li+, 60 pm2e and 3 p

+Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?

Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?

81


• CATIONSCATIONS are are SMALLERSMALLER than the than the atoms from which they come.atoms from which they come.

• The electron/proton attraction has The electron/proton attraction has gone UP and so size gone UP and so size DECREASESDECREASES..

Li,152 pm3e and 3p

Li +, 78 pm2e and 3 p

+Forming Forming a cation.a cation.Forming Forming a cation.a cation.

82


F,64 pm9e and 9p

F- , 136 pm10 e and 9 p

-Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?

Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?

83


• ANIONSANIONS are are LARGERLARGER than the atoms from than the atoms from which they come.which they come.

• The electron/proton attraction has gone DOWN The electron/proton attraction has gone DOWN and so size and so size INCREASESINCREASES..

• Trends in ion sizes are the same as atom sizes. Trends in ion sizes are the same as atom sizes.

Forming Forming an anion.an anion.Forming Forming an anion.an anion.F, 71 pm

9e and 9pF-, 133 pm10 e and 9 p

-

84

Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes

Figure 8.13Figure 8.13

85

Which is Bigger?Which is Bigger?

• Cl or ClCl or Cl-- ? ?

• KK++ or K ? or K ?

• Ca or CaCa or Ca+2+2 ? ?

• II-- or Br or Br-- ? ?

86

Mg (g) + Mg (g) + 738 kJ738 kJ ---> Mg ---> Mg++ (g) + e- (g) + e-

This is called the FIRST This is called the FIRST ionization energy because ionization energy because

we removed only the we removed only the OUTERMOST electronOUTERMOST electron

MgMg+ + (g) + (g) + 1451 kJ1451 kJ ---> Mg ---> Mg2+2+ (g) + e- (g) + e-This is the SECOND IE.This is the SECOND IE.

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).

Ionization EnergyIonization EnergyIonization EnergyIonization Energy

87Trends in Ionization Trends in Ionization EnergyEnergy

Trends in Ionization Trends in Ionization EnergyEnergy

• IE increases across a IE increases across a period because the period because the positive charge increases.positive charge increases.

• Metals lose electrons Metals lose electrons more easily than more easily than nonmetals.nonmetals.

• Nonmetals lose electrons Nonmetals lose electrons with difficulty (they like to with difficulty (they like to GAIN electrons).GAIN electrons).

88



• IE increases UP a IE increases UP a group group

• Because size Because size increases increases (Shielding Effect)(Shielding Effect)

89

Which has a higher 1st ionization energy?

• Mg or Ca ?

• Al or S ?

• Cs or Ba ?

90

Electronegativity, Electronegativity,

is a measure of the ability of an atom is a measure of the ability of an atom in a molecule to attract electrons to in a molecule to attract electrons to itself.itself.

Concept proposed byConcept proposed byLinus PaulingLinus Pauling1901-19941901-1994

Concept proposed byConcept proposed byLinus PaulingLinus Pauling1901-19941901-1994

91Periodic Trends: Electronegativity

• In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements.

• In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.

92

ElectronegativityElectronegativity

93

Which is more electronegative?

•F or Cl ?

•Na or K ?

•Sn or I ?

94

The End !!!!!!!!!!!!!!!!!!!

1 modern atomic theory (a.k.a. the electron chapter!) chapter 4

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