1 modern atomic theory (a.k.a. the electron chapter!) chemistry 1: chapters 5, 6, and 7 chemistry 1...

1Modern Atomic Theory

(a.k.a. the electron chapter!)Chemistry 1: Chapters Chemistry 1: Chapters 5, 6, and 75, 6, and 7

Chemistry 1 Honors: Chemistry 1 Honors: Chapter 11Chapter 11

2

ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATION

ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATION

3Electromagnetic radiation.

4

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.

5

wavelength Visible light

wavelength

Ultaviolet radiation

Amplitude

Node



6

• Waves have a Waves have a frequencyfrequency• Use the Greek letter “nu”, ν, for frequency, and Use the Greek letter “nu”, ν, for frequency, and

units are “cycles per sec” (Hertz, Hz)units are “cycles per sec” (Hertz, Hz)

• All radiation: All radiation: νν • • = c = cwhere c = velocity of light = 3.00 x 10where c = velocity of light = 3.00 x 1088 m/sec m/sec



7

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

8

ElectroElectromagneticmagnetic SpectrumSpectrum

ElectroElectromagneticmagnetic SpectrumSpectrum

In increasing energy, RIn increasing energy, ROOYY GG BBIIVV

9Excited Gases Excited Gases & Atomic & Atomic StructureStructure

10

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

• Problem is that the model only Problem is that the model only works for Hworks for H

Niels BohrNiels Bohr

(1885-1962)(1885-1962)

11

Spectrum of White Spectrum of White LightLight

12Line Emission Line Emission Spectra Spectra

of Excited Atomsof Excited Atoms

Line Emission Line Emission Spectra Spectra

of Excited Atomsof Excited Atoms• Excited atoms emit light of only

certain wavelengths

• The wavelengths of emitted light depend on the element.

13

Spectrum of Spectrum of Excited Hydrogen GasExcited Hydrogen Gas

14

Line Spectra of Other Line Spectra of Other ElementsElements

15

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

lower energy state.

16

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

17

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

18Light Spectrum Lab!

19

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.

20

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.

21

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)

22

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, E. SchrodingerE. Schrodinger

1887-19611887-1961

Quantum or Wave Quantum or Wave MechanicsMechanics

Quantum or Wave Quantum or Wave MechanicsMechanics

23Heisenberg 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

24

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))

25

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)

26

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

27

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

28

Energy LevelsEnergy LevelsEnergy LevelsEnergy Levels

n = 1n = 1

n = 2n = 2

n = 3n = 3

n = 4n = 4

29Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

30

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

31

Types of Orbitals Types of Orbitals ((ll))

s orbitals orbital p orbitalp orbital d orbitald orbital

32

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

33

p Orbitalsp Orbitalsp Orbitalsp Orbitals

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

34

d Orbitalsd Orbitalsd Orbitalsd Orbitals

• d sublevel has 5 d sublevel has 5 orbitalsorbitals

typical d orbital

planar node

planar node

35

The shapes and labels of the five 3d orbitals.

36

f Orbitalsf Orbitalsf Orbitalsf Orbitals

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

37

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

38Diagonal Rule

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.

39

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 orbitals) 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!

40

s orbitalss orbitals d orbitalsd orbitals

Number ofNumber oforbitalsorbitals

Number of Number of electronselectrons

11 33 55

22 66 1010

p orbitalsp orbitals f orbitalsf orbitals

77

1414

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.

41

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.

42

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.

43Let’s Try It!

• Write the electron configuration for the following elements:

H

Li

N

Ne

K

Zn

Pb

44Let’s Try It!

• Write the electron configuration for the following elements:H 1s1

Li 1s2 2s1

N 1s2 2s2 2p3

Ne 1s2 2s2 2p6

K 1s2 2s2 2p6 3s2 3p6 4s1

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

Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2

45Orbitals 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

46

Shorthand Notation

• An abbreviation for 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

47

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.

48

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

49

Practice Shorthand Notation

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

Cl

K

Ca

I

Bi

[Ne] 3s2 3p5

[Ar] 4s1

[Ar] 4s2

[Kr] 5s2 4d10 5p5

[Xe] 6s2 4f14 5d10 6p3

50

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

51

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.

52

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 outermost energy level (not the highest orbital in energy!)

53


• 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 outermost energy level, not the highest energy orbital!

54


• Bromine

Atom: [Ar] 4s2 3d10 4p5

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

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

55Try Some Ions!

• Write the longhand notation for these:

F-

Li+

Mg+2

• Write the shorthand notation for these:

Br-

Ba+2

Al+3

1s2 2s2 2p6

1s2

1s2 2s2 2p6 note this is the same as F- this is called isoelectronic

[Kr]

[Xe]

[Ne]

56

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)

57Exceptions 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)

58Exceptions 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)

59

Try These!

• Write the shorthand notation for:

Cu

W

Au

[Ar] 4s1 3d10

[Xe] 6s1 4f14 5d5

[Xe] 6s1 4f14 5d10

(HONORS only)

60

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.)

61Orbital 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.

62

LithiumLithiumLithiumLithium

Group 1AGroup 1A

Atomic number = 3Atomic number = 3

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

1s

2s

3s3p

2p

63

CarbonCarbonCarbonCarbon

Group 4AGroup 4A


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

64Lanthanide 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

65

Draw these orbital diagrams!

• Oxygen (O)

• Chromium (Cr)

• Mercury (Hg)

66

OxygenOxygenOxygenOxygen

Group 6AGroup 6A


1s1s2 2 2s2s2 2 2p2p44 ---> --->

8 total electrons8 total electrons

1s

2s

3s3p

2p

67

Chromium

68Mercury

69

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

70General 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.Shielding Effect!

71

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 1) additional attraction. This is due to 1) additional energy levels and 2) the energy levels and 2) the shielding shielding effecteffect. Each additional energy level . Each additional energy level “shields” the electrons from being “shields” the electrons from being pulled in toward the nucleus.pulled in toward the nucleus.

• Size goes UPSize goes UP going Right to Left going Right to Left across a period.across a 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 1) additional attraction. This is due to 1) additional energy levels and 2) the energy levels and 2) the shielding shielding effecteffect. Each additional energy level . Each additional energy level “shields” the electrons from being “shields” the electrons from being pulled in toward the nucleus.pulled in toward the nucleus.

• Size goes UPSize goes UP going Right to Left going Right to Left across a period.across a period.

73

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

74

Which is Bigger?Which is Bigger?

• Na or K ?Na or K ?

• Na or Mg ?Na or Mg ?

• Al or I ?Al or I ?

KK

NaNa

II

75

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?

76


• 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.

77


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?

78


• 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 pm

10 e and 9 p

-

79

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

Figure 8.13Figure 8.13

80

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-- ? ?

ClCl--

KK

CaCa

II--

81

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

82Trends 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).

83



• IE increases UP a IE increases UP a group group

• Because size Because size increases (Shielding increases (Shielding Effect & Increased Effect & Increased Distance from Distance from Nucleus)Nucleus)

84

Which has a higher 1st ionization energy?

• Mg or Ca ?

• Al or S ?

• Cs or Ba ?

Mg

S

Ba

85

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

86Periodic Trends: Electronegativity

• In a group: Atoms with fewer energy levels can attract electrons better (less shielding), and are closer to the nucleus. 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.

87

ElectronegativityElectronegativity

88

Which is more electronegative?

•F or Cl ?

•Na or K ?

•Sn or I ?

F

Na

I

89

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

1 modern atomic theory (a.k.a. the electron chapter!) chemistry 1: chapters 5, 6, and 7 chemistry 1...

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