1

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))

2

QUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERS

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

n n (principal)(principal) ---> energy level (1, 2, 3…7)---> energy level (1, 2, 3…7)

ll (orbital) (orbital) ---> shape of orbital (s, p, d, f)---> shape of orbital (s, p, d, f)

mmll (magnetic)(magnetic) ---> designates a particular ---> designates a particular suborbital (px,py,pz) (d- 5 orientations, f-7 )suborbital (px,py,pz) (d- 5 orientations, f-7 )

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

3

QUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERS

So… if two electrons are in the same place at the same So… if two electrons are in the same place at the same time, they must be repelling, so at least the spin time, they must be repelling, so at least the spin quantum number is different!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 same electrons within an atom (or ion) can have the same four quantum numbers.four quantum numbers.

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

Think of the 4 quantum numbers as the address of an Think of the 4 quantum numbers as the address of an electron… State > City > Street> House Numberelectron… State > City > Street> House Number

4

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

5

Energy LevelsEnergy LevelsEnergy LevelsEnergy Levels

n = 1n = 1

n = 2n = 2

n = 3n = 3

n = 4n = 4

6Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

7

Types of Orbitals

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

• So far, we have 4 shapes. They are So far, we have 4 shapes. They are named s, p, d, and f (sharp or named s, p, d, and f (sharp or spherical, principal, diffuse, spherical, principal, diffuse, fundamental). fundamental).

• 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

8

Types of Orbitals Types of Orbitals ((ll))

s orbitals orbital p orbitalp orbital d orbitald orbital

9

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

10

p Orbitalsp Orbitalsp Orbitalsp Orbitals

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

11

d Orbitalsd Orbitalsd Orbitalsd Orbitals

• d sublevel has 5 d sublevel has 5 orbitalsorbitals

typical d orbital

planar node

planar node

12

The shapes and labels of the five 3d orbitals.

13

f Orbitalsf Orbitalsf Orbitalsf Orbitals

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

14Diagonal Rule (aufbau principle)

• 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 /Diagonal rule states that electrons fill from the lowest possible energy to the highest energy

15Diagonal 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.

16

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!

17

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.

18

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.

19

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.

20Let’s Try It!

• Write the electron configuration for the following elements:

H

Li

N

Ne

K

Zn

Pb

21

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

lower energy state.

22An excited H atom returns to a

lower energy level.

23

Determine element when elec.conf. is given

1. 1s2 2s2 2p6 3s2 3p3.

2. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3

3. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2

24Orbitals 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

25

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

26

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.

27

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

28

Practice Shorthand Notation

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

Cl

K

Ca

I

Bi

29

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

30

Electron Dot structures(Lewis structures)

• Shorthand visual method to show valence electrons- dots represent electrons in pairs.

• P162 Practice- a.Draw structures for- Mg, Tl, Xe.

• B. An atom of an element has a total of 13 electrons. What is it, and how many electrons are shown in its dot structure?

• C. Out of the elements - C, Ge, S, Be or Ar, which one has the dot str-

°X °

31

P 167 practice

Q 85- Write orbital diagram and elec. Conf. for

• Beryllium, aluminum, nitrogen, sodium

Q 86- Write shorthand notation of-

- Kr, Zr, P, Pb

Q 87- Which element is shown-

- 1s2 2s2 2p5

- (Ar) 4s2

- (Xe) 6s2 4f4

- (Kr) 5s2 4d10 5p4

- (Rn) 7s2 5f13

Q 90- Draw Lewis structures of-

- C, As, Po, K, Ba

32

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

33

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!)

34

Keep an Eye On Those Ions!

• 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!

35

Keep an Eye On Those Ions!

• 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!

36Try Some Ions!

• Write the longhand notation for these:

F-

Li+

Mg+2

• Write the shorthand notation for these:

Br -

Ba+2

Al+3

37

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.

38Exceptions 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.

39Exceptions 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.

40Exceptions 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.

41

Try These!

• Write the shorthand notation for:

Cu

W

Au

42

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, two electrons in each orbital, etc.)

43Orbital 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

• This rule is nicknamed 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.

44

LithiumLithiumLithiumLithium

Group 1AGroup 1A

Atomic number = 3Atomic number = 3

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

1s

2s

3s3p

2p

45

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

46Lanthanide 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

47

Draw these orbital diagrams!

• Oxygen (O), Chromium (Cr), Mercury (Hg)

• In excited state, electrons may jump to orbitals they would normally not occupy because they have extra energy.

48

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

49Heisenberg Heisenberg Uncertainty PrincipleUncertainty Principle

It is not possible to pinpoint the It is not possible to pinpoint the exact position of an electron exact position of an electron within an atom.within an atom.

Cannot simultaneously define the Cannot simultaneously define the position and momentum (= position and momentum (= m•v) of an electron- since you m•v) of an electron- since you need light energy to spot it.The need light energy to spot it.The electron absorbs this photon of electron absorbs this photon of energy and changes its energy and changes its positionposition

It is not possible to pinpoint the It is not possible to pinpoint the exact position of an electron exact position of an electron within an atom.within an atom.

Cannot simultaneously define the Cannot simultaneously define the position and momentum (= position and momentum (= m•v) of an electron- since you m•v) of an electron- since you need light energy to spot it.The need light energy to spot it.The electron absorbs this photon of electron absorbs this photon of energy and changes its energy and changes its positionposition

W. HeisenbergW. Heisenberg1901-19761901-1976

50

Electron configuration Practice

51Development of the Periodic Table

• In the 1700s, Lavoisier compiled a list of all the known elements of the time.

• The 1800s brought large amounts of information and scientists needed a way to organize knowledge about elements.

• John Newlands proposed an arrangement where elements were ordered by increasing atomic mass.

• Newlands noticed when the elements were arranged by increasing atomic mass, their properties repeated every eighth element. (NEWLANDS OCTAVES)

52

• Meyer and Mendeleev both demonstrated a connection between atomic mass and elemental properties.

• Moseley rearranged the table by increasing atomic number, and resulted in a clear periodic pattern.

53

The Periodic Law

• Dmitri Mendeleev gave us a functional scheme with which to classify elements.

– Mendeleev’s scheme was based on chemical properties of the elements.

– It was noticed that the chemical properties of elements increased in a periodic (repeating after regular intervals) manner.

– The periodicity of the elements was demonstrated by Mendeleev when he used the table to predict to occurrence and chemical properties of elements which had not yet been discovered.

54MENDELEEV- “FATHER OF THE MODERN PERIODIC TABLE”

• Mendeleev left blank spaces in his table when the properties of the elements above and below did not seem to match.

• The existence of unknown elements was predicted by Mendeleev on the basis of the blank spaces.

• When the unknown elements were discovered, it was found that Mendeleev had closely predicted the properties of the elements as well as their discovery.

55

Blank spaces in Mendeleev’s Table

56

The Periodic Law

– Similar physical and chemical properties recur (happen again) periodically when the elements are listed in order of increasing atomic number.

57

The Modern Periodic Table

– The periodic table is made up of rows of elements and columns.

– An element is identified by its chemical symbol.

– The number above the symbol is the atomic number

– The number below the symbol is the rounded atomic weight of the element.

– A row (horizontal) is called a period– A column (vertical) is called a group (or

family)

58

Periodic Patterns

– The chemical behavior of elements is determined by its electron configuration (how electrons are distributed in shells).

– Energy levels are quantized so roughly correspond to layers of electrons around the nucleus.

– A shell is all the electrons with the same value of n.

» n is a row in the periodic table.– Each period begins with a new outer

electron shell

59

Chemical “Families”

– IA are called alkali metals because the react with water to from an alkaline solution

– Group IIA are called the alkali earth metals because they are reactive, but not as reactive as Group IA.

» They are also soft metals like Earth.– Group VIIA are the halogens

» These need only one electron to fill their outer shell

» They are very reactive.– Group VIIIA are the noble gases as they have

completely filled outer shells» They are almost non reactive.

60

Metals, Non-metals and Metalloids

Metal: Elements that are usually solids at room temperature. Most elements are metals.

Non-Metal: Elements in the upper right corner of the periodic Table. Their chemical and physical properties are different from metals.

Metalloid: Elements that lie on a diagonal line between the Metals and non-metals. Their chemical and physical properties are intermediate between the two.

TRANSITION ELEMENTS- D-block elements.

61

P 181 assessment

2. Sketch a simple Periodic table and show the location of metals, non-metals and metalloids on it.

4. Identify the transition metals out of these-

a.Li b. Pt c. Pm d. C

5. For each of the given elements, list 2 other elements with similar chemical properties-

a. Iodine b. Barium c. Iron

6. In one sentence each, describe the contribution of Newlands, Lavoisier, Moseley and Mendeleev.

62General 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.

63

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.

64

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

65

Which is Bigger?Which is Bigger?• Na or K ?Na or K ?• Na or Mg ?Na or Mg ?• Al or I ? (Hint: Atomic size Al or I ? (Hint: Atomic size

shrinks greatly on going shrinks greatly on going across and does not across and does not increase as much on going increase as much on going down a group).down a group).

66

67

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?

68

Ion SizesIon SizesIon SizesIon Sizes

• 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.

69

Ion SizesIon SizesIon SizesIon Sizes

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?

70

Ion SizesIon SizesIon SizesIon Sizes

• 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

-

71

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

Figure 8.13Figure 8.13

72

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-- ? ?

73

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

74Trends 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).

75

Trends in Ionization Trends in Ionization EnergyEnergy

Trends in Ionization Trends in Ionization EnergyEnergy

• IE decreases down a IE decreases down a group group

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

76

Which has a higher 1st ionization energy?

• Mg or Ca ?

• Al or S ?

• Cs or Ba ?

77

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

78Periodic Trends: Electronegativity

• In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity decreases down 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.

79

ElectronegativityElectronegativity

80

Which is more electronegative?

•F or Cl ?

•Na or K ?

•Sn or I ?

81

Trends in reactivity• Metals• Period - reactivity decreases as you go

from left to right across a period.• Group - reactivity increases as you go

down a group

• Non-metals• Period - reactivity increases as you go

from the left to the right across a period.

• Group - reactivity decreases as you go down the group.

82

Periodic Trends Worksheet

• Rank the following elements by increasing atomic radius: C, Al, O, K.

• Rank the following elements by increasing electronegativity: S, O, Ne, Al.

• What is the difference between electron affinity and ionization energy?

• Why does fluorine have a higher ionization energy than iodine?

• Why do elements in the same family generally have similar properties?

83

P 199 practice• Q 41. Why do the elements chlorine and iodine have

similar chemical properties?

• Q 43. How many valence electrons does each noble gas have?

• Q 44. What are the 4 blocks of the periodic table?

• Q 45. What electron configuration has the greatest stability?

• Q64. Which element has the larger ionization energy

• A. Li, N B. Kr, Ne C. Cs, Li

• Q 78. Which element in each pair is more electronegative?

• A. K, As b. N, Sb c. Sr, Be

84

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